General Chemistry

Celsius

(C) A temperature scale defined by having 0 *C equal to the freezing point of water and 100*C equal to the boiling point of water; otherwise known as the centigrade temperature scale

Atomic mass unit

(amu) A unit of mass defined as 1/12 the mass of a carbon-12 atom; approximately equal to the mass of one proton or one neutron

specific heat

(c) the amount of heat required to raise the temperature of one gram of a substance by 1 deg C (or one kelvin); specific heat values will generally be provided on test day but one constant to remember is the specific heat of H20 (l): cH2O = 1 cal/g*K

standard enthalpy of formation

(deltaH&o f) the enthalpy required to produce one mole of a compound from its elements in their standard states; remember that standard state refers to the most stable physical state of an element or compound at 298 K and 1 atm; of note delta H^osubf of an element in its standard state by definition is zero;

standard enthalpy of a reaction

(deltaH^osubrxn) is the enthalpy change accompanying a reaction being carried out under standard conditions; this can be calculated by taking the difference between the sum of the standard heats of formation for the products and the sum of the standard heats of formation for the products and the sum of the standard heats of formation of the reactants

Heat of formation

(deltaHf) The heat absorbed or released during the formation of a pure substance from its elements at a constant pressure

Electron

(e-) A subatomic particle that remains outside the nucleus and carries a single negative charge; in most cases, its mass is considered to be negligible

Electromotive force

(emf) The potential difference developed between the cathode and the anode of an electrochemical cell; also called voltage; if the emf is positive the cell is able to release energy (delta G <0) which means it is spontaneous; if the emf is negative the cell must absorb energy (deltaG > 0) which means it is spontaneous

heat capacity

(mc) the product of mass and specific heat and is the energy required to raise any given amount of a substance one degree Celsius

Density

(p) *fancy; A physical property of a substance, defined as the mass contained in a unit of volume of a substance; we can rearrange PV=nRT where n = m/M therefore PV=(m/M)RT and p=m/v=PM/RT; we can also find p by p=m/V2 from the equation P1V1/T1 = P2V2/T2; remember a mole of ideal gas at STP occupies 22.4 L and temperature is 546 K

Homogeneous catalyst

A catalyst that is in the same phase of matter as the reactants (for example, an aqueous enzyme in the cytoplasm of a cell)

Diamagnetism

A condition that arises when a substance has no unpaired electrons and is slightly repressed by a magnetic field; materials consisting of atoms that have only paired electrons will be slightly repelled by a magnetic field and are said to be diamagnetic, this is shown with a piece of pyrolytic graphite that is suspended in the air over strong neodymium magnets; all of the electrons in this allotrope (configuration) of carbon are paired because of covalent bonding between layers of the material, and are thus opposed to being reoriented; given sufficiently strong magnetic fields beneath an object, any diamagnetic substance can be made to levitate; example of this is the high speed rail networks such as Japan's SCMaglev

Coordinate covalent bond

A covalent bond in which both electrons of the bonding pair are donated by one of the bonded atoms; generally this means that a lone pair of one atom attacked another atom with an unhybridized p-orbital to form a bond; once such a bond forms it is indistinguishable form any other covalent bond; ex of this is BF3 + NH3, BF3 is a lewis acid, meaning it will accept a lone pair of electrons, and NH3 is a lewis base, meaning it will donate a lone pair of electrons

Combined Gas Law

A gas law hat combines Boyle's Law, Charles Law, and Gay-Lussac's Law to state that pressure and volume are inversely proportional to each other, and each is directly proportional to temperature (P1V1)/T1= (P2V2)/T2; where the subscription 1 and 2 refer to the two states of the gas (at STP and at the conditions of actual temperature and pressure for example) this equation assumes the number of moles stays constant; STP mole of gas occupies 22.4 L and temperature is 246K and pressure is 1 atm

Ideal gas

A hypothetical gas with behavior that is described by the ideal gas law under all conditions; assumes that its particles have zero volume and do not exhibit interactive forces; real gases deviate from this ideal behavior at high pressure (low volumes) and low temperatures due to intermolecular forces or volume effects, many compressed real gases demonstrate behavior that is close to ideal

Conductor

A material in which electrons are able to transfer energy in the form of heat or electricity

Equivalent

A mole of charge in the form of electrons, protons, ions, or other measurable quantities that are produced by a substance

Antibonding orbital

A molecular orbital formed by the overlap of two or more atomic orbitals; energy is greater than the energy of the combining atomic orbitals

Bonding orbital

A molecular orbital formed by the overlap of two or more atomic orbitals; energy is less than that of the combining orbitals

Intermediate

A molecule that transiently exists in a multistep reaction; does not appear in the overall balanced equation; reaction intermediates are often difficult to detect because they may be consumed almost immediately after they are formed, but a proposed mechanism that includes intermediates can be supported through kinetic experiments example: A2 + 2 B --> 2AB step 1: A2 + B --> A2B (slow) step 2: A2B + B --> 2AB (fast) -A2B does not appear in the overall reaction and is the intermediate

Isobaric process

A process that occurs at constant pressure; isothermal and isobaric processes are common because it is usually easy to control temperature and pressure; isobaric processes do not alter the first law, but note that an isobaric process appears as a flat line on a pressure-volume graph

Adiabatic process

A process that occurs without the transfer of heat into or out of the system, thus the thermal energy of the system is constant throughout the process; when Q = 0, the first law simplifies to delta U = -W (the change in internal energy of the system is equal to work done on the system, this is the opposite of work done by the system); this process also appears as hyperbolic on a pressure-volume graph

Bronsted-Lowry base

A proton acceptor; a species that accepts hydrogen ions (H+); this is a more inclusive definition as it is not limited to aqueous solutions; every Arrhenius acid or base can also be classified as a bronsted lowry acid or base; every bronsted lowery acid or base can also be classified as a lewis acid or base, however this logic does not always work the other way (for example, NH3 is a bronsted lowry base, but not an Arrhenius base); mnemonic: the brOnsted-lOwry definition revolves around prOtOns; the lEwis definition around Electrons

Bronsted-Lowry acid

A proton donor; a species that donates hydrogen ions (H+); this is a more inclusive definition as it is not limited to aqueous solutions; every Arrhenius acid or base can also be classified as a bronsted lowry acid or base; every bronsted lowery acid or base can also be classified as a lewis acid or base, however this logic does not always work the other way (for example, NH3 is a bronsted lowry base, but not an Arrhenius base); these always occur in pairs as the definition require the transfer of a proton from the acid to the base; these are conjugate acid-base pairs; mnemonic: the brOnsted-lOwry definition revolves around prOtOns; the lEwis definition around Electrons

Compound

A pure substance that can be decomposed to produce elements, other compounds, or both

Complexation reaction

A reaction in which a central cation is bound to one or more ligands

Aqueous solution

A solution in which water is the solvent; this is also called hydration; with acids the hydronium ion (H3O+) can occur; this is facilitated by the transfer of hydrogen ions (H+) from a molecule in solution to a water molecule (H2O); it is important to realize that H+ is never found alone in solution because a free proton is difficult to isolate; rather it is found bonded to an electron pair donor (carrier) molecule such as a water molecule; this is an example of a coordinate covalent bond

Concentrated solution

A solution with a high concentration value; the cut off for the term "concentrated" depends not he purpose and identity of the solution

Dilute solutions

A solution with a low concentration of a given solute

Amphoteric species

A species capable of reacting as either an acid or base, depending on the nature of the reactants; in the bronsted-lowry sense, this can either gain or lose a proton, making it amphiprotic as well

Acid

A species that donates hydrogen ions or accepts electrons

Arrhenius Base

A species that donates hydroxide ions (OH-) in aqueous solution; will dissociate to form an excess of OH- in solution; this generally deals with aqueous bases; easy to identify as bases contain OH at the end of their formula (NaOH, Ca(OH)2, Fe(OH)3, and so on); this definition of acid and bases is very restrictive

Base

A species that donates hydroxide ions or electron pairs or that accepts protons

Arrhenius Acid

A species that donates protons (H+) in aqueous solution; will dissociate to form an excess of H+ in solution; this generally deals with aqueous acids; easy to identify as acids contain H at the beginning of their formula (HCl, HNO3, H2SO4, and so on); this definition of acid and bases is very restrictive

Charging

A state of an electrochemical cell in which an external electromotive force is being used to return a cell to its original state; during this process, electrons are transferred non-spontaneously from cathode to anode

Element

A substance that cannot be further broken down by chemical means; defined by its number of protons (atomic number)

Group

A vertical column of the periodic table containing elements that are similar in their chemical properties; also called a family

Electromagnetic radiation

A wave composed of electric and magnetic fields oscillating perpendicular to each other and to the direction of propagation

Critical temperature

Also known as the critical point. The highest temperature at which the liquid and gas phases of a substance can coexist; above this temperature, the liquid and gas phases are indistinguishable

Ideal bond angle

An angle between nonbonding or bonding electron pairs that minimized the repulsion between them; for example CH4 (methane), NH3 (ammonia) and H20 (water) all have tetrahedral electronic geometry; this is usually associated with an ideal bond angle of 109.5 degrees however nonbonding pairs are able to exert more repulsion than bonding pairs because these electrons reside closer to the nucleus, thus the angle in NH3 is closer to 107 degrees and the angle is water is 104.5 degrees

Calorimeter

An apparatus used to measure the heat absorbed or released by a reaction

Acidic solution

An aqueous solution that contains more H+ ions than OH- ions; pH <7 under standard conditions

Excited state

An electronic state having a higher energy than the ground state; typically attained by the absorption of a photon of a certain energy; at least one electron has moved to a sub shell of higher than normal; energy electrons can be excited to new orbital levels by heat or other energy forms to yield excited states; mnemonic: AHED a= absorb light, h= higher potential, e= excited, d= distant (from the nucleus); because the lifetime of an excited state is brief the electrons will return rapidly to the ground state, resulting in the emission of discrete amounts of energy in the forms of photons

Balanced equation

An equation for a chemical reaction in which the number of atoms for each element in the reaction and the total charge are the same for the reactant and the products

Energy density

An equivalence unit regarding the amount of electrochemical energy capable of being stored per unit weight; a battery with a large energy density can produce a large amount of energy with a small amount of material

Anion

An ionic species with negative charge; a- mean without which is a negative thing

Disproportionation

An oxidation-reduction reaction in which the same species acts as the oxidizing agent and as the reducing agent; also called dismutation; an example of this is oxygen in hydrogen peroxide

Freezing point

At a given pressure, the temperature at which the solid and liquid phases of a substance coexist in equilibrium; identical to the melting point

acetate

C2H3O2-

ethanol

C2H5OH

acetone

C3H6O

Methanol

CH3OH

cyanide

CN-

dichromate

Cr2O7 2-

chromate

CrO4 2-

Atomic Orbital

Describes the region of space where there is a high probability of finding an electron

Energy of electron transition (Bohr model)

E = -Rh [ (1/(ni ^2 ) - (1/nf ^2)]; this equation is stating the energy of the emitted photon corresponds to the difference in energy between the higher energy initial state and the lower energy final state; thus positive E corresponds to emission while negative E corresponds to absorption

Energy of an electron (Bohr model)

E = -Rh/n^2; E is energy of electron; Rh is the experimentally determined Rydberg unit of energy equal to 2.18x10^-18 J/electron; similar to angular momentum the energy of the electron changes in discrete amounts with respect to the quantum number; energy is directly proportional to the principal quantum number, this is due to the negative sign which causes the values to approach zero from a more negative value as n increases (thereby increasing the energy); ultimately the only thing the energy equation is saying is that the energy of an electron increases (becomes less negative) the farther out from the nucleus that it is located (increasing n)

Planck's relation (wavelength)

E= hc/wavelength; E= electromagnetic energy; h= Plancks constant 6.626x10^-34 J*s; c is the speed of light in a vacuum 3 x10^8 m/s; wavelength (upside down y) is wavelength of radiation; this equation is a combination of E=hf and c=fwavelength

Planck's relation (frequency)

E= hf; E is energy; h is the proportionality constant also known as planks constant (h=6.626 x 10^-34 J*s); f is frequency of radiation

Half-reaction

Either the reduction half or oxidation half of an oxidation-reduction; in an electrochemical cell, each half-reaction occurs at one of the electrodes

Bonding Electrons

Electrons located in the valence shell of an atom and involved in a covalent bond

acid equivalent

Equal to one mole of H+ or H3O+ ions

Henderson-Hasselbalch equation

Equation showing the relationship of the pH or pOH of a solution to the pKa or pKb and the ratio of the concentrations of the dissociated species; where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid; note that when [conjugate base] = [weak acid], the pH = pKa because log (1) = 0; this occurs at the half-equivalence points in a titration, and buffering capacity is optimal at this pH for weak base buffer solutions: pOH = pKb + log [B+]/[BOH] where [B+] is the concentration of conjugate acid and [BOH] is the concentration of the weak base; similar to acid buffers, pOH = pKb when [conjugate acid] = [weak base]; buffering capacity is optimal at this point

Excess reagent

In a chemical reaction, any reagent that does not limit the amount of product that can be formed

Galvanization

In electrochemical cells, the precipitation process onto the cathode itself; also called plating

Deposition

In most chemical processes, the direct transition of a substance from the gaseous state to the solid state; in electrochemical reactions, the build up of a solid precipitate onto an electrode; in organic chemistry labs a device known as a cold finger may be used to purify a product that is heated under reduced pressure, causing it to sublime, the desired product is usually more volatile than the impurities, so the gas is purer than the original product and the impurities are left in the solid state; the gas then deposits onto the cold finger, which ahas cold water flowing through it, yielding a purified solid product that can be collected

permanganate

MnO4-

Delocalized orbitals

Molecular orbitals in which electron density is spread over an entire molecule, or a proton thereof, rather than ebbing localized between two atoms

ammonium

NH4+, charge 1+

Heterogenous

Nonuniform in composition

Balmer series

Part of the emission spectrum for hydrogen, representing transitions of an electron from energy levels n>2 to n=2

Compression

Reduction in the volume of a gas

thiocyanate

SCN-

d subshell

Subshell corresponding to the angular momentum quantum number l=2; contains five orbital sand is found in the third and higher principal energy levels

Hydronium ion

The H3O+ ion

Hydroxide ion

The OH- ion

Intermolecular forces

The attractive and repulsive forces between molecules; these interactions can impact certain physical properties such as melting and boiling points; the weakest of the intermolecular interactions are the dispersion forces also known as London forces; the next is dipole-dipole interactions which are intermediate strength, finally the strongest is hydrogen bond which is a misnomer because there is no actual sharing or transfer of electrons; hydrogen bonding is the strongest intermolecular interaction however it is only about 10% of the strength of a covalent bond, therefore these interactions can be overcome with small or moderate amounts of energy

Bond Length

The average distance between two nuclei in a bond; as the number of shared electron pairs increases, the bond length decreases

Hybridization

The combination of two or more atomic orbitals to form new orbitals with properties that are intermediate between those of the original orbitals

Aufbau principle

The concepts that electrons fill energy levels in order of increasing energy, completely filling one sub level before beginning to fill the next

Electron affinity

The energy dissipated by a gaseous species when it gains an electron; halogens are the most "greedy" group of elements due to them being able to complete its octet and achieve a noble gas configuration, this exothermic process expels energy in the form of heat; of note the electron affinity is essentially the opposite concept from ionization energy; because this is exothermic it has a negative deltaHrxn however the electron affinity is reported as a positive number, this is because electron affinity refers to the energy dissipated: if 200 kJ/mol of energy is released, deltaHrxn =-200 kJ/mol, and the electron affinity is 200 kJ/mol; the stronger the electrostatic pull (the higher the zeff) between the nucleus and the valence shell electrons, the greater the energy released will be when the atom gains the electron, thus electron affinity increased across the period from left to right; because the valence shell is farther away from the nucleus as the principal quantum number increased, electron affinity decreased in a group from top to bottom; most metals also have low electron affinity values

Ideal gas law

The equation stating PV=nRT, where R is the gas constant, 8.21 x 10^-2 L*atm/mol*K or 8.314 J/K*mol; can be used to describe the behavior of many real gases at moderate pressures and temperatures significantly above absolute zero

Actual Yield

The experimental quantity of a substance obtained at the end of a reaction

Chemical bond

The interaction between two atoms resulting from the sharing or transfer of electrons

Electron spin

The intrinsic angular momentum of an electron, represented by ms; has arbitrary values of +1/2 and -1/2

Boyle's Law

The law stating that at constant temperature, the volume of a gaseous sample is inversely proportional to its pressure PV=k or P1V1=P2V2; this law shows that pressure and volume are inversely related, when one increases the other decreases

Gay-Lussac's Law

The law stating that the pressure of a gaseous sample at constant volume is directly proportional to its absolute temperature; P/T = k or P1/T1 = P2/T2

Graham's Law

The law stating that the rate of effusion or diffusion for a gas is inversely proportional to the square root of the gas's molar mass; this theory predicts that heavier gases diffuse more slowly than lighter ones because of their differing average speeds; because all gas particles have the same average kinetic energy at the same temperature, it must be true that particles with greater mass travel at a slower average speed r1/r2= square root of (M2/M1); where r1 and r2 are the diffusion rates of gas 1 and 2 and M1 and M2 are the molar masses of gases; from the equation we can see that a gas that has a molar mass four times that of another gas will travel half as fast as the lighter gas; of note diffusion is when gases mix with one another and effusion is when a gas moves through a small hole under pressure, both will be slower for larger molecules, both conditions use the same equation When trying to find the r2 rate you can remember to convert the equation to r2 = r1 square root M1/M2

First Law of Thermodynamics

The law stating that the total energy of a system and its surroundings remains constant; Delta U = Q - W, where delta U is the change in internal energy of the system, Q is the heat added to the system, and W is the work done by the system

Avogadro's Principle

The law stating that under the same conditions of temperature and pressure, equal volumes of different gases will have the same number of molecules; equal amounts of all gases at the same temperature and pressure will occupy equal volumes; of note one mole of any gas at STP will occupy 22.4 L; equation is n/V = k or n1/V1 = n2/V2 where k is a constant, n1 and n2 are the number of moles of gas 1 and gas 2, and v1 and v2 are the volumes of thr gas; this is summarized as the number of moles of a gad increases the volume increases in direct proportion

Atomic mass

The mass of a given isotope of an element; closely related to the mass number

Bohr Model

The model of the hydrogen atom in which electrons assume certain circular orbits around positive nucleus; this model of the hydrogen atom explained the atomic emission spectrum of hydrogen, which is the simplest emission spectrum among all the elements

Atomic number

The number of protons in a given element; top number above element on periodic table

Bond order

The number of shared electron pairs between two atoms; a single bond has a bond order of 1, a double bond has a bond order of 2, a triple bond has a bond order of 3

