General Chemistry

amphoteric species

can act as either acid or base, depending on chemical environment. -in Bronsted-Lowry sense, amphoteric species can either gain or lose a proton. -the partially dissociated conjugate base of a polyprotic acid is usually amphoteric (i.e. HSO4- can either gain H+ to form H2SO4 or lose H+ to form SO4 2-). the hydroxides of certain metals like Al, Zn, Pb, and Cr are also amphoteric. species that can act as either oxidizing or reducing agents are also this, since by accepting or donating electron pairs they act as Lewis acids or bases.

osmotic pressure

because substances tend to flow or diffuse from higher to lower concentrations which increases entropy, water will diffuse from the compartment containing pure water to the compartment containing water-solute mixture. this net flow will cause the water level in the compartment containing the solution to rise above the level in the compartment containing pure water. -the pressure exerted by the water level in the solute-containing compartment due to gravity will eventually oppose the influx of water due to diffusion, and the water will stop flowing once this point is reached. this pressure is known as osmotic pressure. -molarity and osmotic pressure are directly proportional. the osmotic pressure depends on the amount of solute, not the identity.

neutrons

carry no charge and have a mass only slightly larger than that of protons so can still be considered to have a mass of appx 1 amu. Different isotopes of one element have different numbers of neutrons but same number of protons.

electron capture

certain unstable radionucleotides are capable of capturing an inner electron that combines with a proton to form a neutron. the atomic number is now one less than the original, but the mass number remains the same. electron capture is a rare process best thought of as inverse beta - decay, following exact same process as beta- decay but in reverse.

reaction rate

change of concentration of reactant or finished product with respect to time.

periodic law

chemical properties of the elements are dependent in a systematic way upon their atomic numbers.

real gases

deviations due to pressure: as pressure of gas increase, the particles are pushed closer. as condensation pressure for given temperature is approached, intermolecular forces become more significant until gas condenses into liquid state. at moderately high pressure, volume of a gas is less than would be predicted, due to intermolecular forces. at extremely high pressure, size of particles becomes relatively large compared to distance between them, and this causes the gas to take up larger volume than would be predicted. deviations due to temperature: as temperature of gas decreased, average velocity of the gas decreases, and intermolecular forces become significant. as condensation temperature is approached for given pressure, intermolecular forces cause gas to condense into liquid, and increasing intermolecular forces cause gas to have smaller volume than would be predicted. the closer the temperature of a gas to its boiling point, the less ideal.

alpha decay

emission of alpha particle which is a He nucleus that contains 2 protons and 2 neutrons. the alpha particle is massive compared to beta particle and doubly charged since it contains 2 protons and has +2 charge. alpha particles react with matter easily; hence they don't penetrate shielding such as lead sheets very far. -the emission of an alpha particle means that the daughter's atomic number will be 2 less than the parent's atomic number, and the daughter's mass number will be 4 less than parent's mass number.

catalysts

substances that increase reaction rate without themselves being consumed; they do this by lowering the activation energy. catalysts are important biological systems and in industrial chemistry; enzymes are biological catalysts. catalysts may increase the frequency of collision between the reactants, change the relative orientation of the reactants to make a higher percentage of collisions effective, donate electron density to the reactants, or reduce intramolecular bonding within reactant molecules. -the energy barrier for the catalyzed reaction is much lower than the energy barrier for the uncatalyzed reaction. note that the rates of both the forward and the reverse reactions are increased by the catalyst, since Ea of the forward and reverse reactions are lowered by the same amount. therefore, the presence of a catalyst causes the reaction to proceed more quickly toward equilibrium.

Dalton's Law of partial pressure

when two or more gases are found in one vessel without chemical interaction, each gas will behave independently of one another. the pressure exerted by each gas in the mixture will be equal to the pressure the gas would exert if it were the only one in the container. -partial pressure is related to mole fraction and can be determined using this equation.

law of mass action

while the forward and reverse reaction rates are equal at equilibrium, the molar concentrations of the reactants and products are usually not equal. this means that the forward and reverse rates are constant, kf and kr are usually unequal.

gram equivalents

for some substances it is useful to define a measure of reactive capacity. this expresses the fact that some molecules are more potent than others in performing certain reactions. for instance, 1 mol of HCl can donate 1 mol of H+ while 1 mol of H2SO4 contains 2 equivalents of H+ ions. -to determine number of equivalents a compound contains, a new measure of weight called gram-equivalent weight was developed such that equivalents= weight of compound/GEW GEW= molar mass/n n= either number of H ions an Arrhenius or Bronsted-Lowry acid could donate per molecule or the number of hydroxyl groups an Arrhenius base could donate in a reaction. -GEW is dependent on reaction conditions and is determined experimentally. for the PCAT, GEW can be estimated from the molecular structure. by using equivalents, it is possible to say one equivalent of acid will neutralize one equivalent of base, a statement not necessarily true when dealing with moles.

metalloids

found along the line between metals and nonmetals in the PT, and their properties vary. -their densities, boiling points, and melting points vary considerably. -the EN and IE lie between those of metals and nonmetals, so they possess characteristics of both classes. -i.e. silicon (Si) has metallic luster, but is brittle and isn't an effective conductor. -reactivity of metalloids is dependent on the element with which they are reacting. i.e. B behaves as a nonmetal when reacting with Na, but acts as a metal when reacting with F. -B, Si, Ge, As, Sb, Te.

higher order reactants

greater than 2.

reactant concentration

greater the concentrations of the reactants (the more particles per unit volume) the greater the number of effective collisions per unit time, so the reaction time will increase as reactant concentrations increase for all but zero-order reactions. for reactions occurring in the gaseous state, the partial pressures of the reactants can serve as measures of concentrations.

mixed order reactants

has a fractional order

second order reactions

has rate proportional to the product of the concentration of two reactants or to the square of the concentration of a single reactant. i.e k[A]^2 or k[A][B]. -units are M^-1s^-1.