Critical Point

The point in a phase diagram beyond which the phase boundary between liquid and bas no longer exists

Angular momentum

The rotational analog of linear momentum; L= nh/2pi; n is the principal quantum number which can be any positive integer; h is Planks Constant 6.626x10^-34 J*s; because the only variable is the principal quantum number, the angular momentum of an electron changes in discrete amounts with respect to the principal quantum number; this is part of Bohr's model

Half-cell

The separated compartments housing the electrodes and solutions in an electrochemical reaction

Dissociation

The separation of a single species into two separate species; usually used in reference to salts or weak acids or bases

Equilibrium

The state of balance in which the forward and reverse reaction rates of a reversible reaction are equal; the concentrations of all species will remain constant over time unless there is a change in the reaction conditions

Critical pressure

The vapor pressure at the critical temperature of a given substance

Atomic weight

The weighted average mass of the atoms of an element, taking into account the relative abundance of all naturally occurring isotopes; bottom number on periodic table

Chemical Properties

Those properties of a substance related to the chemical changes that it undergoes, such as ionization energy and electronegativity

Colligative properties

Those properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolved particles; these properties - vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure - are usually associated with dilute solutions

Homogeneous

Uniform in composition

Pi bond

a bond with two parallel electron cloud densities formed between two p-orbitals that limits the possibility of free rotation; pi bonds are the second bond in a double bond and both the second and third bonds in a triple bond

Resonance hybrid

a lewis structure that represents the weighted average (by stability) of all possible resonance structures; the more stable resonance structure is the major contributor to the resonance hybrid, in general the more stable the structure the more it contributes to the character of the resonance hybrid

Rate-law

a mathematical expression giving the rate of a reaction as a function of the concentrations of the reactants; must be determined experimentally; rate = k[A]^x[B]^y where k is the reaction rate coefficient or rate constant and the exponents x and y are the order of the reaction

solubility

a measure of the amount of solute that can be dissolved in a solvent at a certain temperature; when the maximum amount of solute has been added, the dissolved solute is in equilibrium with its undissolved state, and we say that the solution is saturated; if more solute is added it will not dissolve; the excess solid form will precipitate at the bottom of the container; a solution in which the proportion of solute to solvent is small is said to be dilute, and one in which the proportion is large is said to be concentrated, note that both dilute and concentrated solutions are still considered unsaturated if the maximum equilibrium concentration (saturation) has not been reached

Titration

a method used to determine the concentration of an unknown solution by gradual reaction with a solution of known concentration; these are performed by adding small volumes of a solution of known concentration (the titrant) to a known volume of a solution of unknown concentration (the titrand) until completion of the reaction is achieved at the equivalence point

Reduction

a reaction involving the net gain of electrons, decreasing the oxidation number

Oxidation

a reaction involving the net loss of electrons, increasing oxidation number

Open System

a system that can exchange both energy (heat and work) and matter with its surroundings; for example a pot of boiling water

Isotopes

atoms containing the same number of protons but different numbers of neutrons; same atomic number but different mass numbers

van der Waals forces

attractive or repulsive forces between molecules that don't arise from covalent or ionic bonds

Pressure

average force per unit area measured in atmospheres (atm), torr or mmHg, or pascals (Pa); 1 atm = 760 torr = 760 mmHg = 101.235 kPa

standard conditions

conditions defined as 25 deg C, 1 atm pressure, and 1 M concentrations; used for measuring the standard Gibbs free energy, enthalpy, entropy, and cell electromotive force; this is not standard temperature and pressure; for this the most stable form of a substance is called the standard state of that substance

standard temperature and pressure (STP)

defined as 0 deg C (273 K) and 1 atm; used for measuring characteristic of an ideal gas; remember that STP is different from standard state

polyvalent acid or base

each mole of the acid or base liberates more than one acid or base equivalent

octet

eight valence electrons in a subshell around a nucleus; imparts great stability to an atom

nonbonding electrons

electrons located in the valence shell of an atom but not involved in covalent bonds

base equivalent

equal to one mole of OH- ions

Reaction Quotient (Q)

has the same form as the equilibrium constant, but the concentrations of products and reactant may not be at equilibrium; when compared to Keq, it dictates the direction a reaction will proceed spontaneously; equation is Qc = ([C]^c[D]^d)/([A]^a[B]^b); this equation is similar to the Keq however it provided different information, while the concentrations used for the law of mass action are equilibrium (constant) concentrations, the concentrations of the reactants and products are not constant when calculating a value for Q of a reaction, thus the utility of Q is not the value itself but rather the comparison that can be made between Q at any given moment in the reaction to the known Keq for the reaction at a particular temperature. Key concept: Q < Keq: delta G < 0, reaction proceeds in forward direction; Q = Keq, delta G = 0, reaction is in dynamic equilibrium, Q > Keq, delta G > 0, reaction proceeds in the reverse reaction

Quanta

in Placnk's theory, discrete bundles of energy that are emitted as electromagnetic radiation from matter

Reaction order

in a calculation of the rate law for a reaction, the sum of the exponents to which the concentrations of reactants must be raised

stoichiometric coefficient

in a reaction, the number placed in front of each compound to indicated relative number of moles of that species involved in the reaction; for example the balanced equation expressing the combustion of nonane is: C9H20 (g) + 14 O2 (g) --> 9CO2 (g) + 10 H2O (l); the coefficients indicate that one mole of C9H20 gas must be reacted with fourteen moles of O2 gas to produce nine moles of carbon dioxide and ten moles of water

Process

in a system, when a change in one or more of the properties (such as concentrations of reactants or products, temperature, or pressure) of the system occurs; while a process is associated with a change of the state of a system some are uniquely identified by the property that is constant throughout the process, such as isothermal processes which occur when the systems temperature is constant

Reducing agent

in an oxidation-reduction reaction, the atom that facilitates the reduction of another species; the reducing agent loses electrons and is thereby oxidized; common reducing agents include: CO, C, B2H6, Sn 2+ and other pure metals, hydrazine, Zn(Hg), Lindlar's catalyst, NaBH4, LiAlH4, NAHD, FADH2 of note NAD+ can act as both oxidizing and reducing agent at different times in metabolic pathways (as can many other biochemical redox reagents), as such they act as mediators of energy transfer during many metabolic processes

Microstate

in thermodynamics, a specific way in which energy of a system is organized; the availability of energy microstates increases as the temperature of the solid increases; this means that the molecules have greater freedom of movement, and energy disperses (in simple terms they are freer to move around in different ways)

spectator ions

ions involved in a reaction that do not change formula, charge, or phase; normally omitted from the net ionic equation

cell diagram

is a shorthand notation representing the reactions in an electrochemical cell ex. for Daniell Cell: Zn (s) | Zn 2+ (1M) || Cu 2+ (1M) | Cu (s) the following rules are used in constructing a cell diagram: 1. the reactants and products are always listed from left to right in this form: anode | anode solution (concentration) || cathode solution (concentration) | cathode 2. A single vertical line indicates a phase boundary 3. A double vertical line indicates the presence of salt bride or some other type of barrier

van der Waals equation of state

one of several real gas laws, which corrects for attractive forces and the volume of gas particles, which are assumed to be negligible in the ideal gas law (P + n^2a/V^2) (V - nb) = nRT where a and b are physical constants experimentally determined for each gas, the a term corrects for the attractive forces between molecules and as such will be smaller for gases that are small and less polarized (such as helium), larger for gases that are larger and more polarized (such as Xe and N2) and largest for polarized molecules such as HCl and NH3; the b term corrects for the volume of the molecules themselves, larger molecules thus have larger values of b, numerical values for a are generally much larger than those for b; note that if both a and b are zero the van der waals equation of a state reduces to the ideal gas law; a mnemonic to remember a and b is a is the van der waals term for the attractive forces and b is the van der waals term for big particles

Phase

one of the three forms of matter; solid, liquid, or gas; also called state

fusion

or melting; the transition from solid to liquid

paschen series

part of the emission spectrum for hydrogen, representing transitions of an electron from energy levels n >= 4 to n=3

saturation point

solute concentration is at its maximum value for the given temperature and pressure; this is defined as the equilibrium position, which is the lowest energy state of a system under a given set of temperature and pressure conditions; systems move spontaneously towards equilibrium and any movement away from equilibrium is non-spontaneous; as the solution becomes more concentrated and approaches saturation the rate of dissolution lessens and the rate of precipitation increases; neither dissolution nor precipitation are more thermodynamically favored at equilibrium; at this point, like all systems at equilibrium, the free energy is zero

yield

the amount of product obtained from a reaction

spectrum

the characteristic wavelengths of electromagnetic radiation emitted or absorbed by an object, atom, or molecule

solvent

the component of a solution present in the greatest amount; the substance in which the solute is dissolved

Solute

the component of a solution that is present in a lesser concentration that the solvent

Vapor pressure depression

the decrease in vapor pressure of a liquid caused b the presence of dissolved solute; a colligative property

Maxwell-Boltzmann distribution curve

the distribution of the molecular speeds of gas particles at a given temperature; as temperature increases, average speed increases and the distribution becomes wider and flatter

subshell

the division of electron shells or energy levels into different values of the azimuthal quantum number (s,p,d, and f); composed of orbitals

standard hydrogen electrode (SHE)

the electrode defined as having a potential of zero under standard conditions; all oxidation and reduction potentials are measured relative to the standard hydrogen electrode at 25 deg C and with 1 M concentrations of each ion in solution

paired electrons

the electrons in the same orbital with assigned spins of +1/2 and -1/2

Latent heat

the enthalpy of an isothermal process

solubility product constant (Ksp)

the equilibrium constant for the ionization reaction of a sparingly soluble salt; key concept: remember that Ksp is just a specialized form of Keq, so you can simplify a lot of problems by using the same concepts that you do for all equilibria equation: Ksp = [A^n+]^m [B^m-]^n for example: you can express the Ksp of silver chloride as Ksp = [Ag+][Cl] remember that pure solids and liquids do not appear in the equilibrium constant

rate order

the exponential effect of a change in concentration of a reactant on the change of rate in a reaction; the overall rate order is the sum of all the individual reactant rate orders

Periodic law

the law stating that the chemical properties of elements depend on the atomic number of the element and change in a period fashion

molar solubility

the molarity of a solute in a saturated solution

acid-base nomenclature

the names of most acids are related to the names of their parent anions (the anion that combines with H+ to form the acid); acids formed from anions with names that end in -ide have the prefix hydro- and the ending -ic example: F- = fluoride ; HF = hydrofluoric acid Cl- = chloride; HCl = hydrochloric acid Br- = bromide; HBr = hydrobromic acid acids formed form oxyanions are called oxyacids; if the anion ends in -ite (less oxygen then the acid will end with -ous acid; if the anion ends in -ate (more oxygen), then the acid will end with -ic acid; prefixes in the names of the anions are retained

valence shell

the outermost shell of an atom

Vapor pressure

the partial pressure of a gaseous substance in the atmosphere above the liquid or solid with which it is in equilibrium

percent yield

the percentage of the theoretical product yield that is actually recovered when a chemical reaction occurs; obtained by dividing the actual yield by the theoretical yield and multiplying by 100%

Standard state

the phase of matter for a certain element under standard conditions; for example, H2 (g), H2O (l), NaCl (s), O2 (g), and C (s, graphite) are the most stable forms of these substances under standard conditions

spectroscopic notation

the shorthand representation of the principal and azimuthal quantum numbers, in which the azimuthal number is designated by a letter rather than a number; l=0 is s, l=1 is p, l=2 is d, l=3 is f; example an electron in the shell n=4 and subshell 2 is said to be in 4d

molecular geometry

the spatial arrangement of only the bonding pairs of electrons around a central atom; example NH3 (ammonia) has tetrahedral electronic structure but has trigonal pyramidal molecular structure; the presence of bond dipoles does not necessarily results in a molecular dipole; that is an overall separation of charge across the molecule; we must fist consider the molecule geometry and vector addition of the bond dipoles based upon that molecular geometry; a compound with nonpolar bonds is always nonpolar however a compound with polar bonds may be polar or nonpolar; depending upon the spatial orientation of the polar bonds in the molecule; if the compound has a molecular geometry such that the bond dipole moments cancel each other out (that is, if the vector sum is zero) then the result is a nonpolar compound; for example CCl4 (carbon tetrachloride) has four polar C-Cl bonds, but because the molecular geometry of carbon tetrachloride is tetrahedral, the four bond dipoles point to the vertices of the tetrahedron and therefore cancel each other out, resulting in a nonpolar compound

Reaction Rate

the speed at which a substance is produced or consumed by a reaction

Liquid

the state of matter in which intermolecular attractions are intermediate between those in gases and in solids, distinguished from the gas phase by having a definite volume and from the solid phase because molecules may mix freely

molecular weight

the sum of the atomic weights of all the atoms in a molecule

Standard Potential

the voltage associated with a half reaction of a specific oxidation-reduction reaction; generally tabulated as reduction potentials, compared to the standard hydrogen electrode

hypo- and per-

this and hyper, written as per- are used to indicate less oxygen and more oxygen respectively; ClO- is Hypochlorite, ClO2 - is chloride, ClO3 - is chlorate, and ClO4 - is perchlorate

Inert

unreactive

valence electrons

valence electrons of an atom are those electrons that are in its outermost energy shell, are most easily removed, and are available for boning; in other words the valence electrons are the "active" electrons of an atom and to a large extent dominate the chemical behavior of the atom. For elements in Groups IA and IIA only the highest s subshell electrons are valence electrons. For elements in Groups IIIA through VIIIA the highest s and p subshells electrons are valence electrons. For transition elements, the valance electrons are those in the highest s and d subshells, even though they have different principal quantum numbers. For the lanthanide and actinide series, the valance electrons are those in the highest s and f subshells, even though they have different principal quantum numbers. All elements in period three (starting with sodium) and below may accept electrons into their d subshell, which allows them to hold more than eight electrons in their valence shell. This allows them to violate the octet rule; example how electrons are the valence electrons of the elemental vanadium, elemental selenium, and the sulfur atom in a sulfate ion? vanadium ahs five valence electrons: two in its 4s subshell and three in its 3d subshell; selenium has 6 valence electrons, 2 in its 4s subshell and four in its 4p subshell, selenium's 3d electrons are not part of its valence electrons; sulfur in a sulfate ion has 12 valence electrons, its original 6 plus 6 from the oxygens to which it is bonded, sulfur's 3s and 3p subshells can contain only eight of these 12 electrons; the other four electrons have entered the sulfur atom's 3d subshells, which is normally empty in elemental sulfur

heating curve

when a compound is heated, the temperature rises until the melting or boiling point is reached ;then, the temperature remains constant as the compound is converted to the next phase (liquid or gas, respectively); once the entire sample is converted then the temperature begins to rise again; during phase change reactions we cannot use q=mcdeltaT because delta T = 0 due to no changes in temperature; we must use enthalpy (or heat) of fusion (delta Hfus) to determine the heat transferred during the phase change; when transitioning from solid to liquid the change in enthalpy will be positive because heat must be added; when transitioning from a liquid to a solid, the change in enthalpy will be negative because heat must be removed; at the liquid-gas boundary, the enthalpy (or heat) of vaporization (delta Hvap) must be used, and its sign convention also follows a similar pattern; use equation q=mL where m is the mass and L is the latent heat; latent heat is the general term for enthalpy of an isothermal process given in units cal/g.

-ic

when this is added to the root of the latin name fo the element it represents the ion with greater charge; example Fe3+ is ferric and Cu2+ is Cupric

-ous

when this is added to the root of the latin name of the element it represents the ion with lesser charge; example Fe2+ is ferrous and Cu+ is cuprous

Anode

The electrode at which oxidation occurs

Cathode

The electrode at which reduction takes places

Bond Energy

The energy (enthalpy change) required to break a particular bond under given conditions

surroundings

all matter and energy in the universe not included in the particular system under consideration

standard free energy (G^o)

the Gibbs free energy for a reaction under standard conditions

pascal (Pa)

the SI unit for pressure, equivalent to N/m^2; 1 atm = 101,325 Pa

Intramolecular forces

the attractive forces between atoms within a single molecule (ionic and covalent bonds

spin quantum number (ms)

the fourth quantum number, which indicates the orientation of the intrinsic spin of an electron in an atom; can only assume values of +1/2 and -1/2; when electrons are in the same orbital they must have opposite spins, in this case they are often referred to as being paired; electrons in different orbitals with the same ms values are said to have parallel spins

Ion product (IP)

the general term for the reaction quotient of a dissolving ionic compound; compared to Ksp to determine the saturation status of a solution key concept: IP < Ksp: unsaturated, solute will continue to dissolve IP = Ksp: saturated, solution is at equilibrium IP > Ksp: supersaturated, precipitation will occur

theoretical yield

the maximal amount of product that can be obtained in a reaction; determined by stoichiometric analysis of the limiting reagent

Pauling electronegativity scale

the most common scale used to express electronegativity of the elements

Heterogeneous catalyst

A catalyst that is not in the same phase of matter as the reactants (for example, a solid platinum catalyst reacting with hydrogen gas)

Covalent Bond

A chemical bond formed by the sharing of an electron pair between two atoms; can be in the form of single bonds, double bonds, or triple bonds; this type of bond is usually nonmetals that have relatively similar values of electronegativity; this is due to the energy required to form ions through the complete transfer of one or more electrons is greater than the energy that would be released upon the formation of an ionic bond; the degree to which the pair of electrons is shared equally or unequally between the two atoms determines the degree of polarity in the covalent bond; for example if the electron pair is shared equally the covalent bond is non-polar; if the pair is shared unequally the bond is polar; if both of the shared electrons are contributed by only one of the two atoms, the bond is called coordinate covalent; covalent compounds contain discrete molecular units with relatively weak intermolecular interactions, as a result compounds like carbon dioxide (CO2) tend to have lower melting and boiling points; in addition, because they do not break down into constituent ions, they are poor conductors of electricity in the liquid state or in aqueous solutions; bond length decreased the more bonds that are shared, hence triple bonds are the shortest, double are second shortest and single bonds are the longest; boney energy is the energy required to break a bond by separating its components into their isolated, gaseous atomic states, triple bonds are the strongest and single bonds are the weakest; polarity occurs when two atoms have a relative difference in electronegativity; when these atoms come together in covalent bonds they must negotiate the degree to which the electron pairs will be shared; the atom with the higher electronegativity gets the larger share fo electron density; a polar bond creates a dipole, with the positive end fo the dipole at the less electronegative atom and the negative end at the more electronegative atom, this creates partial negative charge and partial positive charge, example HCl, H becomes positive and Cl negative, this is represented by an arrow with a cross on the positive end pointing to the negative end; the dipole moment fo the polar bond or polar molecule is a vector quantity given by the equation p=qd where p is the dipole moment, q is the magnitude of charge, and d is the displacement vector separating the two partial charges; when atoms that have identical or nearly identical electronegative share electron pairs they do so with equal distribution of electrons, this is called non-polar covalent bonds, there is no separation of charge across the bond; the most common diatomic molecules are H2, N2, O2, F2, Cl2, Br2, and I2, an easy way to remember this is they form a 7 on the periodic table (except for H) and there are 7 of them ; any bond between atoms with difference in electronegativity less than 0.5 is generally considered non-polar; polar are generally 0.5-1.7;

Arrhenius equation

A chemical kinetics equation that relates the rate constant (k) of a reaction with the frequency factor (A), the activation energy (Ea), the ideal gas constant (R), and temperature (T) in kelvin. k= Ae^(-Ea/RT)

Broken-order reaction

A reaction with noninteger orders in its rate last

Emission Spectrum

A series of discrete lines at characteristic frequencies, each representing the energy emitted when electrons in an atom return from an excited state to their ground state

Common ion effect

A shift in equilibrium of a solution due to the addition of ions of a species already present in the reaction mixture; the solubility of a salt is considerably reduced when it is dissolved in a solution that already contains one of its constituent ions however the presence of the common ion has no effect on the value of the solubility product constant itself; for example a salt such as CaF2 is dissolved into water already containing Ca 2+ ions (from some other salt, perhaps CaCl2) the solution will dissolve less CaF2 than would an equal amount of pure water; this effect is La Chatelier's principle in action due to the solutions already containing one of the constituent ions form the product side of the dissociation equilibrium, the system will shift toward the left side, reforming the solid salt, as a result molar solubility for the solid is reduced, and less of the solid dissolves in the solution although Ksp remains constant

Crystal

A solid in which atoms, ions, or molecules are arranged in a regular, three-dimensional lattice structure

Buffer

A solution containing a weak acid and its salt (which is composed of its conjugate base and a cation) or weak base and its salt (which is composed of its conjugate acid and an anion) that tends to resist changes in pH when small amounts of acid or base are added; the most important buffer in the human body is the H2CO3/HCO3- conjugate pair int he plasma component of the blood, called the bicarbonate buffer system; this is tied to the respiratory system, in conditions of metabolic acidosis (production of excess plasma H+ not caused by the respiratory system itself) the breathing rate will increase to compensate and blow off a greater amount of carbon dioxide gas; this causes the system to shift to the left, thereby reducing [H+] and buffering against dramatic and dangerous changes to the blood pH; it is interesting to note that the bicarbonate buffer system (pKa = 6.37) maintains a pH around 7.4 which is actually slightly outside the optimal buffering capacity of the system; buffers have a narrow range of optimal activity (pKa +- 1); this makes sense as it is far more common for acidemia (too much acid in the blood) to occur than alkalemia (too much base in the blood); as acidemia becomes more severe, the buffering system actually becomes more effective and more resistant to further lowering the pH

Ideal Solution

A solution with an enthalpy of dissolution that is equal to zero; this occurs when the overall strength of the new interactions is approximately equal to the overall strength of the original interaction; in this case the overall enthalpy change for the dissolution is close to zero

Dipole

A species containing bonds between elements of different electronegativities, resulting in an unequal distribution of charge

Indicator

A substance used in low concentrations during a titration that changes color over a certain pH range (acid-base titrations) or at a particular electromotive force (oxidation-reduction titrations); the final color change of an indicator occurs at the endpoint of a titration; indicators are weak organic acids or bases that have different colors in their protonated and deprotonated states; this small structural change, the binding or release of a proton - leads to a change in the absorption spectrum of the molecule, which we perceive as a color change; indicators are generally vibrant and can be used in low concentrations without significantly altering the equivalence point; the indicator must always be weaker than the acid or base being titrated otherwise the indicator would be titrated first; the endpoint when the indicator changes color is not the equivalence point; the most useful titrations combinations involve at least one strong species, weak acid/weak base titrations can be done but are not very accurate and therefore are rarely performed

Closed System

A system that can exchange energy (heat and work) but not matter with its surroundings; for example a steam radiator

Isoelectric Focusing

A technique used to separate amino acids or polypeptides based on their isoelectric points

Electronic geometry

The spacial arrangement of all pairs of electrons around a central atoms, including both the bonding and long pairs; example NH3 (ammonia) has tetrahedral electronic structure but has trigonal pyramidal molecular structure

Empirical formula

The simplest whole-number ratio of the different elements in a compound; for example the empirical formula for benzene is CH, white the molecular formula is C6H6; for some compounds the empirical and molecular formula are identical as is the case for H2O; as previously discussed ionic compounds such as NaCl or CaCO3 will only have empirical formula

Avogadro's Number

The number of atoms or molecules in one mole of a substance: 6.02 X 10^23 mol^-1

Actinide series

The series of chemical elements atomic numbered 89-103 and falling between the s and d blocks on the period table

Absorption spectrum

The series of discrete lines at characteristic frequencies representing the energy required to excite an electron from the ground state

Reaction Mechanism

The series of steps that occur in the course of a chemical reaction, often including the formation and destruction of reaction intermediates

Atom

The smallest unit of an element that retains the properties of the element; it cannot be further broken down by chemical means

Electron Shell

The space occupied by/path followed by an electron around an atom's nucleus. Electron shell (also called principle energy level) for a given electron is indicated by its principle quantum number

metal

one of a class of elements on the left side of the periodic table possessing low ionization energies and electronegativities; readily give up electrons to form cations and possess relatively high electrical conductivity; these include active metals, transition metals, and lanthanide and actinide series of elements; metals are lustrous (shiny) solids, except for mercury, which is a liquid under standard conditions; they generally have high melting points and densities, but there are exceptions, such as lithium, which has a density about half that of water; metals have the ability to be deformed without breaking; the ability of metal to be hammered into shapes is called malleability, and its ability to be pulled or drawn into wires is called ductility; at the atomic level a metal is defined by a low effective nuclear charge, low electronegativity (high electropositivity) large atomic radius, small ionic radius, low ionization energy and low electron affinity; all of these characteristics are manifestations of the ability of metals to easily give up electrons

nonmetal

one of a class of elements with high ionization energies and electron affinities that generally gain electrons to form anions; located in the upper right corner of the period table; nonmetals are generally brittle in the solid state and show little or no metallic luster; they have high ionization energies, electron affinities, and electronegativities, as well as small atomic radii and large ionic radii; they are usually poor conductors of heat and electricity; all of these characteristics are manifestations of the inability of nonmetals to easily give up electrons; these are less unified in their chemical physical properties than the metals; ex carbon is a stereotypical nonmetals that retains a solid structure but is brittle, non lustrous, and generally a poor conductor of heat and electricity

s subshell

subshell corresponding to the angular momentum quantum number l=0 and containing one spherical orbital; found in all energy levels

Solid

the phase of matter possessing the greatest order; molecules are fixed in a rigid structure

nucleus

the small central region of an atom; a dense, positively charged area containing protons and neutrons

molecule

the smallest polyatomic unit of an element or compound that exits with distinct chemical and physical properties

Concentration cell

A cell that creates an electromotive force (emf or voltage) using a single chemical species in half-cells of varying concentration

Electrochemical cell

A cell within which an oxidation-reduciton reaction takes place, containing two electrodes between which there is an electrical potential difference; there are three fundamental types of electrochemical cells: galvanic cells (also known as voltaic cells), electrolytic cells, and concentration cells; galvanic cells and concentration cells are spontaneous reactions, whereas electrolytic cells contain non-spontaneous reactions; all three types contain electrodes where oxidation and reduction take place; for all electrochemical cells, the electrode where oxidation occurs is called the anode, the electrode where reduction occurs is called the cathode (mnemonic: electrodes in a electrochemical cell: AN OX and a RED CAT; the ANode is the site of OXidation; REDuction occurs at the CAThode); for all electrochemical cells, the movement of electrons is from anode to cathode, and the current (I) runs from cathode to anode (electrons move through an electrochemical cell opposite to the flow of current (I)); it is also important to note that all batteries are influenced by temperature changes, for instance lead-acid batteries in cars, like most galvanic cells, tend to fail most in cold weather; of note the electron flow in this cell is A to C (order of alphabet), electrons flow from anode to cathode in all types of electrochemical cells

Ionic Bond

A chemical bond formed through electrostatic interaction between positive and negative ions; one or more electrons from an atom with low ionization energy, typically a metal, are transferred to an atom with a high electron affinity, typically a nonmetal; an example of this is sodium chloride; the resulting electrostatic attraction between opposite charges is what holds the ions together; of note this type of attraction creates lattice structures consisting of repeating rows of cations and anions, rather than individual molecular bonds; the atom that loses the electrons becomes a cation, and the atom that gains electrons becomes an anion; an easy mnemonic is meTals lose electrons to become caTions = posiTive (+) ions and Nonmetals gains electrons to become aNions = Negative (-) ions; these electrons are not shared in ionic bonds, for the electron transfer to occur the difference in electronegativity must be greater than 1.7 on the Pauling scale; because of the strength of the electrostatic force between the ionic constituents of the compound, ionic compounds have a very high melting and boiling point; for example the melting point of sodium chloride is 801 degrees C; many ionic compounds dissolve readily in water and other polar solvents and in the molten or aqueous state are good conductors of electricity; in the solid state the ionic constituents of the compound form a crystalline lattice consisting of repeating positive and negative ions, with this arrangement the attractive forces between oppositely charged ions are maximized, and the repulsive forces between ions of like charge are minimized

Decomposition Reaction

A reaction in which a single compound breaks down into two or more products; usually results from heating, high frequency radiation or electrolysis; an example of this is breakdown of Mercury (II) oxide; note the delta (triangle) over the arrow represents addition of heat; 2HgO (s) -----(heat)----> 2Hg (l) + O2 (g) ; an example of a reaction that utilizes high frequency light is the decomposition of silver chloride crystals in the presence of sunlight, the ultraviolet compound of sunlight has sufficiency energy to catalyze certain chemical reactions; for silver chloride exposure to sunlight results in a decomposition reaction that yields a rust colored product that consists of separated silver and chloride

Amphiprotic species

A species that may either gain or lose a proton

Fluid

A substance that flows due to weak intermolecular attractions between molecules and that takes the shape of its container; liquids and gases are considered fluids

Catalyst

A substance that increases the rates of the forward and reverse directions of a specific reaction by lowering activation every, but is itself left unchanged; homogenous catalyst is in the same phase (solid, liquid, gas) as the reactant; heterogeneous catalyst is in a distinct phase; notice that the only effect the catalyst is the decrease in the energies of activation, Ea, for both forward and reverse reactions; the presence of the catalyst has no impact on the free energies of the reactants or the products or the difference between them; this means that catalyst change only the rates of reactions, and in face, change the forward rate and reverse rate by the same factor; consequently they have no impact whatsoever on the equilibrium position or the measurement of Keq catalyst do not make non-spontaneous reactions into spontaneous ones; they only make spontaneous reactions move more quickly toward equilibrium

Barometer

A tool for measuring pressure

Dipole moment

A vector quantity with a magnitude that is dependent on the product of the charges and the distance between them; oriented from the positive to the negative pole

borate

BO3 3-

Alkali Metals

Elements found in Group IA of the period table; highly reactive, readily losing one valence electron to form ionic compound with nonmetals; these possess most of the classical physical properties of metals except that their densities are lower than those of other metals; they have one loosely bound electron in their outermost shells; they have very low zeff values giving them the largest atomic radii of all the elements in their respective periods; this low zeff also explains the other trends of low ionization energies, low electron affinities, and low electronegative; these metals easily lose one electron to form univalent cations and they react readily with nonmetals especially the halogens as in NaCl

Buffering capacity

The degree to which a system can resist changes in pH

Dipole-dipole interactions

The attractive forces between two dipoles; magnitude is dependent on both the dipole moments and the distance between the two species; polar molecules tend to orient themselves in such a way that the oppositely charged ends of the respective molecular dipoles are closest to each other: the positive region of one molecule is close to the negative region of another molecule; dipole dipole interactions are present in the solid and liquid phases but become negligible in the gas phase because of the significantly increased distance between gas particles; polar species tend to have a higher melting and boiling point than nonpolar species of comparable molecular weight due to these interactions; London forces and dipole dipole interactions are different not in the kind but duration; both are electrostatic forces between opposite partial charges ;the difference is only in the transience or permanence of the molecular dipole

Bond Enthalpy

The average energy that is required to break a particular type of bond between atoms in the gas phase; also known as bond dissociation energies; remember bond dissociation is an endothermic process; in units kJ/mol of bonds broken and is often given in tables on the MCAT; key concept: because it takes energy to pull two atoms apart, bond breakage is generally endothermic, the reverse process, bond formation is generally exothermic; remember atoms generally form bonds to become more stable (often by completing an octet)

Hess's Law

The law stating that the energy change in an overall reaction is equal to the sum of the energy changes in the individual reactions that comprise it delta Hnet = (big E) delta Hr; this law states that enthalpy changes of reactions are additive; when thermochemical equations (chemical equations for which energy changes are known) are added to give the net equation for a reaction, the corresponding heats of reaction are also added to give the net heat of reaction

Henry's Law

The law stating that the mass of a gas that dissolves in a solution is directly proportional to the partial pressure of the gas above the solution; vapor pressure is the pressure exerted by evaporated particles above the surface of a liquid C=kP C= concentration of a dissolved gas k=Henry's Law constant P= partial pressure of the gas of note the solubility of a gas will increase with increasing partial pressure of the gas equation is [A] = ksubH * PsubA or [A]1/P1 = [A]2/P2 =ksubH where [A] is the concentration of A in solution, ksubH is Henrys constant, and PsubA is the partial pressure of A, the value of Henry's constant depends on the identity of the gas; according to this relationship solubility (concentration) and pressure are directly related

Dalton's law of partial pressures

The law stating that the sum of the partial pressures of the components of a gaseous mixture must equal the total pressure of the sample; Psub T = P1 + P2 + P3 +P4 ..... where Psub T is the total pressure in the container; the partial pressure of a gas is related to its mole fraction and can be determined using the following equation: Psub A = Xsub A * Psub T where X sub A = moles of gas A/ total moles of gas

nonelectrolyte

a compound that does not ionize in water

Polyprotic

a molecule capable of donating more than one proton;

Excitation

The promotion of an electron to a higher energy level by absorption of an energy quantum

Ductility

The property of metals that allows a material to be drawn into thinly stretched wires

Diffusion

The random motion of gas or solute particles across a concentration gradient, leading to uniform distribution of the gas or solute throughout the container

Electromagnetic spectrum

The range of all possible frequencies or wavelengths of electromagnetic radiation

Conjugate acid-base pair

The relationship between Bronsted-Lowry acid and its deprotonated form, or a Bronsted-Lowry base and its protonated form; these occur by definition that they require the transfer of a proton from the acid to the base; the conjugate acid is the acid formed when a base gains a proton, and conjugate base is the base formed when an acid loses a proton; when finding the Ka and Kb the reactions are reversible causing the net reaction to simplify to the dissociate of water which means the equilibrium constant for the reaction is Kw = [H3O+][OH-] = 10^-14 which is the product of Ka and Kb; remember: the product of the concentrations of the hydrogen ion and the hydroxide ion must always equal 10^-14 for acidic or base aqueous solutions; because water is an amphoteric species (both a weak acid and a weak base) all acid-base reactivity in water ultimately reduces to the acid-base behavior of water, and all acidic or basic aqueous solutions are governed by the dissociation constant for water; thus if the dissociation constant for one species or its conjugate is known, then the dissociation constant for the other can be determined using the following equations: Ka,acid x Kb, conjugate base = Kw = 10^-14 Kb,base x Ka, conjugate acid = Kw = 10^-14 as the equation shows Ka and Kb are inversely related, in other words if Kb is large, then Kb is small and vice-versa; this means that a very strong acid (a approaching infinity) will produce a very weak conjugate base (ex. HCl is a strong acid and Cl- is a very weak base); a strong base will produce a very weak conjugate acid (for example, NaOH a strong base and H20 is a very weak acid); the conjugate of a very strong acid or base is sometimes termed inert because it is almost completely unreactive; on the other hand a weak acid and base will have weak conjugates

Hund's Rule

The rule that electrons will fill into separate orbitals with parallel spins before pairing within an orbital; the basis of this principle is based off electron repulsion; electrons in the same orbital tend to be closer to each other and thus repel each other more than electrons placed in different orbitals; to add to this we must also remember half filled and fully filled orbits have lower energies (higher stability) than other states; this is shown in chromium and copper (or other elements in their groups) chromium (z=24) should have the configuration [Ar] 4s^2 3d^4 however if you move one electron from the 4s subshell to the 3d subshell allows the 3d orbital to be half filled [Ar] 4s^1 3d^5; while moving 4s electron up to the 3d orbital is energetically unfavorable the extra stability from making the 3d half filled outweighs that cost; similarly copper [Ar] 4s^2 3d^9; a full d subshell out weighs the cost of moving an electron out of the 4s subshell; other elements in the same group have similar behavior, moving one electron from the highest s subshell to the highest d subshell. similar shifts can be seen with the f subshells, but they are never observed from the p subshell; the extra stability does not outweigh the cost

Activation Energy

(Ea) The minimum amount of energy required for a reaction to reach the transition state; also called energy barrier.