vapor pressure lowering (Raoult's Law)

holds only when attraction between molecules of the different components of a mixture is equal to the attraction between molecules of any one component in pure state. when this condition doesn't hold, the relationship between mole fractions and vapor pressure will deviate from this law. solutions that obey this law are called ideal solutions.

exponential decay

let N be the number of radioactive nuclei that haven't yet decayed in a sample. it turns out that the rate at which the nuclei decay (deltaN/delta time) is proportional to the number that remain (N). -the solution of this equation tells us how the number of radioactive nuclei changes with time, which is known as exponential decay. -No is the number of radioactive nuclei changes at time t=0. -the decay constant is related to the half life by gamma = [0.693]/[t(1/2)]

electronegativity

measure of the attraction an atom has for electrons in a chemical bond. the greater the EN, the greater the attraction for bonding electrons. -most common EN scale is the Pauling EN scale, where the value range form 0.7 for most electropositive elements (cesium), to 4.0 for most electronegative elements (fluorine). -EN is related to Zeff. elements with low Zeff will have low EN because their nuclei don't attract electrons strongly, while elements with high Zeff will have high electronegativities because of the strong pull the nucleus has on electrons. -EN increases from left to right and decreases from top to bottom down a group.

reaction quotient Q

measure of the degree to which a reaction has gone to completion. -Q can be calculated based on the concentration of the reactants and products at any point during the course of a reaction, and the values vary considerably depending on when during the reaction the calculation is made. Qc is only a constant at equilibrium, when Qc is equal to Kc

quantum numbers

modern atomic theory states that any electron in an atom can be completely described by four quantum numbers: n, l, ml, and ms. -the pauli exclusion principle states that no two atoms can have the same four quantum numbers. -the position and energy of an electron described by its quantum numbers is known as its energy state. the value of n limits the values of l, which limits the values of ml. the values of three of the quantum numbers qualitatively give information about the orbitals: n about size, l about shape, ml about orientation of the orbital.

molality (m)

number of moles of solute / kg of solvent. -for dilute aq solutions at 25ºC, molality is appx equal to molarity because of density of water at this temperature is 1 kg/L but this is an approximation and true ONLY for dilute aq solutions.

molarity (M)

number of moles of solute / liter of solution. -molarity depends on total volume of the solution, not volume of solvent used to prepare solution.

fusion

occurs when small nuclei combine into larger nucleus. -the sun powers itself by fusing 4 H nuclei to make one He nucleus, by this method the sun produces 4x10^26 J every second. -these can only take place at very high temperatures, which is why they're called thermonuclear reactions.

applications of stoichiometry

once an equation is balanced, the ratio of reactants to moles of product is known, and that info can be used to solve many types of stoichiometry problems. -units should cancel to obtain the answer represent those asked for in the problem.

compound

purse substance composed of two or more elements in fixed proportion. compounds can be broken down chemically to produce their constituent elements or other compounds.

definition of rate

rate = decrease in conc of reactants/ time = increase in conc of products/time -minus signs are added before reactants since they are decreasing. -rate is expressed in units of moles per liter per second or molarity per second.

factors affecting reaction rate

reactant concentrations, temperature, medium, and catalysts

H bonding

specific and strong. -when H is bonded to either F, O, N; the H atoms carries little of electron density of covalent bond. this positively charged H atom interacts with partial negative charge located on EN atoms of nearby molecules. -usually have high boiling points. -important behavior of water, alcohols, amines, and carboxylic acids.

STP

standard temperature and pressure is 273 K and 1 atm. this is different than standard conditions which are used for measuring standard enthalpy, entropy, Gibbs free energy or voltage.

applications of bohr model

Bohr postulated that an electron can exist only in certain fixed-energy states. the energy of the electron is quantized. -the energy of the electron is related to its orbital radius: the smaller the radius, the lower the energy state of the electron. the smallest orbit can be n=1, which is the ground state and the lowest energy level. -the Bohr model is also used to explain the atomic emission spectrum and atomic absorption spectrum of H and is useful in interpretation of spectrum of other atoms.

hydrogen ion equilibria (pH and pOH)

-H ion concentration is measured by pH, and OH- ion concentration is measured by pOH. -dissociation of H2O is an equilibrium reaction and therefore described by constant Kw, the water dissociation constant. [H+][OH-] = 10^-14, so pH+pOH = 14. -in pure H2O, [H+] is equal to [OH-], because for every mole of H2O that dissociates, one mole of H+ and one mole of OH- are formed. -a pH below 7 means excess of H+ ions, and above 7 means excess of OH-.

pnictogens

-N and the elements below it are group VA. -wide properties and mixture of nonmetals (N and P), metalloids (As and Sb), and a metal (Bi). -often forms covalent bonds but most commonly forms three per atom. -N commonly holds positive charge in organic reaction, making several N-containing compounds good bases.

gamma decay

-emission of gamma particles, which are high-energy photons. -usually follows another type of nuclear decay, and is a way for the nucleus to shed excess energy (similar to how an electron in an excited state emits a photon to shed energy). gamma particles carry no charge and simply lower the energy of the emitting nucleus without changing mass number or atomic number.

nuclear binding and mass defect

-every nucleus has a smaller mass than combined mass of constituent protons and neutrons. this difference is mass defect. this is explained by E = mc^2. the mass defect is a result of matter being converted to energy. -this energy is called binding energy and holds nucleons together in nucleus. -binding energy per nucleon peaks at Fe, which indicates it's the most stable atom. in general, intermediate sized nuclei are more stable than large and small ones.

Boyle's Law

-for a given gaseous sample held at constant temperature, the product of pressure and volume is constant PV = k. -pressure and volume are inversely proportional. -individual values of pressure and volume can vary greatly for given sample of gas. as long as temperature remains constant and amount of gas doesn't change, the product of PV will equal the same constant.

noble gases

-group VIIIA -completely nonreactive. -high IE, no EN. -low BP and are gases at room temperature.

radioactive decay half life

-half life is time it takes for half of the sample to decay by any of the above processes. after n half-lives, (1/2)^n of the original sample will remain, whereas 1- (1/2)^n will have decayed.

change in temperature affecting Le Châtelier's Principle

-heat may be considered as a product in an exothermic reaction and as a reactant in an endothermic reaction. -example: A<-> B + heat; if this system were placed in an ice bath, the temperature would decrease driving the reaction to the right to replace the heat lost. if the system were placed in a hot water bath, the reaction would shift to the left due to the increased concentration of heat. -not only does a temperature change alter to position of the equilibrium, it also alters the numerical value of the equilibrium constant. in contrast, changes in the concentration of a species in the reaction, in the pressure, or in the volume alter the position of the reaction quotient without changing the numerical value of the equilibrium constant itself.