Faraday constant

(F) The total charge on 1 mole of electrons (F= 96,485 C/mole-); not to be confused with the farad (also denoted F), a unit of capacitance

Gram equivalent weight

(GEW) The amount of a compound that contains 1 mole of reacting capacity when fully dissociated; one GEW equals the molar mass divided by the reactive capacity (how many of the species of interest is obtained) per formula unit; equation is Gram Equivalent Weight = molar mass / n, where n is the number of particles of interest produced or consumed per molecule of the compound in the reaction; for example one would need 31 grams of H2CO3 (molar mass =62 grams/mol) to produce one equivalent of hydrogen ions because each molecule of H2CO3 can donate two hydrogen ion (n=2); simply put the equivalent weight of a compound is the mass that provides one mole of the particle of interest; if the amount of a compound in a reaction is known and we need to determine how many equivalents are present use the equation Equivalents = mass of compound (g) / gram equivalent weight (g)

Enthalpy

(H) The heat content of a system at constant pressure; the change in enthalpy (deltaH) in the course of a reaction is the difference between the enthalpies of the products and the reactants; enthalpy is a state function delta Hrxn = H products - H reactants a positive deltaHrxn corresponds to an endothermic process and a negative deltaHrxn corresponds to an exothermic process; it is not possible to measure enthalpy directly; only delta H can be measured, and only for certain fast and spontaneous processes

Acid dissociation constant

(Ka) The equilibrium constant that measures the degree of dissociation of an acid under specific conditions

Base dissociation contant

(Kb) The equilibrium constant that measures the degree of dissociation for a base under specific conditions

Equilibrium constant

(Keq) The ratio of the concentrations of the products to the concentrations of the reactants for a certain reaction at equilibrium, all raised to their stoichiometric coefficients; when dealing with gases, the equilibrium constant is referred to as Kp, the p indicates that it is pressure; for dilute solutions, Kc or Keq are interchangeably; while forward and reverse reactions rates are equal at equilibrium, the concentrations of the reactants and products are not usually equal; this means that the forward and reverse reaction rate constant, kf and kr, respectively are not usually equal to each other; the rate of kf to kr is Kc: Kc = Keq = kf/kr when a reaction occurs in more than one step, the equilibrium constant for the overall reaction is found by multiplying together the equilibrium constants for each step of the reaction, when this is done the equilibrium constant for the overall reaction is equal to the concentrations of the products divided by the concentrations of the reactants in the overall reaction, with each concentration term raised to the stoichiometric coefficient for the respective species, the forward and reverse rate constants for the nth step are designated kn and k-n, respectively; for example, if the reaction aA + bB <---> cC = dD occurs in three steps, each with a forward and reverse rate, then: Kc = (k1k2k3)/(k-1,k-2,k-3) = ([C]^c[D]^d)/([A]^a[B]^b)

Gas constant

(R) A proportionality constant that appears in the ideal gas law equation, PV=nRT. Its value depends on the units of pressure, temperature, and volume used in a given situation

Entropy

(S) A property related to dispersion of energy through a system or the degree of disorder in that system; the change in entropy (delta S) in the course of a reaction is the difference between the entropies of the products and the reactants; for example, hot tea cools down, frozen drinks melt, iron rusts, buildings crumble, balloons deflate, living things die and decay; the thermal energy in the tea is spreading out to the cooler frozen drink; the chemical energy in the bonds of elemental iron and oxygen is released and dispersed as a result of the formation of the more stable, lower energy bonds of iron oxide (rust) (Fe2O3), the potential energy of the building is released and dispersed in the form of light, sound, and heat as the building crumbles and falls; the energy of the pressurized air is released to the surrounding atmosphere as the balloon deflates; the chemical energy of all the molecules and atoms in living flesh is released into the environment during the process of death and decay; we must not think of entropy as disorder but as the measure of the spontaneous dispersal of energy at a specific temperature: how much energy is spread out, or how widely spread out energy becomes; the calculation for entropy is: S = Qsubrev/T, where delta S is the change in entropy, Qsubrev is the heat that is gained or lost in a reversible process, and T is the temperature in kelvin, units are usually J/mol*K; when energy is distributed into a system at a given temperature, its entropy increases; when energy is distributed out of a system at a given temperature its entropy decreases; entropy is a state function, so a change in entropy from one equilibrium state to another is pathway independent and only depends upon the difference in entropies of the final and initial states; further the standard entropy change for a reaction, delta S^osubrxn, can be calculated using the standard entropies of the reactants and products much like enthalpy: delta S^osubrxn = sigma deltaS^osubf,products - sigma deltaS^osubf,reactants; entropy is maximized at equilibrium

Effective nuclear charge

(Zeff) The charge perceived by an electron from the nucleus; applies most often to valence electrons and influences periodic trends such as atomic radius and ionization energy; as the positivity of the nucleus increases the electrons surrounding the nucleus including those in the valence shell, experience a stronger electrostatic pull towards the center of the atom, this causes the electron cloud to move closer and bind more tightly to the nucleus; for elements in the same period zeff increases from left to right; as one moves down the elements of a given group the n increased by one each time, this means the the valence electrons are increasingly separated form the nucleus by a greater number of filled principal energy levels, the results of this increased separation is a reduction in the electrostatic attraction between the valence electrons and the positively charged nucleus; this the zeff is more or less constant among the elements within a given group

Calorie

(cal) A unit of thermal energy

Gibbs free energy

(delta G) The energy of a system available to do work; this will determine whether or not a reaction will occur by itself without outside assistance; of note spontaneous does not mean quickly, most biological reactions are so slow without the aid of enzymes and other catalysts measurable reaction progress might not actually occur over the course of an average human lifetime; The change in Gibbs free energy, delta G, can be determined for a given reaction equation from the enthalpy change, temperature, and entropy change; a negative delta G denotes a spontaneous reaction, while a positive delta G denotes a non spontaneous reaction; the free energy change of the reaction (delta Grxn) is the difference between the free energy of the products and the free energy of the reactants, a negative free energy change indicates an exergonic reaction (energy is given off) and a positive free energy change indicates an endergonic reaction (energy is absorbed); the change in Gibbs free energy, delta G, is a measure of the change in the enthalpy and the change in entropy as a system undergoes a process, and it indicates whether a reaction is spontaneous or non-spontaneous, the change in the free energy is the maximum amount of energy released by a process- occurring at constant temperature and pressure - that is available to perform useful work; the equation for this is: delta G = deltaH - T * delta S, where T is the temperature in kelvins, and TdeltaS represents the total amount of energy that is absorbed by a system when its entropy increases reversibly; mnemonic: Goldfish ARE (equals sign) Horrible WITHOUT (minus sign) Tartar Sauce; a way to visualize Gibbs free energy is picturing a ball rolling into a valley between two hills the ball would eventually rest at the lowest point in the valley, any system - including chemical reactions - will move in whichever direction results in a reduction of the free energy of the system, the bottom of the valley represents equilibrium, and the sides of the hill represent the various points in the pathway toward or away from equilibrium; movement toward the equilibrium position is associated with a decrease in Gibbs free energy (delta G < 0) and is spontaneous; when a system releases energy it is said to be exergonic; on the other hand movement away from the equilibrium position is associated with an increase in Gibbs free energy (delta G> 0) and is non spontaneous, such a reaction is said to endergonic; of note endergonic/exergonic (describing Gibbs free energy) are not the same as endothermic/exothermic (describing enthalpy (heat/thermic)); when delta G is zero, the system is in a state of equilibrium; delta H = TdeltaS; of note phase equilibria are states in which more than one phase exists, therefore delta G = G (g) - G (s) = 0 so G (g) = G (s); temperature is in kelvin therefore always positive; + H & + S = spontaneous at high T; + H & - S = non-spontaneous at all T; - H & +S = spontaneous at all T; -H & -S = spontaneous at low T; another way to calculate standard free energy change is delta G^osubrxn = -RTlnKeq, where R is the ideal gas constant, T is the temperature in kelvins, and Keq is the equilibrium constant; this equation allows us to make not only quantitative evaluations of the free energy change of a reaction but also the qualitative assessments of the spontaneity of the reaction; the more positive the natural logarithm, the more negative the standard free energy change; the more negative the standard free energy change, the more spontaneous the reaction

Heat of fusion

(delta Hfus) The enthalpy change for the conversion of 1 gram or 1 mole of a solid to a liquid at constant temperature and pressure

Heat of sublimation

(delta Hsub) The enthalpy change for the conversion of 1 gram or 1 mole of a solid to a gas at constant temperature and pressure

Heat of vaporization

(delta Hvap) The enthalpy change for the conversion of 1 gram or 1 mole of a liquid to a gas at constant temperature and pressure

Isovolumetric process

(isochoric) A process that occurs at constant volume in which the system performs no work; also called an isochoric process; gas neither expands nor compresses; the first law simplifies to delta U = Q (the change in internal energy is equal to the heat added to the system); an isovolumetric process is a vertical line on a pressure-volume graph; the area under the curve which represents the work done by the gas, is zero

Azimuthal quantum number

(l) The quantum number denoting the sub level or sub shell in which an electron can be found; reveals the shape of the orbital; this is very important because it has important implications for chemical bonding and bond angles; the value of n limits the value of l in the following way: for any given value of n, the range of possible values for l is 0 to (n-1); (ex. n=1 the only possible value of l is 0; for n=2 the values of l are 0 and 1) a simple say to remember this is the n-value also tells you the number of possible subshells, ex. there is only one subshell (l=0) in the first princial energy level; there are two subshells (l=0 and 1) in the secondary principal energy level; within each subshell there is a capacity to hold a certain number of electrons given by 4l+2; the energies of the subshells increase with increasing l value; however, the energies of subshells from different principal energy levels may overlap, for example the 4s subshell will have a lower energy than the 3d subshell

Solubility Rules (7)

1. All salts containing ammonium (NH4+) and alkali metal (group 1) cations are water-soluble 2. All salts containing nitrate (NO3-) and acetate (CH3COO-) anions are water-soluble 3. Halides (Cl-, Br-, and I-) excluding fluorides, are water soluble, with the exceptions of those formed with Ag+, Pb2+, and Hg2 2+ 4. All salts of the sulfate ion (SO4 2+) are water soluble, with the exceptions of those formed with Ca 2+, Sr 2+, Ba 2+, and Pb 2+ 5. All metal oxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides 6. All hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and Ca 2+, Sr 2+, and Ba 2+, 7. All carbonates (CO3 2+, phosphates (PO4 3+), sulfides (S 2-) and sulfites (SO3 2-) are insoluble, with the exception of those formed with the alkali metals and ammonium the MCAT will not expect you to memorize all of the solubility rules but it is worth knowing: all salts of Group 1 metals, and all nitrate salts are soluble

Factors that Affect Reaction Rate

1. reaction concentrations: the greater the concentration of the reactants the greater the number of effective collisions per unit time 2. Temperature: for nearly all reactions the reaction rate will increase as the temperature increases; because the temperature of a substance is a measure of the particles' average kinetic energy, increasing the temperature increases the average kinetic energy of the molecules; if temperatures elevate too high a catalyst may denature and then the reaction rate plummets 3. Medium: the rate at which a reaction takes place may also be affected by the medium in which it takes place, some molecules are more likely to react with each other in aqueous environments, while others are more likely to react with each other in non-aqueous solvents, such as in dimethyl sulfoxide (DMSO) or ethanol; furthermore, the physical state of the medium (liquid, solid, or gas) can also have significant effect; generally polar solvents are preferred because their molecules dipole tends to polarize the bonds of the reactants, thereby lengthening and weakening them, permitting the reaction to occur faster 4. Catalyst: are substances that increase reaction rate without themselves being consumed in the reaction; these interact with the reactants, either by adsorption or through the formation of intermediates, and stabilize them so as to reduce the activation energy necessary for the reaction to proceed; while many catalyst including all enzymes, chemically interact with the reactants, they return to their original chemical state upon formation of the products; they may increase the frequency of collisions between the reactants, change the relative orientation of the reactant, making a higher percentage of the collision effective; donate electron density to the reactants; or reduce intramolecular bonding within reactant molecules

Law of Mass Action

: the form of the equilibrium contestant; has the concentrations of products over concentrations of reactants, each raised to their stoichiometric coefficients; if the system is at equilibrium at a constant temperature, than the following ratio is constantKeq = ([C]^c [D]^d) / ([A]^a [B]^b) Key concept: at equilibrium that rate of the forward reaction equals the rate of the reverse reaction, entropy is at a maximum, and Gibbs free energy is at a minimum; this links the concepts of thermodynamics and kinetics; properties of law of mass action and equilibrium constant expression: concentrations of pure solids and pure liquids do not appear in the equilibrium constant expression, this is due to equilibrium expression being based on the activities of compounds, not concentrations; the activities of pure solids and liquids are defined to be 1. for the purpose of the MCAT, there is a negligible difference between concentration and activity; Keq is characteristic of a particular reaction at a given temperature; the equilibrium constant is temperature-dependent; the larger the value of keq, the farther to the right the equilibrium position; if the equilibrium constant for a reaction written in one direction is keq, the equilibrium constant for the reverse reaction is 1/keq

Electrolyte

A compound that ionizes in water and increases the conductance of the solution; solid ionic compounds tend to be poor conductors of electricity because the charged particles are rigidly set in place by the lattice arrangement of the crystalline solid. In aqueous solutions however the lattice arrangement is disrupted by the ion-dipole interactions between the ionic components and the water molecules. The cations and anions are now free to move, and as a result, the solution of ions is able to conduct electricity; solutes that are enable solutions to carry currents are called electrolytes; the electrical conductivity of aqueous solutions is governed by the presence and concentration of ions in the solutions; subsequently the number of electron equivalents being transferred in such a system, such as in electrochemical cells, varies pure water which has no ions other than the very few hydrogen ions and hydroxide ions that result from water's low level autodissociation, is a very poor conductor

Combustion reaction

A reaction in which an oxidant (typically oxygen) reacts with a fuel (typically a hydrocarbon) to yield water and an oxide (such as carbon dioxide if between a hydrocarbon and oxygen); for example the balanced equations expression the combustion of methane CH4 + 2O2 ---> CO2 + 2H2O

First-order reaction

A reaction in which the rate is directly proportional to the concentration of only one reactant, such that doubling the concentration of that reactants results in an doubling of the rate of formation of the product, the rate law for a first order reaction is: rate = k [A]^1 or rate = k [B]^1 it is important to recognize that a first order rate law with a single reactant suggests that the reaction begins when the molecule undergoes a chemical change all by itself, without a chemical interaction, and usually without a physical interaction with any other molecule; plotting a first order reaction on a concentration vs time curve results in a nonlinear graph, the curve shows that the rate of formation of product is dependent on the concentration of reactant; plotting ln [A] vs time reveals a straight line; the slope of such a line is the opposite of the rate constant k; k=-slope; temperature, changing reactant concentrations, and catalyst affect rate of first order; lowering temperatures decreases rate; doubling reactants doubles rate; and adding catalyst increases rate

Combination reaction

A reaction in which two or more reactants form a single product; the formation of water by burning hydrogen gas in air is an example of a combination reaction 2H2 (g) + O2 (g) -> 2H2O (g)

Hydrolysis

A reaction in which water is consumed during the breakdown of another molecule; in neutralization reactions this is the reverse of the reaction ex: HA (aq) + BOH (aq) <--> BA (s) + H2O (l)

Endothermic reaction

A reaction that absorbs heat from the surroundings as the reaction proceeds (positive deltaH)

Exothermic reaction

A reaction that gives off heat to the surroundings (negative delta H) as the reaction proceeds

Irreversible reaction

A reaction that proceeds in one direction only and goes to completion

Collision Theory of Chemical Kinetics

A theory that states that the rate of a reaction is proportional to the number of collisions per second between reacting molecules that have sufficient energy to overcome the activation energy barrier; implies that only a fraction of collisions are sufficient, not all collisions result in a chemical reaction; an effective collision (one that leads to the formation of products) occurs only if the molecules collide with each other in the correct orientation and with sufficient energy to break their existing bonds and form new ones; the minimum energy of collision necessary for a reaction to take place is called the activation energy or energy barrier; only a fraction of colliding particles have enough kinetic energy to exceed the activation energy; this means that only a fraction of all collisions are effective; the rate of a reaction can therefore be expressed as follows: rate = Z x f where Z is the total number of collisions occurring per second and f is the fraction of collisions that are effective A much more quantitatively rigorous analysis of the collision theory can be accomplished through the Arrhenius equation which is given: k = Ae^(-Ea/RT) where k is the rate constant of a reaction, A is the frequency factor, Ea is the activation energy of the reaction, R is the ideal gas constant, and T is the temperature in kelvins The frequency factor also known as the attempt frequency of the reaction is a measure of how often molecules in a certain reaction collide, with the units s^-1; it is important to understand the relationship between the variables in Arrhenius equation; for example the relationship between A and K is evident in the equation, as the frequency factor of the reaction increases, the rate constant of the reaction also increase in a direct relationship; a more complex relationship cab be seen with temperature of a chemical system were to increase to infinity, while all other variables are held constant, the value of the exponent would have a magnitude less than 1. however before assuming that the rate constant is going to decrease as a result, note the presence of a negative sign; as the magnitude of the exponent gets smaller it actually moves from a more negative value towards zero, the exponent thus becomes less negative which means the rate constant actually increases; this should make sense conceptually because the rate of a reaction increases with temperature; of note the frequency factor can be increased by increasing the number of molecules in a vessel, when there are more molecules, the opportunities for collision are increased

Basic solution

An aqueous solution that contains more OH- ions than H+ ions; pH >7 under standard conditions

Freezing point depression

Amount by which a given quantity of solute lowers the freezing point of a liquid; a colligative property; the presence of solute particles in a solution interferes with the formation of the lattice arrangement of solvent molecules associated with the solid state; thus a greater amount of energy must be removed form the solution (resulting in a lower temperature) in order for the solution to solidify equation: delta Tsubf = i*Ksubf*m where deltaTsubf is the freezing point depression, i is the van't hoff factor, Kf is the proportionality constant characteristic of a particle solvent, and m is the molality of the solution; example of why we put salt on roads in the winter is on pg 326

Electrode

An electrical conductor through which an electrical current is used to power an otherwise non spontaneous decomposition reaction

Daniell Cell

An electrochemical cell in which the anode is the site of Zn metal oxidation and the cathode is site of Cu2+ ion reduction

Galvanic cell

An electrochemical cell that uses a spontaneous oxidation-reduction reaction to generate an electromotive force; also called a voltaic cell; all of the non-rechargeable batteries you own are galvanic cells; these reactions are spontaneous (think of the batteries providing energy to flashlight or remote control), this means free energy is decreasing (delta G < 0) as the cell releases energy to the environment; if free energy change is negative the electromotive force (Esubcell) must be positive; (of note: the free energy change and electromotive force always have opposite signs); as shown in a Daniell Cell as the spontaneous reaction proceeds toward equilibrium, the movement of electrons results in a conversion of electrical potential energy into kinetic energy; by separating the reduction and oxidation half reactions into two compartments we are able to harness this energy and use it to do work by connecting various electrical devices into the circuit between the two electrodes

Electrolytic cell

An electrochemical cell that uses an external voltage source to drive a non spontaneous oxidation-reduction reaction; comparing electronic cell to galvanic

Chemical equation

An expression used to describe the quantity and identity of the reactants and products of a reaction

Cation

An ionic species with a positive charge; cats are positive

Propane

C3H8

Alkaline earth metals

Elements found in Group IIA of the period table; chemistry I similar to that of the alkali metals, except that they have two valence electrons and, thus, form +2 (divalent) cations; they have slightly higher effective nuclear charges and thus slightly smaller atomic radii compared to alkali metals; together the alkaline earth metals and the alkali metals are called the active metals because they are so reactive that they are not naturally found in their elemental (neutral) state

Chalcogens

Elements found in Group VIA or Group 16 of the period table with diverse chemistry; the group contains metals, nonmetals (like oxygen), and metalloids; typically form -2 anions; these are not as reactive as the halogens, they are crucial for normal biological functions; they each have 6 electrons in their valence electron shell and due to their proximity to the metalloids, generally have small atomic radii and large ionic radii; oxygen is the most important element in this group for many reasons, it is one of the primary constituents of water, carbohydrates, and other biological molecules; sulfur is also an important component of certain amino acids and vitamins; selenium also has an important nutrient for microorganisms and has a role in protection from oxidative stress; the remainder of this group is primarily metallic and generally toxic to living organisms; it is important to note that at high concentrations many of these elements-no matter how biologically useful can be toxic or damaging