Le Châtelier's Principles

-if an external stress is applied to a system currently at equilibrium, the system will attempt to adjust itself to partially offset the stress. -this rule is used to determine the direction in which a reaction at equilibrium will proceed when subjected to a stress, such as change in concentration, pressure, temperature, and volume.

change in pressure or volume affecting Le Châtelier's Principle

-in a system at constant temperature, a change in pressure causes a change in volume and vice versa. since liquids and solids are practically incompressible, a change in the pressure or volume of a system involving only these phases has little or no effect on their equilibrium. reactions involving gases may be greatly affected by changes in pressure or volume, since gases are highly compressible, changing the volume of a container effectively changes the concentration of gases it contains. -pressure and volume are inversely related. an increase in the pressure of a system will shift the equilibrium so as to decrease the number of moles of gas present. this reduces the volume of the system and relieves the stress of the increased pressure.

strong acid and strong base titration

-in the early part of the curve, the acidic species predominates, so the addition of small amounts of base won't change either the OH or the pH. -in the last part of the titration curve, when an excess of base has been added, the addition of small amounts of base won't change OH significantly, and the pH remains relatively constant. the addition of base most alters the concentration of H+ and OH- near the equivalence point, and thus the pH changes most drastically in that region.

changes in concentration affecting Le Châtelier's principle

-increasing the concentration of a species will tend to shift the equilibrium in the direction that will reestablish the equilibrium concentration of the species that is added, and decreasing its concentration will shift the equilibrium in the opposite direction. -this effect is often used in industry to increase the yield of a useful product or drive a reaction to completion. if D were constantly removed from the above reaction, the net reaction would produce more D and more C. Likewise, using excess of the least expensive reactant helps to drive the reaction forward.

weak acid and strong base titration

-initial pH of weak acid solution is greater than the initial pH of strong acid solution. the pH changes most significant early on in the titration, and the equivalence point is in the basic range. although the original weak acid and strong base are completely neutralized at the equivalence point, the side product of the weak conjugate base from the original weak acid will be able to react with water in solution consuming H+, thereby lowering H+ and raising pH.

oxidation reduction reactions

-involves transfer of electrons from one species to another. -OIL RIG -must occur simultaneously, resulting in electron transfer known as redox reaction. -oxidizing agent causes another atom in redox reaction to undergo oxidation, and itself is reduced. -reducing agent causes the other atom to be reduced, and itself oxidized.

fission

-large heavy (mass number > 200) atom splits to form smaller, more stable nuclei (especially noble gases) and one or more neutrons. it's important to note that because the original large nucleus is more unstable than its products, there is a release of a large amount of energy. spontaneous fission rarely occurs. -by absorption of low-energy neutron, fission can be induced in certain nuclei. of special interest are those fission reactions that release more neutrons since those other neutrons will cause other atoms to undergo fission. this in turn releases more neutrons causing a chain reaction.

double displacement reactions

-metathesis reactions -elements from two different compounds displace each other to form two new compounds. -occurs when one of the products is removed from solution as precipitate or gas, or when two of the original species combine to form a weak electrolyte that remains undissociated in solution. -example: when solutions of CaCl2 and 2 AgNO3 are combined, insoluble 2 AgCl forms in solution of Ca(NO3)2.

buffers

-mixture of weak acid and its salt (which consists of its conjugate base and a cation) or a mixture of weak base and its salt (consists of conjugate acid and an anion). two examples of buffers are a solution of acetic acid and its salt sodium acetate, and a solution of ammonia and its salt ammonium chloride. buffer solutions have useful property of resisting changes in pH when small amounts of acid or base are added.

radioactive decay

-naturally occurring spontaneous decay of certain nuclei accompanied by emission of specific particles. it can be classified as certain type of fission. there are three types of decay problems: 1. integer arithmetic of particle and isotope species. 2. radioactive half-life problems 3. use of exponential decay curves and decay constants.

weak acids and bases

-partially dissolve in aqueous solution. weak monoprotic acid in aq solution will achieve the following equilibrium after dissociation. -the acid dissociation constant, Ka measures the degree to which an acid dissociates by showing ratio of concentrations of the products (the conjugate base and the H+ donated) to that of the reactant (the original acid). -the weaker the acid, the smaller the Ka. Ka doesn't contain an expression for pure liquid and pure solids since they aren't included in Keq equations. -a weak monovalent base undergoes dissociation to give B+ and OH-. the base dissociation constant Kb is a measure of the degree that a base dissociates. the weaker the base, the smaller the Kb. -conjugate acid is the acid formed when base gains proton. conjugate base is formed when acid loses proton. -the equilibrium constant Kw is equal to the product of Ka and Kb, 10^-14.

properties of equilibrium constant Keq

-pure solids and liquids don't appear in equilibrium constant. -Keq is characteristic of a given system at a given temperature. -if the value of Keq is much larger than 1, an equilibrium mixture of reactants and products will contain very little fo the reactants compared to the products. -if the value of Keq is much smaller than 1, an equilibrium mixture of reactants and products will contain very little of the products compared to the reactants. -if the value of Keq is close to 1, an equilibrium mixture of products and reactants will contain approximately equal amounts of the two.

freezing-point depression

-pure water freezes at 0ºC, but for every mole of solute particles dissolved in 1 L of water, the freezing point is lowered by 1.86ºC. this is because the solute particles interfere with crystal formation that occurs during freezing; the solute particles lower the temperature which the molecules can align themselves to a crystalline structure.

combination reactions

-two or more reactants form one product. -can also occur when two compounds react to form a new compound.

boiling-point elevation

a liquid boils when its vapor pressure equals the atmospheric pressure. if the vapor pressure of a solution is lower than that of the pure solvent, more energy and a higher temperature will be required before vapor pressure equals atmospheric pressure. the extent to which the boiling point of a solution is raised relative to the pure solvent is given by this formula.

solubility product constant

-a sligtly soluble ionic solid exists in equilibrium with its saturated solution. AgCl (s) <-> Ag+ (aq) + Cl- (aq). -the ion product (Qsp) of a compound in solution is defined as : Qsp = [A]^m[B]^n -the same expression for a saturated solution at equilibrium defines the solubility product constant Ksp. -Qsp is defined with respect to initial concentrations and doesn't represent an equilibrium or saturated solution while Ksp does. -each salt has its own Ksp at a given temperature. if at a given temperature a salt's Qsp is equal to its Ksp, the solution is saturated and the rate at which the salt dissolves equals the rate at which it precipitates out of solution. if a salt's Qsp exceeds its Ksp, the solution is supersaturated and unstable. if the supersaturated solution is disturbed by adding more salt, other solid particles, or jarring the solution by a sudden decrease in temperature, the solid state will precipitate until Qsp = Ksp. if the Qsp is less than Ksp, the solution is unsaturated and no precipitate will form.

Lewis definition

-acid as an electron pair acceptor, and base as an electron pair donor. -as every Arrhenius acid is a Bronsted-Lowry acid, every Bronsted-Lowry acid is a Lewis acid and likewise for bases. -i.e. BCl3 and AlCl3 can each accept an electron pair but don't donate protons.

Bronsted-Lowry Definition

-acid is species that donates protons, while a base is one that accepts protons. -NH3 (g) can be called a base in this definition, but can't be an Arrhenius base because it isn't in aqueous solution. -these always occur in pairs, called conjugate acid-base pairs. the two members of a conjugate piar are related by transfer of a proton.

polyvalence and normality

-acidity or basicity of an aqueous solution is determined by relative concentrations of acid and base equivalents. an acid equivalent is equal to one mole of H+ a base equivalent is equal to one mole of OH- ion. some acids and bases are polyvalent. -the acidity or basicity of a solution depends on concentration of acidic or basic equivalents that can be liberated. this is directly indicated by normality which is N = molarity x equivalents/mol -equivalent weight: taking molecular weight and dividing it by the number of equivalents/mole.

John Dalton's atomic theory

-all elements are composed of very small particles called atoms. all atoms of a given element are identical in size, mass, and chemical properties (with the exception that we now know about isotopes). the atoms of one element are different from atoms of another element. -all compounds are composed of atoms of more than one element. for any given compound, the ratio of the number of atoms of any two of the elements present is either an integer or a simple fraction. -a given chemical reaction involves only the separation, combination, or rearrangement of atoms; it does NOT result in the creation or destruction of atoms.

Arrhenius Definition

-an acid is a species that produces H+ ions in aqueous solution, and a base is one that produces OH-.

nuclear chemistry

-an amount of energy called binding energy is required to break up a given nucleus into its constituent protons and neutrons. that energy is converted to mass via Einstein's E = mc^2 equation, resulting in larger mass for constituent protons and neutrons that that of the original nucleus; this difference is called the mass defect.

strong acids and bases

-completely dissociate into their component ions in aqueous solution. -[H+] = normality of strong acid, [OH-] = normality of strong base.

multiple components phase diagram

-complicated by the requirement that the composition of the mixture, as well as the temperature and pressure must be specified. -consider solution of two liquids A and B, the vapor above the solution is a mixture of the vapors A and B. the pressures exerted by vapor A and B on the solutions are the vapor pressures that each exert above its individual liquid phase. Raoult's Law enables one to determine the relationship between vapor A and the concentration of liquid A in the solution.

relative strengths of acids and bases

-depends largely on ability to ionize. the strength of an acid, can be measured by fraction of molecules of that acid undergoing ionization. -when an acid or base is strong, its conjugate acid or base will be weak. for incredibly strong acids, the conjugate base is so weak it's basically inert.

applications of kinetic molecular theory of gases: Graham's law of diffusion and effusion

-diffusion: molecules of gases mix with one another by virtue of individual kinetic properties. diffusion occurs when gas molecules move through a mixture. diffusion accounts for fact that an open bottle of perfume can be smelled across a room. it predicts that heavier gas molecules diffuse more slowly than lighter ones because of their different average speeds. -under isothermal and isobaric conditions, the rates at which two gases diffuse are inversely proportional to the square root of their molecular masses. -effusion is the flow of gas particles under pressure from one compartment to another through a small opening. for two gases at same temperature, the rates of effusion are proportional to the average speeds. the rates of effusion in terms of molar mass and found that the relationship is the same as that of diffusion.

gases

-display similar laws and behavior regardless of their identity. the atoms of gases move rapidly and are far apart from each other, and there is only very weak intermolecular forces exist between gas particles; this results in characteristic properties: ability to expand to fill any volume, to take on shape of container, and to flow as fluids. gases are also easily compressible.

single component phase diagrams

-gas phase is found at high temperatures and low pressures; the solid phase is found at low temperature and high pressure; the liquid phase is found at high temperature and high pressure. -the three phases are demarcated by lines indicating the temperatures and pressures at which two phases are in equilibrium. -line A is freezing/melting. -line B is vaporization/condensation. -line C is sublimation/deposition. -intersection of three lines is triple point. at this temperature and pressure, unique for a given substance, all three phases are in equilibrium. -the point at B is known as the critical point, the temperature and pressure above which the liquid gas phases aren't possible and supercritical fluids exist instead.

assumptions of kinetic molecular theory

-gases are made up of particles whose volumes are negligible compared to container volume. -gas atoms are inert and exhibit no intermolecular forces. -gas particles are in continuous, random motion undergoing collisions with other particles and container walls. -collisions between any two gas particles are elastic, meaning there's no overall gain or loss of energy. -average kinetic energy of gas particles is proportional to the absolute temperature of the gas and is same for all gases at given temperature.