Halogens

The active nonmetals in Group VIIA of the period table, which have high electronegativities and high electron affinities, this means they are highly active nonmetals with seven valence electrons; the physical properties of this group are variable; they are especially reactive towards the alkali and alkaline earth metals; F is the highest electronegative of all elements; these are so reactive they are not naturally found in their elemental state but rather as ions (called halides) or diatomic molecules

Boiling Point Elevation

The amount by which a given quantity of solute raises the boiling point of a liquid; a colligative property; the boiling point elevation formula calculates the amount the normal boiling point is raised, the value calculated is not the boiling point itself; when a nonvolatile solute is dissolved into a solvent to create a solution, the boiling point of the solution will be greater than that of the pure solvent; the boiling point is the temperature at which the vapor pressure of the liquid equals the ambient (incident) pressure; adding solute to a solvent results in a decrease in the vapor pressure of the solvent, then more energy (and consequently a higher temperature) will be required before its vapor pressure equals the ambient pressure; the extent to which the boiling point of a solution is raised relative to that of the pure solvent is given by the formula deltaTsubb= iKsubb*m where delta Tsubb is the increase in boiling point, i is the van't hoff factor, Ksubb is a proportionality constant characteristic of a particular solvent (which will be provided for MCAT), and m is the molality of the solution

Concentration

The amount of solute per unit of solvent or the relative amount of one component in a mixture; there are many different ways of expressing concentration, EtOH is expressed in volume percent (volume of solute divided by volume of solution times 100%; alcohol proof is twice the volume percent); the MCAT will express concentration by percent composition by mass, mole fraction, molarity, molality, and normality

Atomic radius

The average distance between a nucleus and its outermost electron; usually measured as one-half the distance between two nuclei of an element in its elemental form; the atomic radius cannot be measured by examining a single atom because the electrons are constantly moving around, making it impossible to mark the outer boundary of the electron cloud; as you move from left to right on the period table protons and electrons are added one at a time to the atoms, the electrons are added to the outermost shell and the number of inner shell electrons remains constant, the increasing positive charge of the nucleus pulls the outer electrons more closely inward and holds them more tightly; the zeff increases left to right across a period, as a result, atomic radius decreases from left to right across a period; as you move down a group the increasing principal quantum number implies that the valence electrons will be found farther away from the nucleus because the number of inner shells is increasing, separating the valence shell form the nucleus; the zeff essentially remains the same however the atonic radius increased down the group; for reference the largest atomic radius in teh period table is Cs (cesium) and the smallest is He (helium), francium is typically not considered because it is exceptionally rare in nature

Heisenberg uncertainty principle

The concept that states that it is impossible to determine both the momentum and position of an electron simultaneously with perfect accuracy; if we want to assess the position of an electron, the electron has to stop (thereby removing its momentum); if we want to assess its momentum, the electron has to be moving (thereby changing its position)

Formal charge

The convention assignment of charges to individual atoms of a Lewis structure for a molecule; the total number of valence electrons in the free atom minus the total number of electrons when the atom is bonded (assuming equal splitting of the electrons in bonds)

Charle's Law

The law stating that the volume of a gas at constant pressure is directly proportional to its absolute (kelvin) temperature V1/T1= V2/T2; this law states volume and temperature are directly proportional, when one increases the other increases in direct proportion

Effusion

The movement of gas from one compartment to another under pressure through a small opening; follows Graham's Law

Coordination number

The number of atoms that are bound to a central atoms, is the relevant factor when determining molecular geometry; for example consider that CH4 (methane), NH3 (ammonia) and H2O all have the same electronic geometry: in each compound, four pairs of electrons surround the central atom, this is tetrahedral electronic geometry; however because each molecule has a different coordination number they have different molecular geometries; in molecular geometry, methane has tetrahedral geometry, ammonia has trigonal pyramidal geometry, and water is identified as angular or bent

Gas

The physical state of matter possessing the most disorder, in which molecules interact through very weak attractions (low intermolecular interactions); have far apart molecules; found at relatively low pressure and high temperatures; gases are fluid and easily - although not indefinitely - compressible which distinguishes them from liquids; we can define the state of a gaseous sample by four variables - pressure (P), volume (V), temperature (T) and number of moles (n); gas pressure is usually defined in atmospheres (atm) or in millimeters of mercury (mmHg) which are equivalent to torr; the SI unit for pressure however is the pascal (Pa); 1 atm = 760 mmHg = 760 torr = 101.325 kPa

Half-equivalence point

The point at which half a given species within a titration has been protonated or deprotonated

Endpoint

The point in a titration at which the indicator changes to its final color

Equivalence Point

The point in a titration at which the moles of acid present equal the moles of base added, or vice-versa; it is important to emphasize that, while a strong acid/strong base titration will have its equivalence point at a pH of 7, the equivalence point does not always occur at pH 7; when titrating polyprotic acids or bases, there are multiple equivalence points as each acidic or basic conjugate species is titrated separately; at the equivalence point the number of equivalents of acid and base are equal, this fact allows us to calculate the unknown concentration of the titrand through the equation: NaVa = NbVb where Na and Nb are the acid and base normalities and Va and Vb are the volumes of acid and base solutions; note that as long as the units are consistent for volume it does not have to be liters the equivalence point is calculated in two common ways: evaluated by using a graphical method, plotting the pH of the unknown solution as a function of added titrant by using a pH meter, or estimated by watching for a color change of an indicator of note: strong acid + weak base: equivalence point pH < 7 strong acid + strong base: equivalence point pH 7 weak acid + strong base :equivalence point pH > 7

Buffer region

The portion of a titration curve in which the concentration of an acid is approximately equal to that of its conjugate base; pH remains relatively constant through this region

Autoionizaiton

The process by which a molecule (usually water) spontaneously dissociates into cations and anions example: H2O (l) + H2O (l) <--> H3O+ (aq) + OH- (aq) (H3O+ is hydronium ion and OH- is hydroxide ion)

Condensation

The process in which a gas transitions to the liquid state; in a covered or closed container the escaping molecules are trapped above the solution, these molecules exert a countering pressure, which forces some of the gas back into the liquid phase; this is facilitated by lower temperatures or higher pressure, atmospheric pressure acts on a liquid in a manner similar to that of an actual physical lid; as evaporation and condensation proceed, the respective rates of the two processes become equal, and equilibrium is reached, the pressure that the gas exerts over the liquid at equilibrium is the vapor pressure of the liquid; vapor pressure increase as temperature increase because more molecules have sufficient kinetic energy to escape into the gas phase; the temperature at which the vapor pressure of the liquid equals the ambient (also known as external, applied, or incident) pressure is called the boiling point

Freezing

The process in which a liquid transitions to the solid state; also known as solidification or crystallization; the temperature at which these processes occur is called the melting point or freezing point, depending on the direction of the transition; whereas pure crystalline solids have distinct, very precise melting points, amorphous solids, such as glass, plastic, chocolate, and candle wax, tend to melt (or solidify) over a larger range of temperature due to their less-ordered molecular structure

Electrolysis

The process in which an electrical current is used to power an otherwise non spontaneous decomposition reaction

Chelation

The process of binding metal ions to the same ligand at multiple points; this generally requires large organic ligands that can double back to form a second or even third, bond with the central cation; chelation therapy is often used to sequester toxic metals (lead, arsenic, mercury, and so on) even biological necessary metals such as iron, can be toxin in overload

Discharging

The state of a rechargeable electrochemical cell that is providing an electromotive force by allowing electrons to flow spontaneously from anode to cathode

Hydrogen Bonding

The strong attraction between a hydrogen atom bonded to a highly electronegative atoms (such as nitrogen, oxygen, or fluorine) In one molecule and a highly electronegative atom in another molecule; hydrogen bonds is a specific, unusually strong form of dipole dipole interaction that may be intra- or intermolecular; these are not actual bonds, there is no sharing or transferring of electrons between two atoms; the hydrogen atom essentially acts as a naked proton; the positively charged hydrogen atom interacts with the partial negative charge fluorine, oxygen, or nitrogen on nearby molecules; substances that display hydrogen bonding tend to have unusually high boiling points compared to compounds of similar molecule weights that do not exhibit hydrogen bonding; the difference derives from the energy required to break the hydrogen bonds

f subshell

The sub shell corresponding to the angular momentum quantum number l=3; contains seven orbitals and is found int eh fourth and higher principal energy levels

Formula weight

The sum of the atomic weights of constituent ions according an ionic compound's empirical formula

Evaporation

The transition from a liquid to a gaseous state; some of the molecules neat the surface of the liquid may have enough kinetic energy to leave the liquid phase and escape into the gaseous phase; each time the liquid loses a high energy particle, the temperature of the remaining liquid decreases; evaporation is an endothermic process for which the heat source is the liquid water; of course the liquid water itself may be receiving thermal energy from some other source, as in the case of a puddle of water drying up under the hot summer sun or a pot of water on a stovetop; given enough energy the liquid will completely evaporate; boiling is a very specific type of vaporization that occurs only under certain conditions; any liquid will lose some particles to the vapor phase over time; however boiling is the rapid bubbling of the entire liquid with rapid release of the liquid as gas particles, while evaporation happens in all liquids at all temperatures, boiling can only occur above the boiling point of a liquid and involves vaporization through the entire volume of liquid

Electron configuration

The symbolic representation used to describe the electron arrangement within the energy sub levels in a given atom; this uses spectroscopic notation, where in the first number denotes the principal energy level, the letter designates the subshell, and the superscript give the number of electrons in that subshell; for example 2p^4 indicated that there are four electrons in the second p subshell of the second principal energy level, this also implies that the energy levels below 2p (that is 1s, and 2s) have already been filled; electrons fill from lower to higher energy subshells according to the Aufbau principle (also called building up principle) and each subshell will fill completely before electrons begin to enter the next one; an easy way to remember this is the n+l rule, for example which will fill first the 5d subshell or the 6s subshell, for 5d, n=5 and l=2 so n+l=7, for 6s n=6 and l=0 so n+l =6 so therefore 6s subshell has lower energy and will fill first; electron configuration can be abbreviated by placing the noble gas that precedes the element of interest in brackets, for example, any element in period four can be abbreviated by starting with [Ar]; example what is the electron configuration of osmium, [Xe] 6s^2 4f^14 5d^6; for ions such as anions (negatively charged ions) have additional electrons that fill according to the same rules as above; for example fluorine is [He] 2s^2 2p^5, F- is [He] 2s^2 2p^6; positively charged ions (cations) are more complicated, start with the neutral atom and remove electrons from the subshells with the highest value for n first, if multiple subshells are tied for the highest n value, then electrons are removed from the subshell with the highest l value; example Fe^3+, start with [Ar] 4s^2 3d^6, then take 2 electrons from the 4s^2 and one from 3d^6 and you get [Ar] 3d^5

Absolute zero

The temperature at which all substances have no thermal energy; 0K or -273.15 C

Boiling Point

The temperature at which the vapor pressure of a liquid is equal to the incident pressure; the normal boiling point of any liquid is defined as its boiling point at a pressure of 1 atmosphere

Kinetic molecular theory

The theory proposed to account for the observed behavior of gases; considers gas molecules to be point like, volume less particles exhibiting no intermolecular forces that are in constant random motion and undergo only completely elastic collisions with the container or other gas particles; this theory was developed in reference to ideal gases however it can be reasonably accurate to real gases as well; assumptions can be made about gases due to this theory, such as: 1. gases are made up of particles with volumes that are negligible compared to the container volume 2. gas atoms or molecules exhibit no intermolecular attractions or repulsions 3. gas particles are in continuous, random motion, undergoing collisions with other particles and the container wall 4. Collisions between any two gas particles (or between particles and the container walls) are elastic, meaning that there is conservation of both momentum and kinetic energy 5. the average kinetic energy of gas particles is proportional to the absolute temperature of the gas (in K), and it is the same for all gases at a given temperature, irrespective of chemical identify or atomic mass according to the kinetic molecular theory of gases, the average kinetic energy of a gas particle is proportional to the absolute temperature of the gas: KE - 1/2 mv^2 = 3/2 ksubB*T; where ksubB is the Boltzmann constant (1.38x10^-23J/K) which serves as a bridge between the macroscopic and microscopic behaviors of gases; this equation shows the speed of a gas particle is related to its absolute temperature; key concept: the higher the temperature the faster the molecules move, the larger the molecules the slower they move

Ground state

The unexcited state of an electron; state of the lowest energy of an atom; electrons are all in the lowest possible orbitals

atomic absorption spectrum

The unique spectrum of light absorbed when an atom's electrons are excited to higher energy levels; for an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy precisely equal to the energy difference between the two levels

Reversible reaction

a reaction that can proceed in either the forward or reverse direction, and typically does not go to completion

strong base

a base that undergoes complete dissociation in an aqueous solution; for example when sodium hydroxide is added to water, the ionic compound dissociates according to the net ionic equation: NaOH (s) --> Na + (aq) + OH- (aq) hence, in a 1 M NaOH solution, complete dissociation such as NaOH, will remain in the solution; this is why the dissociation of strong acids and bases is said to go to completion; the contribution of OH- and H+ ions from the auto-ionization of water is negligible if the concentration of the acid or base is significantly greater than 10^-7 M; on the other hand if the concentration of acid or base is close to 10^-7 M, then the contribution form the auto-ionization of water is important; strong bases commonly seen on the MCAT include NaOH (sodium hydroxide), KOH (potassium hydroxide), and other soluble hydroxides of Group IA metals calculations of the pH and pOH of strong acid and bases assume complete dissociation of the acid or base in solution

weak base

a base that undergoes partial dissociation in an aqueous solution; weak bases are also known as dilute, while strong bases are known as concentrated; the base dissociation constant (Kb) can be calculated as: Kb = [B+][OH-]/[BOH] the smaller the Kb, the weaker the base; and consequently the less it will dissociate; as with the acid dissociation expression, water is not included because it is a pure liquid; we can characterize a species as a weak base if its Kb is less than 1.0; on the MCAT molecular (nonionic) weak bases are almost exclusively amines

Sublimation

a change of phase from solid to gas without passing through the liquid phase; an example of this is dry ice (solid CO2) sublimes at room temperature and atmospheric pressure, the absence of the liquid phase makes it a convenient dry refrigerant

Ion

a charged atom or molecule that results from the loss of or gain of electrons

salt bridge

a component of an electrochemical cell composed of an inert electrolyte that allows the charge gradient that builds up in the half-cells to be dissipated as a reaction occurs; contains ions that will not react with electrodes or ions in solution and that can move to balance charge

normality (N)

a concentration unit equal to the number of equivalents per liter of solution; on the MCAT it is most commonly used for hydrogen ion concentration; thus 1 N solution of acid contains a concentration of hydrogen ions equal to 1 mole per liters; a 2 N solution of acid contains a concentration of hydrogen ions equal to 2 moles per liter; the actual concentration of the acidic compound may be the same or different from the normality because different compounds are able to donate different numbers of hydrogen ions; In a 1 N HCl solution, the molarity of HCL is 1 M because HCl is a monoprotic acid; in a 1 N H2CO3 solution the molarity of H2CO3 is 0.5 M because H2CO3 is a diprotic acid; note that normality calculations always assume that a reaction will proceed to completion; while carbonic acid does not fully dissociate in solutions, it can be reacted with enough base for each molecule to give up both of its protons; the conversion from normality to molarity of a given solute is Molarity = Normality / n, where n is the number of protons, hydrogen ions, electrons, or ions produced or consumed by the solute of note simple ideas on the MCAT make things easier; when thinking of normality think of it as molarity of the stuff of interest in the reaction

Molality (m)

a concentration unit equal to the number of moles of solute per kilogram of solvent; equation: m = moles of solute/kilograms of solvent you wont have to use molality often on the MCAT however be mindful of the specific situations when it is requried: boilign point elevation and freezing point depression

molarity (M)

a concentration unit equal to the number of moles of solute per kilogram of solvent; this is the most common way of expressing concentration on the MCAT equation: M = moles of solute/liters of solution of note the volume term in the denominator of molarity refers to the solution volume, NOT the volume of solvent used to prepare the solution - although the two values are often close enough to proximate the solution volume using the solvent volume

paramagnetism

a condition that rises when a substance has unpaired electrons and is slightly attracted to a magnetic field; the presence of paired or unpaired electrons affects the chemical and magnetic properties of an atom or molecule; material composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field, and the material will thus be weakly attracted to the magnetic field, these materials are considered paramagnetic; an example of this is a set of iron orbs that are influenced by a magnetic, the metallic spheres that are close enough to be induced by the magnet are attracted to the magnet and move towards it; a good mnemonic for this is to remember PARAmagnetic means that a magnetic filed will cause PARAllel spins in unpaired electrons and therefore cause an attraction

Polar Covalent Bond

a covalent bond between atoms with different electronegativities in which electron density is unevenly distributed, giving the bond positive and negative ends

nonpolar covalent bond

a covalent bond between elements of similar electronegativity; contains no charge separation

Potentiometer

a device used to measure electromotive force (voltage); these can be used in titrations: redox titrations with no indicator

pH meter

a device used to measure the concentration of hydrgoen ions in solution and report it as a pH value

Resonance

a difference in the arrangement of electron pairs but not the bond connectivity or overall charge within a Lewis structure; resonance structures are represented with double headed arrows between them, the actual electronic distribution in the compound is a hybrid or composite of all of the possible resonance structures; ex SO2 has three resonance structures, :O: = S: = :O: <-> :O: = S:+1 - O:::-1 <-> :::O-1 - S:+1 = :O:; the minor contributors to the resonance hybrid contain formal charges indicating decreased stability; one can use formal charge to access the stability of the resonance structures according to the following guidelines: 1. A lewis structure with small or no formal charge is preferred over a lewis structure with large formal charges 2. A lewis structure with less separation between opposite charges is preferred over a lewis structure with a large separation of opposite charged 3. A lewis structure in which negative formal charges are placed on more electronegative atoms is more stable than one in which the negative formal charges are placed on less electronegative atoms

stoichiometry

a form of dimensional analysis focusing on the relationships between amounts of reactants and products in a reaction; these fractions demonstrate an underlying three step process: convert from the given units to moles, use the mole ratio, convert from moles to the desired units; common conversions used in stoichiometry include: 1 mole of any ideal gas at STP = 22.4 L; 1 mole of any substance = 6.22 x 10^23 particles (avogadro's number); 1 mole of any substance = its molar mass in grams (from the periodic table)