gas-liquid equilibrium

-temperature of a liquid = average kinetic energy, but the kinetic energy of molecules will vary. few molecules near surface may have enough energy to leave liquid and enter gas phase. this is evaporation. each time the liquid loses a high energy particle, the temperature of the remaining liquid decreases; thus evaporation is a cooling process. -if a cover is placed on a beaker of liquid, the escaping molecules are trapped above the solution. these molecules exert a countering pressure, which forces some of the gas back into liquid, called condensation. -atmospheric pressure acts on liquids as a solid lid. as evaporation and condensation proceed, an equilibrium is reached in which the rates of the two processes become equal. once equilibrium is reached, the pressure that the gas exerts over the liquid is called vapor pressure of the liquid. vapor pressure increases as temperature increases, since many molecules have sufficient kinetic energy to escape into the gas phase. the temperature at which the vapor pressure of the liquid equal the external pressure is called the boiling point.

solution equilibria

-the process of solvation tends towards equilibrium. -immediately after a solute has been introduced into a solvent most of the change taking place is dissociation because no dissolve solute is initially present. -according to Le Châtelier's principle, as solutes dissociate, the reverse reaction (precipitation of the solute) also begins to occur. eventually, an equilibrium is reached and the rate of solute dissociation is equal to the rate of precipitation. at equilibrium, the net concentration of the dissociated solute remains unchanged regardless of the amount of solute added.

Gibbs function for phase changes

-the thermodynamic criterion for each equilibria is that the change in Gibbs free energy must equal 0, meaning no net energy is required for or released from the forward or reverse reaction.

titration and buffers

-used to determine molarity of acid or base. -done by reacting known volume of a solution of unknown concentration with known volume of solution with known concentration. when number of acid equivalents equals number of base equivalents, the equivalence point is reached. -when titrating polyprotic acids or bases, there are several equivalence points. -while a strong acid/base titration will have an equivalence point of pH 7, it need not always occur at 7. -you can either plot the pH of the solution as a function of adding titrant on a graph or watching for color change by indicator. these are weak organic acids or bases that have different colors in their undissociated and dissociated states. these are in low concentration and don't significantly alter equivalence point. -the point at which the indicator changes color isn't the equivalence point but the end point, this isn't significant mathematically, but shows the reaction is done. -the volume difference between end and equivalence point is small and may be corrected or ignored if done correctly.

nuclear vs. chemical reactions

-when the nucleus of atom is unstable, it may spontaneously emit particles or electromagnetic radiation. -nuclei may also change composition when nuclear transmutation occurs. this process involves the bombardment of the nucleus by electrons, neutrons, or other nuclei. these are all specific types of nuclear reactions, which involve changes to nuclei instead of electrons.

stoichiometric coefficients

used to indicate number of moles of a given species involved in the reaction.

hydrogen

H does resemble alkali metals because it has a single s valence electron and forms H+ ion, which is hydrated in solution. it can also form the hydride ion H-, which is too reactive to exist in water. in this respect, hydrogen resembles halogens in that it only requires one additional electron to reach the next noble gas configuration.

salt formation

acids and bases may react with one another, forming a salt and often but not always water in what is termed a neutralization reaction. -the salt may precipitate out or remain ionized in solution depending on solubility and the amount produced. neutralization reactions generally go to completion. the reverse reaction, where the salt reacts with water to give back acid or base is called hydrolysis. -four combinations of strong and weak acids and bases are possible: 1. strong acid + strong base 2. strong acid + weak base 3. weak acid + strong base 4. weak acid + weak base -the products of a reaction between equal concentration of a strong acid and a strong base are salt and water. -product of a strong acid and weak base is salt but no water is formed since weak bases aren't usually hydroxides. the cation of the salt is the conjugate base, which won't be inert and will attribute to pH. instead of neutral salt and water, an inert anion and weak acid cation are formed creating a slightly acidic solution. the opposite occurs for weak acid and strong base. -the pH of a solution containing both weak acid and base depends on relative strengths of the reactants.

Charles's Law

at constant pressure, the quotient of the volume and temperature of gas is constant. the volume of a gas is directly proportional to its absolute temperature.

exceptions to octet rule

atoms found in or beyond third period can have more than 8 valence electrons, since some valence occupy d orbitals. these can be assigned more than four bonds in Lewis structures.

liquids

atoms or molecules are held closely together with little space between them. they have definite volumes and can't be easily compressed. molecules can still move around and are in state of relative disorder. liquid can still change shape to fit its container, and molecules are able to diffuse and evaporate. -liquids are able to mix with each other and with other phases to form solutions. the degree to which they can mix is called miscibility. oil and water are almost completely immiscible, the molecules tend to repel each other due to their polarity difference. these normally form separate layers when mixed, and oil on top because it's less dense. under extreme conditions, such as violent shaking, two immiscible liquids can form fairly homogenous mix called an emulsion. although they look like solutions, emulsions are actually mixtures of discrete particles too small to be seen distinctly.

solids

attractive forces between atoms or molecules are strong enough to hold them rigidly together. thus the particles' only motion is vibration about fixed positions, and the kinetic energy of solids is predominantly vibrational energy. as a result, solids have definite shapes and volumes. a solid can be crystalline or amorphous. these are specific 3D arrangements with repeating patterns of atoms, ions, or molecules. most solids are crystalline in structure. the two most common forms of crystals are ionic and metallic. -ionic crystals: aggregates of positively and negatively charged ions; there are no discrete molecules. the physical properties of ionic solids include high melting points, high boiling points, strong electrostatic interactions, and poor electrical conductivity in the solid phase. these properties are due to the compounds' strong electrostatic interactions, which also cause the ions to be relatively immobile. ionic structures are given by empirical formulas. -metallic crystals: metal atoms packed densely together. they have high melting and boiling points as a result of strong covalent bonds. pure metallic structures are usually described as layers of spheres of roughly similar radii. -crystals are defined by their unit cells, which can be simple cubic, body-centered cubic, and face-centered cubic.

beta decay

emission of beta particle, which could either be a beta electron or beta positron from the nucleus. -positron is similar to electron with minimal mass but has a positive charge. -electrons and positrons don't normally reside in nucleus but are emitted when a proton or neutron in nucleus decays. this is because protons and neutrons are composed of elementary particles called quarks, which can recombine to form different particles. -in beta - decay, a neutron decays into a proton (and an antineutrino). -in beta + decay, a proton decays into a neutron and a B+ particle (and a neutrino). -neither of these reactions is concerned with electrons in orbitals outside of nucleus, which is ignored during radioactive decay. -beta - decay means neutron is consumed and proton takes place. the parent's mass number is unchanged, and the parent's atomic number is increased by 1. -beta + decay a proton is consumed and a neutron takes its place, the mass number is unchanged and the atomic number is decreased by 1. -note that beta - decay is the only radioactive decay on the test where the atomic number increases. -since beta particles are slightly charged and about 1,836x lighter than protons, the beta radiation from radioactive decay is more penetrating than alpha radiation.

Henderson Hasselbalch equation

estimates pH of solution in buffer region where concentration of species and its conjugate are present in appx equal amounts. when conjugate base concentration = weak acid concentration, it's halfway to equivalence point and pH = pKa.

liquid-solid equilibrium

even though atoms of solids are confined to definite locations, each atom or molecules can undergo motions about some equilibrium point. these vibrations increase when heat is applied. if atoms or molecules in solid phase absorb enough energy this way, the solid's 3D structure breaks down, and the liquid phase begins. the transition from sold to liquid is called melting or fusion. the reverse process is solidification, crystallization, or freezing. the temperature at which these processes occur is called melting point or freezing point. pure crystals have distinct melting points, amorphous solids, like glass, tend to melt over a large range of temperatures due to less ordered molecular distribution.

collision theory of chemical kinetics

for a reaction to occur, molecules must collide with each other. this theory states that the rate of reaction is proportional to number of collisions/s between reacting molecules. -reaction rates almost always increase with increasing temperature. -not all collisions result in a reaction. effective collusion (one that leads to formation of products) only occurs if molecules collide with correct orientation and sufficient force to break existing bonds to form new ones. -minimum energy of collision necessary for reaction to take place is the activation energy. -only a fraction of colliding particles have enough kinetic energy to exceed activation energy. this means only a fraction of collisions are effective. the rate of the reaction is therefore expressed as: rate=fZ where Z is total collisions/s and f is fraction that are effective.

Avogadro's principle

for all gases at constant temperature and pressure, the volume of the gas will be directly proportional to the number of moles of gas present, therefore all the gases have the same number of moles in the same volume.

temperature

for nearly all reactions the reaction rate will increase with temperature increases. since 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. consequently, the proportion of molecules having energies greater than Ea (and thus capable of undergo reaction) increases with higher temperature.

chemical equilibrium

when there is no net change in the concentration of products and reactants during a reversible chemical reaction, equilibrium exists. this is not to say that a reaction is static; change continues to occur in both the forward and reverse directions. equilibrium can be thought of as a balance between the two reaction directions.

applications of kinetic molecular theory of gases: average molecular speeds

the average kinetic energy and therefore average velocity of a gas particle is proportional to the absolute temperature of a gas. KE = 3/2kT -because of the large number of rapidly and randomly moving gas particles, the speed of individual gas molecule is nearly impossible to define. instead it is the average speed of all the gas particles that can be related exactly to temperature. some particles will move at either higher or lower speeds. -Maxwell-Boltzmann distribution curve shows distribution of speeds of gas particles at a given temperature. the curve below shows a distribution curve of molecular speeds at two temperatures, where T2 > T1. notice that the bell-shaped curve flattens and shifts to the right as temperature increases, indicating that, at higher temperatures, more molecules are moving at high speeds.

factors affecting solubility

the quantity of a salt that can be dissolved is considerably reduced when it is dissolved in a solution that already contains one of its ions rather than in a pure solvent. this reduction in solubility, called the common ion effect, is another example of Le Châtelier's principle. the common ion effect won't change Ksp but it will change molar solubility (the concentration of the individual ions).

medium

the rate of a reaction may also be affected by the medium in which takes place. due to differences in intermolecular forces and other stabilizing factors, certain reactions proceed more rapidly in aqueous solution, whereas other reactions proceed more rapidly in benzene. the state of the medium (liquid, solid, or gas) can also have a significant effect.

coordinate covalent bonds

the shared electron pair comes from lone pair of one of the atoms in the molecule. once such a bond forms, it is indistinguishable from any other covalent bond, so identifying such a bond is used only in keeping track of valence electrons. these are typically found in Lewis acid-base compounds. a Lewis acid is a compound that can accept an electron pair

polyprotic acids and bases

titration curve looks difference between there are multiple equivalence points.

ideal gas law

under STP conditions, volume of one mole of ideal gas is 22.4 L. -R = 8.21 x 10^-2

gas-solid equilibrium

when a solid goes directly into gas phase, it is called sublimation. i.e. dry ice. the reverse transition is called deposition.

transition state theory

when molecules collide with sufficient energy, they form transition state in which the old molecules are weakened and new bonds are beginning to form. the transition state then dissociates into products and new bond are fully formed. -transition state, also called the activation complex has greater energy that the reactants and products. -once an activated complex is formed, it can either dissociate into products or revert to reactants without any additional energy input. transition states are distinguished from intermediates in that, existing as they do at energy maxima, transition states don't have finite lifetime. -potential energy diagram illustrates relationship between activation energy, heats of reaction, and potential energy of the system before and after the reaction. the most important factors are the relative energies of the products and reactants. the enthalpy change is the difference between potential energies of the products and reactants. a negative enthalpy change indicate exothermic reaction and a positive enthalpy change indicates endothermic reaction. -the activated complex exists at the top of the energy barrier. the difference in potential energies between the activated complex and the reactants is the activation energy of the forward reaction; the difference in potential energies between the activated complex and the products in the activation energy of the reverse reaction.