Molecular formula

a formula showing the actual number and identity of all atoms in each molecule of a compound; always a whole-number multiple of the empirical formula; for example the molecular formula for benzene is C6H6 while the empirical formula is CH

state function

a function that depends on the state of a system but not on the path used to arrive at that state; state functions are independent of the path (process) taken, they are not necessarily independent of one another, for example Gibbs free energy is related to enthalpy, temperatures, and entropy; includes pressure, density, temperature, volume, enthalpy, internal energy, Gibbs free energy, and entropy; mnemonic: when im under PRESSURE and feeling DENSE, all I want to do is watch TV and get HUGS;

Real Gas

a gas that exhibits deviations from the ideal gas law due to molecular attractions and the actual volume of the gas molecules themselves; at high temperature and low pressure (high volume) deviations from idealists are usually small; good approximations can still be made from the ideal gas law; when gas atoms or molecules are forced into close proximity under high pressure (at low volume) or at low temperatures deviations from ideal gas behavior are noticeable; deviations occur due to pressure because as the pressure of gas increases the particles are pushed closer and closer together, as the condensation pressure for a given temperature is approached intermolecular forces become more and more significant until the gas condense into liquid, also at moderately high pressure a gas's volume is less than would be predicted by the ideal gas law due to intermolecular attraction, at extremely high pressure however the size of particles becomes relatively larger compared to the distance between them and this causes the gas to take up a larger volume; as temperature decreases the average speed of the gas molecules decreases and the attractive intermolecular forces become increasingly significant, as the condensation temperature approaches for a given pressure intermolecular attraction eventually cause the gas to condense to a liquid state, as the temperature of a gas is reduced toward its condensation point intermolecular attractions cause the gas to have a smaller volume than that which would be predicted by the ideal gas law, the closer the gas is to its boiling point the less ideal it acts, at extremely low temperatures gas will again occupy more space than predicted by ideal gas law because the particles cannot be compressed to zero volume

potential energy diagram

a graph that shows the potential energies of the reactants and products of a reaction during the course of the reaction; by convention, the x-axis shows the progress of the reaction and the y-axis shows potential energy

Sigma bond

a head to head bond between two orbitals of different atoms that allows free rotation about its axis

Solution

a homogeneous mixture of two or more substances that combine to form a single phase, usually liquid; it may be solid (brass), liquid (HCL (aq)), or gas (air); solutions are made up of solute (such as NaCl, NH3, C6H12O6, or CO2) dissolved (dispersed) in a solvent (such as H2O, benzene, or ethanol); the solvent is the component of the solution that remains in the same phase after mixing; if the two substances are already in the same phase (for example two liquids) the solvent is the component present in greater quantity; if the two same-phase components are in equal proportions in the solution, then the component that is more commonly used as a solvent in other contexts is considered the solvent; solute molecules move about freely in the solvent and interact with it by way of intermolecular forces such as ion-dipole, dipole-dipole, or hydrogen bonding; dissolved solute molecules are also relatively free to interact with other dissolved molecules of different chemical identities; consequently, chemical reactions occur easily in solution

Period

a horizontal row of the periodic tablet containing elements with the same number of electron shells

Raoult's Law

a law stating that the partial pressure of a component in a solution is proportional to the mole fraction of that component in the solution; provides an explanation for vapor pressure depression seen in solutions; as a solute is added to a solvent, the vapor pressure of the solvent decreases proportionately; on a molecular level the presence of the solute molecules can block the evaporation of solvent molecules but not their condensation; this reduces the vapor pressure of the solution compared to the pure solvent equation: PsubA = XsubA*PsubA where Pa is the vapor pressure of the solvent A when solutes are present, XsubA is the mole fraction of the solvent A in the solution, and PsubA is the vapor pressure of the solvent A in its pure states this law holds only when the attraction between the molecules of the different components of the mixture is equal to the attraction between the molecules of any one component in its pure state; when this condition does not hold, the relationship between mole fraction and vapor pressure will deviate from Raoult's law; solutions that obey this law are called ideal solutions; this law states that ideal solution behavior is observed when solute-solute, solvent-solvent, and solute-solvent interaction are all very similar

Electronegativity

a measure of the attractive force that an atom will exert on an electron in a chemical bond; the greater the electronegative of an atom, the more it attracts electrons within a bond; electronegative values are related to ionization energies: the lower the ionization energy, the lower the electronegativity; the higher the ionization energy the higher the electronegativity; the first three noble gases are exception: despite their high ionization energies, these elements have negligible electronegativity because they do not often form bonds commonly measured with the Pauling scale which ranges from 0.7 for cesium, the least electronegative (most electropositive) element to 4.0 for fluorine, the most electronegative element; electronegativity increased across a period from left to right and decreased in a group from top to bottom

temperature

a measure of the average kinetic energy of the particles in a system; temperature is the way that we scale how hot or cold something is; temperature scales are Fahrenheit, Celsius, and Kelvin; the average kinetic energy of the particles in a substance is related to the thermal energy (enthalpy) of the substance, but because we must also include consideration of how much substance is present to calculate total thermal energy content, the most we can say about temperature is that when a substance's thermal energy increases, its temperature also increases; nevertheless, we cannot say that something that is hot necessarily has greater thermal energy (in absolute terms) that a substance that is cold. for example we might determine that a large amount of lukewarm water has a greater total heat content than a very small amount of hot water; the absolute temperature scale, Kelvin, was determined via the third law of thermodynamics which elucidated that there is a finite limit to temperature below which nothing can exist; there can be no temperature below 0 K because by definition the system is said to be usable to lose any more heat energy

pH

a measure of the hydrogen ion content of an aqueous solution, defined as the negative log of the H+ (H3O+) concentration; for pure water at equilibrium and 298K the concentration of hydrogen ions equals the concentration of hydroxide ions (10^-7 M) therefore pure water at 298K has a pH of 7 and a pOH of 7 (-log10^-7 = 7); by taking the negative logarithm of the entire water dissociation constant expression ([H3O+][OH-] = 10^-14) we find pH + pOH = 14

pOH

a measure of the hydroxide (OH-) ion content of an aqueous solution, defined as the negative log of the OH- concentration; for pure water at equilibrium and 298K the concentration of hydrogen ions equals the concentration of hydroxide ions (10^-7 M) therefore pure water at 298K has a pH of 7 and a pOH of 7 (-log10^-7 = 7); by taking the negative logarithm of the entire water dissociation constant expression ([H3O+][OH-] = 10^-14) we find pH + pOH = 14 (this comes from for an aqueous solution at 298K a pH less than 7 (or pOH greater than 7) indicates a relative excess of hydrogen ions, and the solution is acidic; a pH greater than 7 (or pOH less than 7) indicates a relative excess of hydroxide ions and the solution is basic; a pH equal to 7 indicates an equal concentration of hydrogen and hydroxide ions resulting in a neutral solution

Lewis Structure

a method of representing the shared and unshared electrons of an atom, molecule, or ion; also called a Lewis dot diagram; it is the chemical symbol of an element surrounded by dots, each representing one of the s or p valence electrons of the atom; example Carbon has four valence electrons so it is resented by C surrounded by a dot above, below, and on both sides; Lithium has one valence electrons so it is Li with 1 dot; when drawing lewis dot structures remember that some atoms can expand their octets by utilizing the d-orbitals in their outer shell, this will only take place with atoms in period 3 or greater; Rules for drawing lewis structure: 1. Draw out the backbone of the compound (arrangement of atoms), in general the least electronegative atoms is the central atom; hydrogen (always) and the halogens F, Cl, Br, and I (usually) occupy a terminal position; ex in HCN, H must occupy and end position, of the remaining two atoms, C is the least electronegative and, therefore, occupies the central position. Therefore, the skeletal structures is as follows: H-C-N 2. Count all the valence electrons of the atoms. The number of valence electrons of the molecule is the sum of the valence electrons of all the atoms present; using HCN as an example H has 1 valence electron, C has four, N has 5, therefore HCN has 10 valence electrons; 3. Draw single bonds between the central atoms and the atoms surrounding it. Each single bond corresponds to a pair of electrons; ex H : C : N 4. Complete the octets of all atoms bonded to the central atom, using the remaining valence electrons left to be assigned. Recall that H is an exception to the octet rule because it can only have two valence electrons. In this example, H already has two valence electrons from its bond with C; ex H : C : N : with :: below and above N 5. Place any extra electrons on the central atom. If the central atom has less than an octet, try to write double or triple bonds between the central and surrounding atoms using the long pairs on the surrounding atoms; ex HCN does not satisfy the octet rule for C because C only has four valence electrons, therefore, two lone electron pairs from the N atoms must be moved to form two more bonds with C, creating a triple bond between C and N so you get H - C -= N: ; now the octet rule is satisfied for all three atoms; C and N have eight valence electrons, and H has two valence electrons Formal charge is the difference between the number of electrons assigned to an atom in a lewis structure and the number of electrons normally found in that atom's valence shell; an equation to calculate formal charge is: formal charge = V - Nnonbonding - 1/2 Nbonding where V is the normal number of valence electrons, Nnonbonding is the number of nonbonding electrons, and Nbonding is the number of bonding electrons (double the number of bonds because each bond has two electrons). The charge of an ion or compound is equal to the sum of the formal charge fo the individual atoms comprising the ion or compound; another way to remember this equation is formal charge = valence electrons - dots - sticks brief note: the difference between formal charge and oxidation number is that the formal charge underestimates the effect of electronegativity differences whereas oxidation numbers overestimate the effect of electronegativity differences; assuming that the more electronegative atom has a 100% share of the bonding electrons pair; for example in a molecule of CO2, the formal charge on each of the atoms is 0, but the oxidation number of each fo the oxygen atoms is -2 and the carbon is +4. in reality the distribution of electron density between the carbon and oxygen atoms lies somewhere between the extremes predicted by the formal charge and the oxidation states

mole

a mole is a quantity of any substance (atoms, molecules, dollar bills, kitten, anything) equal to the number of particles that are found in 12 grams of carbon-12; this number of particles is defined as Avogadro's number (NA 6.022 x 10^23 mol-1); the mass of 1 mole of substance in grams is the same as the mass of one molecule or atom in atomic mass units (amu); for example H2CO3 (carbonic acid) has a mass of 62 amu; one mole of the compound has a mass of 62 grams

Ligand

a molecule bonded to a metal ion in a coordination compound; ligands are lewis bases that form coordinate covalent bonds with the central metal ion

Polar molecule

a molecule possessing one or more polar covalent bonds and a geometry that allows the bond dipole moments to sum to a det dipole moment

nonpolar molecule

a molecule that exhibits no net separation of charge and, therefore, no net dipole movement

Quantum Number

a number used to describe the energy levels in which electrons reside; all electrons in an element are described by a unique set of four quantum numbers; the values of quantum numbers qualitatively give information about the size, shape, and orientation of the orbitals

Radioactivity

a phenomenon exhibited by certain unstable isotopes in which they undergo spontaneous nuclear transformation via emission of one or more particles

Malleability

a physical property of metals that defines how well an element can be shaped using a hammer

Mass

a physical property representing the amount of matter in a given sample

Titration curve

a plot of the pH of a solution vs the volume of acid ro base added in an acid-base titration, or a plot of the electromotive force of a solution vs the volume of oxidizing or reducing agent added in an oxidation-reduction titration

Mixed-order reaction

a reaction in which the reaction order changes over time in the rate law; sometimes refer to non-integer orders (fractions); fractions are more specifically described as broken order; in recent times, the term mixed order has come to refer solely to reactions that change order over time; an example: rate = (k1 [C] [A]^2)/k2 + k3 [A]) where A represents a single reactant and C a catalyst The result of the large value for [A] at the beginning of the reaction is that k3 [A] >> k2, a reaction will appear to be first order with respect to A; at the end of the reaction when [A] is low, k2>> k3[A]; making the reaction appear second order with respect to A;

Phase Diagram

a plot, usually of pressure vs. temperature, showing which phases of a compound will exist under any set of conditions; graphs that show the standard and nonstandard states of matter for a given substance in an isolated system; lines on the phase diagram are called the lines of equilibrium or the phase boundaries and indicate the temperature and pressure values for the equilibria between phases; the lines of equilibrium divide the diagram into three regions corresponding to the three phase - solid, liquid and gas - and they themselves represent the phase transformations; the triple point is the point at which the three phase boundaries meet, this is where the three phases exist in equilibrium; the critical point is the point where the phase boundary between the liquid and gas phases is terminated, this is the temperature and pressure above which there is no distinction between the phases, although this may seem to be an impossibility after all its always possible to distinguish between the liquid and solid phase such as supercritical fluids are perfectly logical, as a liquid is heated in a closed system its density decreases and the density of the vapor sitting above it increases, the critical point is the temperature and pressure at which the two densities become equal and there is no distinction between the two phases, the heat of vaporization at this point and for all temperatures and pressures above the critical point values is zero

Lyman series

a portion of the emission spectrum from hydrogen representing electronic transitions from energy levels n> 1 to n=1

Isothermal process

a process that occurs at constant temperature; constant pressure implies that the total internal energy of the system (U) is constant throughout the process, this is because temperature and internal energy are directly proportional; when U is constant, delta U = 0 and the first law simplifies to Q = W (the heat added to the system equals the work done by the system); an isothermal process appears as a hyperbolic curve on a pressure-volume graph

nonspontaneous process

a process that will not occur on its own without energy input from the surroundings; has a positive change in free energy

spontaneous process

a process that will occur on its own without energy input from the surroundings; defined by a negative change in free energy; calculating the change in Gibbs free energy (delta G) for a process, such as chemical reaction, allows us to predict whether the process will be spontaneous or nonspontaneous; spontaneous reactions will not necessarily happen quickly and may not go to completion; many spontaneous reactions have very high activation energies and therefore rarely take place; for example a match will not ignite itself however providing a quantity of thermal energy (generated by the friction associated with striking the match) that equals or exceeds the activation energy will allow the match to light and burn spontaneously, at this point the combustion of the chemical components of the match using molecular oxygen in the air will not need any additional external energy once the activation energy has been supplied; some spontaneous reactions proceed very slowly, the role of enzymes is to selectively enhance the rate of certain spontaneous (but slow) chemical reactions so that the biologically necessary products can be formed at a rate sufficient for sustaining life, some reactions do not go to completion but settle into a low-energy state called equilibrium; spontaneous reactions may go to completion, but many simply reach equilibrium with dynamically stable concentrations of reactants and products; a common method for supplying energy for nonspontaneous reactions is by coupling nonspontaneous reactions to spontaneous ones

Physical Property

a property of a substance unrelated to it chemical behavior, such as melting point, boiling point, density, or odor

semipermeable

a quality of a membrane allowing only some components of a solution to pass through, usually including the solvent, while limiting the passage of other species

Net ionic equation

a reaction equation showing only the species actually participating in the reaction; this equation does not include spectator ions; for some examples use the half reaction to determine the net ionic equation

single displacement reaction

a reaction in which an ion of own compound is replaced with another ion; also known as a single replacement reaction; for example solid copper metal will displace silver ions in a clear solution of silver nitrate to form a blur copper nitrate solution and solid silver metal Cu (s) + AgNO3 (aq) ---> Ag (s) + CuNO3 (aq); these are often classified further as oxidation reduction reactions

zero-order reaction

a reaction in which the concentrations of reactants have no effect on the overall rate; these reactions have a constant reaction rate equal to the rate constant (rate coefficient), k. The rate law for a zero order reaction is Rate = k [A]^0 [B]^0 = k where k has units of M/s Remember that the rate constant itself is dependent on temperature; thus it is possible to change the rate for a zero order reaction by changing the temperature; the only other way to change the rate of a zero order reaction is by addition of a catalyst which lowers the activation energy, thereby increasing the value of k; plotting a zero order reaction on a concentration vs time results in a linear graph; this line shows that the rate of formation of product is independent of the concentration of reactant; the slope of such a line is the opposite of the rate constant k; temperature and the addition of a catalyst are the only factors that can change the rate of a zero-order reaction; zero order is unaffected by doubling reactant concentrations

second order reaction

a reaction in which the rate is directly proportional to the concentration of two reactants, or to the square of one single reactant; example: rate = k [A]^1 [B]^1 or rate = k [A]^2 or rate = k [B]^2 it is important to recognize that a second order rate law often suggests a physical collision between two reactant molecules, especially if the rate law is first order with respect to each of the two reactants; plotting a reaction that is second order with respect to a single reactant on a concentration vs time curve results in a nonlinear graph, this curve shows that the rate of formation of product is dependent on the concentration of reactant; plotting 1/[A] vs time reveals a linear curve; the slope of such a curve is equal to the rate constant k; k=slope; changing temperature, changing reactant concentrations, and adding catalyst all change the rate of second order reactions; lowering temperature decreased the rate, doubling concentrations quadruples the rate, and adding catalyst increases the rate

Oxidation-reduction (redox) reaction

a reaction that involves the transfer of electrons from one chemical species to another; oxidation is the loss of electrons and reduction is the gain of electrons (mnemonic: LEO GER: Loss of Electrons is Oxidation, Gain of Electrons is Reduction

Nickel-cadmium battery

a rechargeable electrochemical call in which the anode is the side of Cd metal oxidation and the cathode is the site of Ni2+ ion reduction

Nickel-metal hydride battrey

a rechargeable electrochemical cell in which the anode is the side of metal hydride oxidation and the cathode is the site of nickel ion reduction; the nickel may be in one of many oxidation states

Orbital

a region of electron density around an atom or molecule containing no more than two electrons of opposite spin;

Octet rule

a rule stating that bonded atoms tend to undergo reactions that will produce a complete octet of valence electrons forming a stable electron configuration similar to that of the noble gases; applies without exception only to C, N, O, and F, sodium and magnesium also almost always abide by this rule; this rule is hardly a rule at all due to there being many exceptions; it is noted that elements, especially ones in biological roles, tend to be most stable with eight electors in their valence shell; exceptions to this rule are hydrogen which can only have two valence electrons, lithium and beryllium which bond to attain two and four valence electrons respectively, boron which bonds to attain six valence electrons; and all elements in period 3 and greater which can expand the valence shell to include more than eight electrons by incorporating d-orbitals; for example in certain compounds chlorine can form seven covalent bonds, thereby holding 14 electrons in its valence shell; rules to remember this: incomplete octet (elements stable with fewer than 8 electrons in valence shell include hydrogen, helium, lithium, beryllium, and boron) expanded octet (any element in period 3 and greater can hold over 8 electrons due to adding them to d-orbital) and odd numbers of electrons (any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom; for example nitric oxide (NO) has eleven valence electrons)

atomic emission spectrum

a set of frequencies of electromagnetic waves given off by atoms of an element; consists of a series of fine lines of individual colors; every element has its own characteristic emission spectrum; sometimes the electromagnetic energy emitted corresponds to a frequency in the visible light range