zero order reactions

- have a constant rate that does not depend on the concentration of reactant -have units of Ms^-1 - rate can only be affected by changing the temperature or adding a catalyst - with respect to administration of medication, this is one in which the amount of drug administered/eliminated per unit time remains constant. -concentration of drug A can be calculated by this: [A] = [A0]-(k0)(t) where A0 is initial [A], k0 is rate constant, t is time. -zero order half-life changes with time and is proportional to the initial drug concentration. it is inversely proportional to the zero-order rate constant and can be represented with half-life = (1/2)(A0/k0)

net ionic equations

-net equations can be written in ionic form. -when displacement occurs, there are usually spectator ions that don't take part in overall reaction but simply remain in solution throughout. the spectator ion in this equation is sulfate which doesn't undergo a transformation. a net ionic reaction can be written showing only the species that actually participate in the reaction. -net ionic equations are important for demonstrating the actual reaction that occurs during a displacement reaction.

chemistry history

-1911: Ernest Rutherford provided experimental evidence that an atom has a dense, positively charged nucleus that accounts for only a small portion of the volume of that atom. -1900: Max Planck developed the first quantum theory, proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles called quanta. the energy value of a quantum is given by the equation E=hf, where h is Planck's constant, 6.626 x 10^-34 Js, and f is the frequency of the radiation.

for redox reactions in basic instead of acidic solution

-an additional step is required because H+ won't be readily available and shouldn't appear as reactant. use same initial steps, then add enough OH- to both sides to completely combine with free H+ forming water. then, remove any species that appears on both sides and complete step 5 as usual to confirm everything is still balanced.

assigning oxidation numbers

-assigned to atoms to keep track of redistribution of electrons during chemical reaction. -from the oxidation numbers of reactants and products, it is possible to determine how many electrons are gained and lost by each atom. -oxidation number of number of charges an atom would have in a molecule if electrons were completely transferred in the direction indicated by the difference in EN. -element is said to be oxidized if its oxidation increases in a reaction, and an element is said to be reduced if the oxidation number decreases in a reaction.

nonpolar covalent bond

-atoms have same EN. the bonding electron pair is shared equally such that there is no separation of charge across the bond. they occur in diatomic molecuules.

balancing equations

-chemical equations must be balanced so that there are the same number of atoms of each element in the products as there are in the reactants.

decomposition reactions

-compound breaks down into two or more substances, usually as a result of heating or electrolysis. -when compounds that contain oxygen are heated, most will decompose to form O2. -electrolysis is a specific processes that causes decomposition of compound by passing electrical current through reactant.

electrolytes

-electrical conductivity of aq solutions is governed by presence and concentrations of ions in solution. -pure water doesn't conduct electrical current well since concentration of H+ and OH- ions are very small. -solutes that make conductive solutions are called electrolytes. -examples of strong electrolytes are ionic compounds like NaCl or KI, and molecular compounds with highly polar covalent bonds that dissociate into ions when dissolved, like HCl and water. -a weak electrolyte ionizes or hydrolyzes incompletely in aq solution, and only some of the solute is present in ionic form. examples include acetic acid and other weak acids and bases. -many compounds don't ionize at all in aq solution, retaining their molecular structure in solution which limits their solubility. these are called nonelectrolytes and include many nonpolar gases and organic compounds.

empirical and molecular formulas

-empirical: gives simplest whole number ratios of the elements in the compound. -molecular: gives exact number of atoms of each element in the compound and is a multiple of the empirical formula. -example: benzene. empirical: CH; molecular: C6H6. -some will have the same like H2O. -an ionic compound will only have empirical formula. -note that given a molecular formula, the empirical formula for that molecule can be calculated by simplifying the ration of the subscripts next to each other. -given only an empirical formula, molecular formula can't be determined.

electron configuration and orbital filling

-for an atom or ion, the pattern by which subshells are filled and the number of electrons within each principal level and subshell are designed by an electron configuration. -in configuration notation, the first number denotes the principal energy level, the letter designates the subshell, and the superscript gives the number of electrons within that subshell. -Aufbau principle: subshells are filled from lowest to highest energy and each subshell will fill completely before electrons begin to enter the next one. -(n+l) rule: used to rank subshells by increasing energy. this rule states that the lower the sum of the first and second quantum numbers, the lower the energy of the subshell. if two subshells posses the same value, the subshell will the lower n value has a lower energy and will fill first. -subshells with more than 1 orbital will fill according to Hund's rule: within a subshell, orbitals are filled such that there are a max number of half-filled orbitals with parallel spins. electrons prefer empty orbitals to half filled ones because a pairing energy must be overcome for two electrons carrying repulsive negative charges to exist in same orbital. -the presence of paired or unpaired electrons affects the chemical and magnetic properties of an atom or molecule. if the material has unpaired electrons, a magnetic field will align the spins of these electrons and weakly attract the atom to the field. these are said to be paramagnetic. materials that have no unpaired electrons and are slightly repelled by a magnetic field are said to be diamagnetic.

covalent bonds

-more similar EN differences. -the binding force between the two atoms results from the attraction that each electron of the shared pair has for the two positive nuclei. covalent compounds generally contain discrete molecular units with weak intermolecular forces. they have low melting points and don't conduct electricity. -can have single, double, triple bond. the number of shared pairs between atoms is called the bond order, a single bond has a bond order of one, and so on. -bond length: average distance between the two nuclei. as number of shared pairs increases, the two atoms are pulled closer together, leading to decrease in bond length. -bond energy: energy required to separate two bonded atoms. for a given pair of atoms, the strength of the bond increases as number of shared electron pairs increases.

cations and anions

-for elements that can form more than one positive ion (usually transition metals), the charge is indicated by a roman numeral in parentheses following the name of the element. -older but still commonly used way is to add -ous or -ic to the root of the latin name of the element to represent the ions with lesser or greater charge, respectively. -monatomic anions are named by dropping the ending and adding -ide. H- hydride, F- fluoride, O2- oxide -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. NO2- nitrite, NO3- nitrate, SO3 2- sulfite, SO4 2- sulfate. -when the series of oxyanions contains four, prefixes are also used. hypo- and per- are used to indicate less and more oxygen, respectively. ClO- hypochlorite, ClO2- chlorite, ClO3-, chlorate, ClO4- perchlorate. -polyatomic anions often gain one or more H+ 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.