Ionic Solid

a solid consisting of positive and negative ions arranged into crystals that are made up of regularly repeated units held together by ionic bonds

saturated solution

a solution containing the maximum amount of solute that can be dissolved in a particular solvent at a given temperature

Unsaturated solution

a solution in which more solute may be dissolved before reaching saturation

Titrant

a solution of known concentration that is slowly added to a solution of unknown concentration to determine its concentration

Titrand

a solution of unknown concentration to which a solution of known concentration is added to determine its concentration

supersaturated

a solution which is beyond equilibrium, where ion product is greater than the solubility product constant; these are thermodynamically unstable; it is possible to create this by dissolving solute into a hot solvent and then slowly cooling the solution

Lewis acid

a species capable of accepting an electron pair; the difference between lewis acid and bronsted acid and bases follow the exchange of hydrogen ion (H+), which is essentially a naked proton, in the lewis definition the focus on the reaction is no longer on the proton but instead the electrons forming the coordinate covalent bond

Lewis base

a species capable of donating an electron pair; the difference between lewis acid and bronsted acid and bases follow the exchange of hydrogen ion (H+), which is essentially a naked proton, in the lewis definition the focus on the reaction is no longer on the proton but instead the electrons forming the coordinate covalent bond

Redox Titration

a specific method used to determine the concentration of an unknown solution using reducible titrants or titrands, typically by measuring voltage changes

neutron

a subatomic particle contained within the nucleus of an atom; carries no charge and has a mass slightly larger than that of a proton

Proton (p+)

a substance particle that carries a single positive charge and has a mass slightly less than 1 amu

supercritical fluid

a substance whose current state is simultaneously a liquid and a gas-there is no distinction between the two phases

surge current

an above average current transiently released at the beginning of the discharge phase of a battery

mixture

a system containing multiple substances (2+) that have been physically combined but are not chemically combined; gases "dissolved" into other gases can be thought of as solutions however they are more properly defined as mixtures because the gas molecules do not interact all that much chemically; as a point of clarification, all solutions are considered mixtures, but not all mixtures are considered solutions

Isolated system

a system that can exchange neither matter nor energy (heat or work) with its surroundings; ex. an insulated bomb calorimeter

valence shell electron pair repulsion (VSEPR) theory

a system that reflects the geometric arrangement of a molecule based on its lewis dot structure; three dimensional structure is determined by the repulsions between bonding and nonbonding electron pairs in the valence shells of atoms; these electron pairs arrange themselves as far apart as possible, thereby minimizing repulsion forces. Steps to predict geometrical structure of a molecule using VSEPR theory: 1. Draw the lewis dot structure of the molecule 2. Count the total number of bonding and nonbonding electron pairs in the valence shell of the central atom 3. Arrange the electron pairs around the central atom so that they are as far apart as possible; example AX2 has a lewis structure X : A : X, the A atom has two bonding electron pairs in its valence shell, to position these electron pairs as far apart as possible their geometric structure should be linear of note electronic geometry is different from molecular geometry, example HN3 (ammonia) has a tetrahedral electronic structure but is considered to have a molecule structure that is trigonal pyramidal

Kelvin (K)

a temperature scale with units equal to the units of the Celsius scale and absolute zero defines as 0 K; also called the absolute temperature scale

sphygmomanometer

a tool for measuring blood pressure

Mole fraction (X)

a unit of concentration equal to the ratio of the number of moles of a particular component to the total number of moles for all species in the system equation: XsubA = moles of A/total moles of all species the sum of the mole fractions in a system will always equal 1; the mole fraction is used to calculate the vapor pressure depression of a solution, as well as the partial pressures of gases in a system

Millimeters of Mercury (mmHg)

a unit of pressure defined as the number of millimeters that mercury in a barometer is raised above its surface in a capillary tube by an external pressure; 1 torr is equal to 1 mmHg by definition, and 1 atmosphere is equal to 760 mmHg

Complex ion

also known as coordination compound; A polyatomic molecule in which a central cation is bonded to electron pair donors called ligands; complexes are help together with coordinate covalent bonds, in which an electron pair donor (a lewis base) and an electron pair acceptor (a lewis acid) form very stable lewis acid-base adducts; many active sites for proteins utilize complex ion binding and transition metal complexes to carry out their function; one classic example is the ion cation in hemoglobin, which can carry oxygen, carbon dioxide, and carbon monoxide as ligands; physical and chemical properties of complex ions are diverse, including an wide range of solubilities and varied chemical reactions; inorganic complex ions are often tend to have vibrant, distinctive colors; in some complexes the central cation can be bonded to the same ligand in multiple places, this is known as chelation

Ionization Energy

also known as ionization potential; The energy required to remove an electron from the valence shell of a gaseous atom; removing an electron from an atom always requires an input of heat, which makes it an endothermic process; the greater the atoms zeff or the closer the valence electrons are to the nucleus the more tightly bound they are, this makes it more difficult to remove one or more electrons increasing the ionization energy, thus ionization energy increasing from left to right across the periodic table and from bottom to top in a group; the subsequent removal of a second or third electron requires increasing amount of energy because the removal of more than one electron means that the electrons are being removed from an increasingly cationic (positive) species; the energy required to remove the first electron is called the first ionization energy, the energy required to remove the second electron is called the second ionization energy; elements in group IA and IIA such as lithium and beryllium have such low ionization energies that they are called the active metals; the active metals do not exist naturally in their neutral forms, they are always found in ionic compound, minerals, or ores; the loss of one electron from the alkali metals (group IA) or the loss of two electrons in the alkaline earth metals (group IIA) results in the formation of a stable, filled valence shell; if losing a certain number of electrons gives an element a noble gas like electron configuration, then removing a subsequent electron will cost much more energy, for example Mg2+ (g) --> Mg3+ (g) +e- third ionization energy = 7730 kj/mol, first ionization energy=738 and second ionization energy = 1450; group VIIIA elements, or noble or inert gases, are the least likely to give up electrons; they already have a stable electron configuration and are unwilling to disrupt that stability by giving up an electron; therefore noble gases are among the elements with the highest ionization energies

Double-displacement reaction

also known as metathesis reaction; A reaction in which ions from two different compounds swap their associated counter ions; this type of reaction occurs when one fo the products is removed from the solution as a precipitate or gas or when two of the original species combine to form a weak electrolyte that remains undissociated in solution; for example when solutions of calcium chloride and silver nitrate are combined, insoluble silver chloride forms in a solution of calcium nitrate; CaCl2 (aq) + 2 AgNO3 (aq) --> Ca(NO3)2 (aq) + 2 AgCl (s)

Inert Gases

also known as noble gases; the elements in Group VIIIA, which contain a full octet of valence electrons in their outermost shells and are therefore very unreactive; they have high ionization energies, little or no tendency to gain or lose electrons and (for He, Ne, and Ar at least) no measurable electronegative; noble gases have extremely low boiling points and exist as gases at room temperature; noble gases have found a commercial niche as lighting due to their lack of reactivity

strong acid

an acid that undergoes complete dissociation in an aqueous solution; for example when sodium hydroxide is added to water, the ionic compound dissociates according to the net ionic equation: NaOH (s) --> Na + (aq) + OH- (aq) hence, in a 1 M NaOH solution, complete dissociation such as NaOH, will remain in the solution; this is why the dissociation of strong acids and bases is said to go to completion; the contribution of OH- and H+ ions from the auto-ionization of water is negligible if the concentration of the acid or base is significantly greater than 10^-7 M; on the other hand if the concentration of acid or base is close to 10^-7 M, then the contribution form the auto-ionization of water is important; strong acids commonly seen on the MCAT include HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), H2SO4 (sulfuric acid), HNO3 (nitric acid) and HClO4 (perchloric acid); one thing that is important to note for acid strength is the effect of induction; electronegative elements positioned near an acidic proton increase acid strength by pulling electron density out of the bond holding the acidic proton; this weakens proton bonding and facilitates dissociation; thus acids that have electronegative elements nearer to acidic hydrogens are stronger than those that do not

Weak acid

an acid that undergoes partial dissociation in an aqueous solution; a weak monoprotic acid, HA, will dissociate partially in water to achieve an equilibrium state: HA (aq) + H2O (l) <--> H3O+ (aq) + A- (aq) because the system exists in an equilibrium state, we can write the dissociate equation to determine the acid dissociation constant (Ka) as: Ka= [H3O+][A-]/[HA] water is not included because it is a pure liquid; generally we can characterize a species as a weak acid if its Ka is less than 1.0

neutral solutions

an aqueous solution in which the concentration of H+ and OH- ions are equal (pH= 7 at 298 K)

Lead-acid battery

an electrochemical cell in which the anode is the side of Pb metal oxidation and the cathode is the site of Pb4+ ion reduction; the electrolyte is a strong acid, usually sulfuric acid

Rechargeable Battery

an electrochemical cell that can undergo a reversible oxidation-reduction process; when discharging, it functions as a galvanic (voltaic) cell, and when charging, it functions as an electrolytic cell

valence electron

an electron in the highest occupied energy level of an atom; these are more likely to become involved in bonds with other atoms because they experience the least electrostatic pull from their own nucleus; the tendency of a given valence electron to be retained or lost determines the chemical properties of an element

Metalloid

an element possessing properties intermediate between those of a metal and those of a nonmetal; also called a semimetal; these separate the metals and nonmetals in a stair step group of elements on the period table; their physical properties -densities, melting points, and boiling points- vary widely and can be combinations of metallic and nonmetallic characteristics; ex. silicon (Si) has a metallic luster but is brittle and a poor conductor; the reactivates of these are dependent on the elements with which they are reacting; ex. Boron (B) behaves like a nonmetal when reacting with sodium (Na) and like a metal when reacting with fluorine (F); these include B, Si, Ge, As, Sb, Te, Po, and At;

nernst equation

an equation that relates the voltage of an electrochemical cell to the concentrations of the reactants and products within that cell

Precipitate

an insoluble solid that separates from a solutions; generally the results of mixing two or more solutions or of a temperature change

sparingly soluble salt

an ionic compound that has a low solubility at a given temperature

salt

an ionic substance consisting of cations and anions

Transition Metal

any of the elements in the B groups of the periodic table (groups IB to VIIIB or groups 3-12); these elements have two or more oxidation states (charges when forming bonds with other atoms); because the valence electrons of all metals are only loosely held to their atoms, they are free to move, which makes metals good conductors of heat and electricity; the valence electrons of the active metals are founds in the s subshell; those of the transition metals are found in the s and d subshells; and those of the lanthanide and actinide series elements are in the s and f subshells; some transition metals -copper, nickel, silver, gold, palladium, and platinum- are relatively nonreactive, a property that makes them ideal for the production of coins and jewelry; one of the unique properties of the transition metals is that many of them can have different possible charged forms or oxidation states because they are capable of losing different numbers of electrons from the s- and d- orbitals in their valence shells; for example copper (cu) can exist in either the +1 or the +2 oxidation state, and manganese (Mn) can exist in the +2, +3, +4, +6, or +7 oxidation state; because of this ability to attain different positive oxidation states, transition metals form many different ionic compounds; these different oxidation states often correspond to different colors; solutions with transitional metals-containing complexes are often vibrant; these complex ions tend to associate in solutions either with molecules of water (known as hydration complexes, such as CuSO4 * 5H2O) or with nonmetals (such as {Co(NH3)6}Cl3). This ability to form complexes contributes to the variable solubility of certain transition metal containing compounds; for example AgCl is insoluble in water but quite soluble in aqueous ammonia due to the formation of the complex ion {Ag(NH3)2}+; the formation of complexes causes the d-orbitals to split into two energy sublevels; this enables many of the complexes to absorb certain frequencies of light- those containing the precise amount of energy required to raise electrons from the lower to higher energy d-orbitals. The frequencies not absorbed (known as the subtraction frequencies) give the complexes their characteristic colors; colors are seen as the color that is not absorbed but reflected off the object; our brain mixes these subtraction frequencies and we perceive the complementary color of the frequency that was absorbed; this is best illustrated with an example, carotene is a photosynthetic pigment that strongly absorbs blue light but reflexes other colors, thus our brains interpret the color of carotene as the result of white light minus blue light, which is yellow, the complementary colors are shown in a color wheel

Representative elements

elements in Groups 1,2, and 13 through 18 in the modern IUPAC table (the s- and p-blocks of the table, also called A group elements); these elements tend to have valence shells that follow the octet rule

nonrepresentative element

elements with an expanded valence shell that includes d- and f-block electrons; also called group B or transition elements

Limiting reagent

in a chemical reaction, the reactant presents in such quantity as to limit the amount of product that can be formed; this occurs because the one reactant is used up or consumed first which limits the amount of product that can be formed in the reaction; the reactants that remain after all the limiting reagent is used up are called excess reagents; for problems involving the determination of the limiting reagent, keep in mind two principles: 1. All comparisons of reactant must be done in units of moles. Gram to gram comparisons will be useless and may even be misleading 2. it is not the absolute mole quantities of the reactants that determine which reactant is limiting reagent. Rather, the rate at which the reactants are consumed (the stoichiometric ratios of the reactants), combined with the absolute mole quantities determine which reactant is the limiting reagent

parallel spin

in quantum mechanics, electrons in different orbitals fo an atom with the same ms values

London dispersion forces

intermolecular forces arising from interactions between temporary dipoles in molecules; the bonding electrons in nonpolar covalent bonds may appear to be shared equally between two atoms, but at any point in time, they will be located randomly throughout the orbital; in a given moment the electron density may be unequally distributed between the two atoms; this results in a rapid polarization and couterpolarization of the electron cloud and the formation of short lived dipole moments; subsequently these dipoles interact with the electron clouds of neighboring compounds, inducing the formation of dipoles; the momentarily negative end of one molecule will cause the closest region in any neighboring molecule to become temporarily positive itself; this causes the other end of the neighboring molecule to become temporarily negative, which in turn induces other molecules to become temporarily polarized and the cycle beings again; the attractive or repulsive interactions of these short lived and rapidly shifting dipoles are known as London dispersion forces or van der Waals force; dispersion forces are the weakest of all of the intermolecular interactions because they are the result of induced dipoles that change and shift moment to moment; they do not extend over long distances and are significant only when molecules are in close proximity; the strength of London force also depends on the degree and ease by which the molecules can be polarized, that is, how easily the electrons can be shifted around; large molecules are more easily polarizable than comparable smaller molecules and thus posses greater dispersion forces; despite the weak nature of these interactions they are important, if it weren't for them the noble gases would not liquefy at any temperature because no other intermolecular forces exist between the noble gas atoms; the low temperatures at which noble gases liquefy are indicative of the very small magnitude of the dispersion forces between the atoms

-ite and -ate

many polyatomic anions contain oxygen and are therefore called oxyanions; when an element forms two oxyanions, the name of the one with less oxygen ends in -ite, and the one with more oxygen ends in -ate; example NO2- Nitrite, NO3- is Nitrate, SO3 2- is sulfide, SO4 2- is sulfate; mnemonic: the l-ITE-st anions have the fewest oxygen; the heaviest anions ATE the most oxygen

-ide

monatomic anions are named by dropping the ending of the name of the element and adding -ide; example H- is hydride, F- is fluoride, O2- is oxide, S2- is sulfide, N3- is nitride, P3- is phosphide

hydrogen - or dihydrogen -

polyatomic anions often gain one or more H+ ions to form anions of lower charge. The resulting ions are named by adding the word hydrogen or dihydrogen to the front of the anion's name. An older method uses the prefix bi- to indicate the addition of a single hydrogen ion; example HCO3 - is hydrogen carbonate or bicarbonate, HSO4 - is hydrogen sulfate or bisulfate, H2PO4 - is dihydrogen phosphate

Phase Change

reversible transition between solid, liquid, and/or gas phase caused by shifts in temperature or pressure; for example at 0 C and 1 atm in an isolated system, ice and water exist in an equilibrium, in other words some of the ice may absorb heat (From the liquid water) and melt, but because that heat is being removed from the liquid water, and equal amount of the liquid water will freeze and form ice, thus the relate amounts of ice and water remain constant;

Oxidation potential

the ability of a substance to be spontaneously oxidized; a more positive oxidation potential (measured in volts) is indicative of a substance that is easier to oxidize and will therefore more likely act as an anode in an electrochemical cell

Reduction potential

the ability of a substance to be spontaneously reduced; a more positive reduction potential (measured in volts) is indicative of a substance that is easier to reduce and will therefore more likely act as a cathode in an electrochemical cell

oxidizing agent

the atom that facilitates the oxidation of another species; the oxidizing agent gains electrons and is thereby reduced; common oxidizing agents include: O2, H2O2, F2, Cl2, Br2, I2 (halogens), H2SO4, HNO3, NaClO, KmnO4, CrO3, Na2Cr2O7, pyridinium chlorochromate (PCC), NAD+, FADH

Ionic Radius

the average distance from the center of the nucleus to the edge of its electron cloud; cationic radii are generally smaller than their parent metal, whereas anionic radii are generally larger than their parent nonmetal; in order to understand ionic radii, we must make two generalizations, one is that metals lose electrons and become positive, while non-metals gain electrons and become negative, the other is that metalloids can o in either direction, but tend to follow the trend based on which side of the metalloid line they fall on, thus Si behaves like a nonmetal while Ge tends to act like a metal; for nonmetals close to the metalloid line, their group number indicates that they require more electrons than nonmetals to achieve the electronic configuration seen in Group VIIIA (Group 18); these nonmetals gain electrons while their nuclei maintain the same charge; therefore, these nonmetals close ot the metalloid line possess a larger ionic radius than their counterparts closer to Group VIIIA; for metals the trend is similar but opposite, metals closer to teh metalloid line have more electrons to lose to achieve the electronic configuration see in group VIIIA; because fo this the ionic radius of metals near the metalloid line is dramatically smaller than that of other metals; metals closer to group IA have fewer electrons to lose and therefore experience a less drastic reduction in radius during ionization; these changes a

Root-mean-square speed (urms)

the average speed of a gas molecule at a given temperature; as a scalar, it does not take direction into account; usubrms = square root 3RT/M, where R is the ideal gas constant, T is temp and M is molar mass; for the equation make sure to convert molar mass to kg/mol due to Joules also being derived from kilograms

solvation

the electrostatic interaction between solutes and solvent molecules; also called dissolution; the term hydration can be used when water is the solvent; solvation involves breaking intermolecular interactions between solute molecules and between solvent molecules and forming new intermolecular interactions between solute and solvent molecules together; when the new interactions are stronger than the original ones, solvation is exothermic, and the process is favored at low temperatures; when the new interactions are waker than the original ones, solvation is endothermic and the process is favored at high temperatures; most dissolutions are of this type; the second property besides enthalpy change is entropy change that occurs in the process; at constant temperature and pressure, entropy always increases upon dissolution; as any process the spontaneity of dissolution depends on the change in Gibbs free energy: spontaneous processes are associated with a decrease in free energy, while non-spontaneous processes are associated with an increase in free energy; thus, whether or not dissolution will happen spontaneously depends on both the change in enthalpy and the change in entropy for the solute and solvent of the system

standard heat of combustion

the enthalpy change associated with the combustion of a fuel; because measurements of enthalpy change require a reaction to be spontaneous and fast, combustion reactions are the ideal process for such measurements; most combustion reactions presented on the MCAT occur in the presence of atmospheric oxygen, but keep in mind that there are other combustion reactions in which oxygen is not the oxidant; diatomic fluorine can be used as an oxidant; hydrogen gas will combust with chlorine gas to form gaseous hydrochloric acid and in the process will evolve a large amount of heat and light as it characteristic of combustion reactions; the reactions listed in the CH4 (g) example shown earlier are combustion reactions with O2 (g) as the oxidant; therefore the enthalpy change listed for each of the three reactions is the delta H comb for each of the reactions; the glycolytic pathway is also a combustion reaction that utilizes a fuel (glucose) mixed with an oxidant (oxygen) to produce carbon dioxide and water C6H12O6 + 6O2 --> 6CO22 + 6H2O; key concept: the larger the alkane reactant, the more numerous the combustion products

water dissociation constant (Kw)

the equilibrium constant of the water dissociation reaction at a given temperature; equal to 10^-14 at 25 deg C (298) K; each mole of water that autoionizes produces one mole each of hydrogen and hydroxide ions, so the concentration of the hydrogen ions and hydroxide ions are always equal in pure water at equilibrium, thus the concentration of each of the ions in pure water at equilibrium at 298K is 10^-7 M, they will only be equal when the water is neutral; it is important to note that Kw is an equilibrium constant: unless the temperature of the water is changes, the value of Kw cannot be changed; therefore isolated changes in concentration, pressure, or volume will not affect Kw

Principal quantum number (n)

the first quantum number, which defines the energy level or shell occupied by an electron; the larger the integer value of n, the higher the energy level and radius of the electrons shell; within each shell, there is a capacity to hold a certain number of electrons given by 2n^2; the difference in energy level between two shells decreased as the distance from the nucleus increases because the energy difference is a function of [(1/ni^2)-(1/nf^2)]; for example n=3 to n=4 is less than the energy difference between the n=1 and n=2 shells

Photon

the form of light which displays particulate and quantal behavior

Structural formula

the graphic representation of a molecule depicting how its atoms are arranged

second law of thermodynamics

the law stating that all spontaneous processes lead to an increase in the entropy of the universe; energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so; notice this law states that energy will spontaneously disperse, it does not say that energy can never be localized or concentrated; however the concentration of energy will rarely happen spontaneously in a closed system; work usually must be done to concentrate energy; for example refrigerators work against the direction of spontaneous heat flow (that is, they counteract the flow of heat from the "warm" exterior of the refrigerator to the "cool" interior) thereby "concentrating" energy outside of the system in the surroundings; as a result, refrigerators consume a lot of energy to accomplish this movement of energy against the temperature gradient; this saw is described as time's arrow because there is a unidirectional limitation on the movement of energy by which we recognize before and after or new and old; for example, you would instantly recognize whether a video recording of an explosion was running forward or backward

Law of constant composition

the law stating that the elements in a pure compound are found in specific mass ratios; for example every sample of water will contain two hydrogen atoms for every one oxygen atom, or in term of mass for every one gram of hydrogen there will be eight grams of oxygen

Law of conservation of charge

the law stating that, in a given reaction, the charge of ions in the products is equal to the charge of ions in the reactants

Law of conservation of mass

the law stating that, in a given reaction, the mass of products is equal to the mass of the reactants

Molar Mass

the mass of one mole of a compound; usually expressed g/mol; to find moles of a sample use equation moles = mass of sample (g)/ molar mass (g/mol); example how many moles are in 9.53 g of MgCl2, first find molar mass of MgCl2 = 95.3 g/mol, now solve for moles 9.53 g/95.3g/mol = 0.1 mol MgCl2; we can take the ideal gas law and arrange it to find molar mass as well by M= (psubSTP)(22.4L/mol)

system

the matter and energy under consideration; the total amount of reactants and products in a chemical reaction; it could be the amount of solute and solvent used to create a solution; it could be the gas inside a balloon; the surroundings or environment are everything outside the system; the boundary between system and environment can be moved, it is not permanently fixed

Osmosis

the movement of water through a semipermeable membrane down its concentration gradient, from low solute concentration to high solute concentration

oxidation number

the number assigned to an atom in an ion or molecule that denotes its real or hypothetical charge, assuming that the most electronegative element in a bond is awarded all of the electrons in that bond; also called oxidation state oxidation number of an atom in a compound is assigned according to the following rules: 1. the oxidation number of a free element is zero, for example, the atoms in N2, P4, S8 and He all have oxidation numbers of zero 2. the oxidation number for a monatomic ion is equal to the charge of the ion, for example, the oxidation number for Na+ = +1, Cu 2+ = +2, Fe 3+ = +3, Cl- = -1, and N 3- = -3 3. The oxidation number of each Group IA element in a compound is +1 4. The oxidation number of each Group IIA element in a compound is +2 5. The oxidation number of each Group VIIA element in a compound is -1, except when combined with an element of higher electronegativity, for example, in HCl, the oxidation number of Cl is -1; in HOCl however the oxidation number of Cl is +1 6. The oxidation number of hydrogen is usually +1, however its oxidation number is -1 in compounds with less electronegative elements (Groups IA and IIA). Hydrogen is +1 in HCl, but -1 in NaH (of note the conventions of formula writing put cations first and anion second, thus HCL implies H+, and NaH implies H-) 7. In most compounds, the oxidation number of oxygen is -2, the two exceptions are peroxides (O2 2-) for which the charge on each oxygen is -1, and compounds with more electronegative elements such as OF2, in which oxygen has a +2 charge 8. The sum of the oxidation numbers of all the atoms present in a neutral compound is zero, the sum of the oxidation numbers of the atoms present in a polyatomic ion is equal to the charge of the ion, thus for (SO4 2-) the sum of the oxidation numbers must be -2

van't Hoff factor

the number of particles into which a compound dissociates in solution; for example i = 2 for NaCl because each formula unit of sodium chloride dissociates into two particles - a sodium ion and a chloride ion - when it dissolves; covalent molecules such as glucose do not readily dissociate in water and thus have i values of 1

Le Chatelier's Principle

the observation that when a system at equilibrium is disturbed or stressed, the system will react in such a way as to relieve the stress and restore equilibrium; the bicarbonate buffer system is a classic ex of Le Chatelier's princliple applied to physiology: CO2 (g) + H2O (l) <--> H2CO3 (aq) <--> H+ (aq) + HCO3- (aq); Pressure changes: only chemical reactions that involve at least one gaseous species will be affected by change in the system's pressure and volume; when system is compressed its volume decreases and its total pressure increases, this increase in total pressure is associated with an increase in the partial pressure of each gas in the system, and this results in the system no longer being in the equilibrium state such that Qp does not equal Keq; The system will move forward or in reverse; always toward whichever side has the lower total number of moles of gas, this is a consequence of the ideal gas law, which tells us that there is a direct relationship between the number of moles of gas and the pressure of the gas; if one increases the pressure of a system it will respond by decreasing the total number of gas moles, thereby decreasing the pressure; EX. N2 (g) + 3H2 (g) <--> 2NH3 (g), if pressure is increased it will proceed to the side with fewer moles of gas which would be the 2NH3, if pressure is decreased it will proceed in the side that produces more moles which would be N2 + 3H2 Temperature Changes: unlike the effect of changing concentrations or pressures, the result of changing temperature is not a change in the reaction quotient Qc or Qp, but a change in Keq; change in temperature does not cause the concentrations or partial pressures of the reactants and products to change immediately, so Q immediately after the temperature change is the same as before the temperature change; thus because Keq is now a different value, Q no longer equals Keq; the system has to move in whichever direction allows it to reach its new equilibrium state at the new temperature; that direction is determined by the enthalpy of the reaction; if a reaction is endothermic (delta H > 0), heat functions as a reactant; if a reaction is exothermic (delta H<0), heat functions as a product. EX: H204 (g) <-heat-> 2 NO2 (g); the equilibrium position can be shifted by changing the temperature; when heat is added and the temperature increases the reaction shifts to the right, and the flask turns reddish brown due to an increase in (NO2); when heat is removed and the temperature decreases the reaction shifts to the left, and the flask turns more transparent due to an increase in (N2O4) Key Concept: A (aq) + 2 B (g) <--> C (g) + heat Will shift to the right if...... A or B is added; C is removed; the pressure is increased or the volume is reduced; the temperature is reduced Will shift to the left if.......C is added, A or B is removed; the pressure is reduced or the volume is increased; the temperature is increased

percent composition

the percentage of the total formula weight of a compound attributable to a given element; to determine the percent composition of an element in a compound, the following formula is used: Percent Composition = mass of elements in formula / molar mass x 100%; one can calculate the percent composition of an element by using either the empirical or molecular formula; it is also possible to determine the molecular formula given both the percent compositions and molar mass of a compound; for example, what is the percent composition of chromium in K2Cr2O7? the molar mass is about 292 g/mol, Cr atomic weight is 52 g/mol, so percent composition = (2x 52)/294 x 100%, rounds to (2x50)/300 x 100% = 100/300 x100% = 33%

Transition state Theory

the point during a reaction in which old bonds are partially broken and new bonds are partially formed; has a higher energy than the reactants or products of the reaction and is also called the activated complex; the energy required to reach this transition state is the activation energy; once the activated complex is formed, it can either dissociate into the products or revert to reactants without any additional energy input; transition states are distinguished from reaction intermediates in that transition states are theoretical constructs that exist at the point of maximum energy, rather than distinct identities with finite lifetimes

triple point

the pressure and temperature at which the solid, liquid, and gas phases of a particular substance coexist in equilibrium

Osmotic pressure

the pressure that must be applied to a solution to prevent the passage of water through a semipermeable membrane down its concentration gradient; best thought of as a "sucking" pressure drawing water into solution; the osmotic pressure is the amount of pressure that must be applied to counteract this attraction of water molecules for the solution equation: where II is the osmotic pressure, i is the van't hoff factor, M is the molarity of the solution, R is the ideal gas constant and T is the temperature; water moves in t he direction of higher solute concentration, for instance pure water (no solute concentration) will traverse a semipermeable membrane to a solution containing solute particles (such as NaCl) and increase the level of the water as a result

partial pressure

the pressure that one component of a gaseous mixture would exert if it were alone in the container

Pauli exclusion principle

the principle stating that no two electrons within an atom may have an identical set of quantum numbers; the position and energy of an electron described by its quantum numbers are known as its energy state. The value of n limits the values of l, which in turn limit the values of ml. In other words for a given value of n, only particular values of l are permissible; given a value of l, only particular values of ml are permissible

thermodynamic product

the product of a reaction that is formed favorably at a higher temperature because thermal energy is available to form the transition state of the more stable product; has a larger overall difference in free energy between the products and reactants than the kinetic product; thermodynamic products are therefore associated with greater stability, and with more negative delta G than kinetic products

Kinetic Product

the product of a reaction that is formed favorably at a lower temperature because thermal energy is not available to form the transition state required to create a more stable thermodynamic product; has a smaller overall difference in free energy between the products and reactants than the thermodynamic product; because less free energy is needed for kinetic products these often form faster than the thermodynamic products and are sometimes called "fast" products; kinetic products are less stable than thermodynamic products

Rate Constant

the proportionality constant in the rate law of a reaction; specific to a particular reaction at a given temperature

molecular orbital

the region of electron density in chemical bonding that results from the overlap of two or more atomic orbitals; molecular orbits are obtained by combining the wave functions fo the atomic orbitals; qualitatively the overlap of two atomic orbitals describes this molecular orbitals; if the signs of the two atomic orbitals are the same a bonding orbital forms; if the signs are different an anti-bonding orbital forms; two different patterns of overlap are observed in the formation of molecular bonds; when orbits overlap head to head the resulting bond is a sigma bond; sigma bonds allow for free rotation about their axes because the electron density of the bonding orbital is a single linear accumulation between the atomic nuclei; when the orbits overlap in such a way that there are two parallel electron cloud densities a pi bond is formed; a pi bond do not allow for free rotation because the electron densities of the orbitals are parallel and cannot be twisted in such a way that allows continuous overlapping of the clouds of electron densities

Lanthanide series

the series of chemical elements atomic numbered 57-71 and falling between the S adn D blocks on the periodic table

Mechanism

the series of steps involved in a given reaction; knowing the accepted mechanism of a reaction may help explain the reactions rate, position of equilibrium, and thermodynamic characteristics; example generic reaction A2 + 2 B --> 2 AB however it actually occurs in two steps step 1: A2 + B --> A2B (slow) step 2: A2B + B --> 2 AB (fast)

Rate-determining step

the slowest step of a reaction mechanism; this step serves as a bottleneck on the progress of the reaction because it prevents the overall reaction from proceeding any faster than the slowest step

p subshell

the subshell corresponding to the angular momentum quantum number l = 1; contains three dumbbell-shaped orbitals oriented perpendicular to each other (px,py. and pz) and is found in the second and higher principal energy levels

Mass Number

the sum of protons and neutrons in an atom's nucleus; can also be called atomic mass number; in the convention A over Z by X where X is the element A is the mass number and Z is the atomic number

melting point

the temperature at which the solid and liquid phases of a substance coexist in equilibrium; identical to the freezing point

Magnetic quantum number (ml)

the third quantum number, defining the particular orbital of a subshell in which an electron resides; conveys information about the orientation of the orbital in space; each orbital can hold a maximum of two electrons; the possible values of ml are the integers between -l and +l, including 0. for example the s subshell with l=0 limits the ml values to 0 and because there is a singular value of ml there is only one orbital in the s subshell; the p subshell with l=1 limits the ml values to -1, 0, and 1, and because there are three values of ml, there are three orbitals in the p subshell, the d subshell have five orbits (-2 to +2) and the f subshell has seven orbitals (-3 to +3); the shape fo the orbits like the number of the orbits is dependent on the subshell in which they are found; the orbits in the s subshell are spherical, while the three orbits in the p subshell are each dumbbell shaped and align along the x, y, and z axes

Heat

the transfer of energy from one substance to another as a result of their differences in temperature; zeroth law of thermodynamics implies that objects are in thermal equilibrium only when their temperatures are equal; heat is a process function, not a state function: we can qualify how much thermal energy is transferred between two or more objects as a result of their difference in temperatures by measuring the heat transferred; processes that absorb heat are endothermic (delta Q (heat) > 0), processes that release heat are exothermic (delta Q < 0); the unit of heat is the unit of energy: joule (J) or calorie (cal) for which 1 cal - 4184 J; the process of measuring transferred heat is called calorimetry, the two basic types include constant-pressure calorimetry and constant-volume calorimetry; to measure heat transferred use the equation q = mc deltaT, mnemonic is q equals MCAT, where m is the mass, c is the specific heat of the substance, and delta T is the change in the temperature (in celsius or kelvin); it requires less heat to raise the temperature of a glass of water the same amount of a swimming pool, while these two items have the same specific heat, c, they have different heat capacities - the product mc (mass times specific heat); it is important to understand heat can transfer from a system to the surrounding on calorimetry questions, remember colder objects gains thermal energy and the hotter object loses it, use the equation q cold = - qhot

vaporization

the transformation of a liquid to gas

Joule (J)

the unit of energy; 1 J = 1 (Kg*m^2)/s^2

Periodic table

the visual display of all known chemical elements arranged in rows (periods); there are 7 periods representing the principal quantum numbers n=1 through n=7; and columns (groups) according to their atomic number and electron structure; groups contain elements that have the same electronic configuration in their valence shell and share similar chemical properties; the Roman numeral above each group represents the number of valence electrons elements in that group have in their natural state; roman numeral is combined with the letter A or B to separate the elements into two larger classes; the A elements are known as the representative elements and include groups IA through VIIIA, the elements in these groups have their valence electrons in the orbits of either s or p subshells; the B elements are known as non-representative elements and include both the transition elements which have valance electrons in the s and d subshells, and the lanthanide and actinide series, which have valence electrons in the s and f subshells; the first periodic table was designed by Dmitri Mendeleev and arranged by atomic weight; Henry Moseley later reorganized the period table my atomic number (the number of protons in an element); the periodic table creates a visual representation of the periodic law which states: the chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers

higher-order reactions

there are very few noteworthy reactions in which a single reaction step involves a termolecular process, in other words there are few processes with third-order rates; this is because it is far more rare for three particles to collide simultaneously with the correct orientation and sufficient energy to undergo a reaction

neutralization reaction

this is a specific type of double-displacement reaction in which a reaction between an acid and base in which a salt is formed (and sometimes water); for example hydrochloric acid and sodium hydroxide will react to form sodium chloride and water: HCl (aq) + NaOH (aq) ----> NaCl (aq) + H2O (l); reactions between acids and bases are not always visible, the addition of an indicator or use of indicator strips can determine when the reaction has occurred

percent composition by mass

this is used for aqueous solutions but also for metal alloys and other solid-in-solid solutions equation: mass of solute/mass of solution x 100%

Dilution

this occurs to a solution when solvent is added to a solution of higher concentration to produce a solution of lower concentration equation: MiVi = MfVf where M is the molarity, V is volume, and the subscripts i and f refer to the initial and final values of note the term parts-per can be used to indicate concentration of a dissolved substance in a solution (most commonly water); parts per million (ppm, 10^-6) is the most common usage; if a problem states there is one ppm of substance X in water, as there would be 1 millionth of a gram per gram of water, and the density of water is 1 g/mL

estimating scale values for pH (or PpOH)

when the original value of pH is a power of ten the operation is relatively easy: changing the sign on the exponent gives the corresponding p scale value directly; for example if [H=] = 0.001 or 10^-3 then the pH = 3 and pOH = 11; or if Kb= 1 x 10^-12, then pKb = 12 one can obtain a relatively close proximation of a p scale value using the following shortcut: if the non-logarithmic value is written in proper scientific notation, it will be in the form n x 10^-m, where n is the number between 1 and 10; taking the negative logarithm and simplifying, the p value will be: -log(n x 10^-m) = - log (n) - log (10^-m) which = m - log(n) because n is a number between 1 and 10 its logarithm will be a decimal between 0 and 1 (log 1 = 0 and log 10 = 1) the closer n is to 1, the closer log n will be to 0; the closer n is to 10 the closer log n will be to 1; a reasonable approximation one can say that: p value approximates (=) m - 0.n where 0.n represents sliding the decimal point of n one position to the left (dividing it by ten)