rate law

-for nearly all forward, irreversible reactions, rate is proportional to the prospect of the concentration of reactants, each raised to some power. -k is rate constant, which is the constant of proportionality between chemical reaction rate and concentration of reactants. -multiplying the units of k by the concentration factors raised to appropriate powers gives the rate in units of conc/time. -x and y are orders of reaction; x is order respect to A and y is respect to B. the exponents may be integers, decimals and must be determined experimentally. -exponents of rate law aren't necessarily equal to coefficients in overall reaction equation. the exponents are equal to coefficients of rate determining step. -if one of the reactants or products of the rate-determining step is not included in overall reaction, then calculating the rate law can be complex. -overall order of a reaction or reaction order is defined as sum of exponents.

nonmetals

-generally brittle in the solid state and show little or no metallic luster. -high IE and EN and are poor conductors of heat and electricity. -show ability to gain electrons easily but otherwise display a wide range of chemical behaviors and reactivities. -upper right side of PT and are separated from the metals by a line cutting diagonally through the region of the PT containing elements with partially filled p orbitals.

alkali metals

-group IA. -possess most of the physical properties common to metals, yet their densities are lower. -have only one loosely bound valence electron in the outermost shell, giving them the largest atomic radii of all the elements in their respective periods. -metallic properties and high reactivity are due to low IE. -readily lose their valence to form univalent cations. -low EN and react readily with nonmetals, especially halogens.

alkaline earth metals

-group IIA. -possess many characteristic metal properties. -properties dependent on the ease with which they lose electrons. -have two valence in outermost shell, and smaller atomic radii. the valence aren't held tightly by nucleus, so they can be removed to form divalent cations. -have low EN and positive EA.

carbon group

-group IVA -wide range of characteristics and includes nonmetal (C), metalloid (Si and Ge), and metals (Sn and Pb). -they all have 2 electrons in their outermost p subshell, leading to a configuration that is distant from a noble gas. this is why C doesn't form ions (which would have to be +4 or -4 to reach noble gas configuration) but rather participates in electron sharing. most stable with four covalent bonds.

Chalcogens

-group VIA, which contains oxygen. -elements requiring two additional valence electrons to complete outermost shell. -fairly EN and form -2 anions. -can participate in covalent bonds, preferring to have two shared electron pairs and two non-bonded pairs.

halogens

-group VIIA -highly reactive nonmetals with one valence electron less than the closest noble gas. -commonly -1 anions. -halogens are otherwise variable in properties. -halogens range from gases (F2 and Cl2) to liquid (Br2) and solid (I2) at room temperature. -EN of halogens are really high, and they are reactive towards alkali metals and alkaline earths, which want to donate e- to halogens to form stable ionic crystals. -F has highest EN.

transition elements

-groups IB to VIIIB, all considered metals. -elements are very hard and have high melting points and boiling points. -from left to right across a period the five d orbitals become progressively more filled. the d electrons are held only loosely by the nucleus and are relatively mobile, contributing to the malleability and high electrical conductivity of these elements. -have low IE and may exist in variety of positively charged forms or oxidation states because transition elements are capable of losing various number of electrons from s and d orbitals of their valence shells. -because of ability to attain positive oxidation states, the transition metals form many different ionic and partially ionic compounds. the dissolve ions can form complex ions either with molecules of water (hydration complexes) or with nonmetals, forming highly colored solutions and compounds. this complexation can enhance the relatively low solubility of certain compounds. -the formation of complexes causes the d orbital to be 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 the higher d sublevel. the frequencies not absorbed- known as the subtraction frequencies give the complexes their characteristic colors.

first order reactions

-has a rate proportional to concentration of one reactant. -have units of s^-1 -classic example is process of radioactive decay. -in medicine, first order is one where % of drug administered/eliminated per unit time remains constant. the amount of drug administered/eliminated is proportional to the amount of drug remaining. -concentration of drug A can be calculated with equation: ln [A] = ln[A0] -ke(t) where ke is first order elimination rate constant.

solvation

-interaction between solute and solvent molecules is known as solvation or dissolution. -when water is solvent it is called hydration, and resulting solution is known as aqueous solution. -solvation is possible when attractive forces between solute and solvent are stronger than those between solute particles. example: Na and Cl interact better with water than they do each other due to polar ion-dipole interactions between Na+ and Cl- and the water molecules, which are stronger and more favorable than H bonds between the water molecules. -for nonionic solutes, solvation involves VDW interactions between the two. like dissolves like.

oxidation number rules

-number of free element (in elemental state) is 0. -for monoatomic ion is equal to charge of ion. -oxidation number of each Group IA element is +1. -oxidation number of each Group IIA element is +2. -oxidation number of each Group VIIA element is -1, except when combined with an element of higher EN. in HCl, the oxidation number of Cl is -1; in HOCl, the oxidation number of Cl is +1 because oxygen is more EN and has oxidation state of -2. -oxidation number of H is +1. it can be -1 when paired with Group IA and IIA. example: NaH. -oxidation number of O is -2. no the case for OF2, because F is more EN and has oxidation state of -1, so O has +2. in peroxides, the oxidation number is -1 because of structure of peroxide ion. -sum of oxidation numbers of all atoms present in neutral compound is 0. sum of atoms present in polyatomic ion is equal to charge of ion. for SO4 2-, sum must equal -2. -F has oxidation number of -1 in all compounds because it has highest EN. -metallic elements have only positive oxidation numbers; nonmetallic elements can have either.

polar covalent bonds

-occur between atoms with small differences in EN, generally range of 0.4 to 1.7 on Pauling scale. the bonding e- pair isn't shared equally but pulled towards atom with higher EN. the more EN atom always has a partial negative charge, while the less EN atom has a partial positive charge. -these are called polar molecules and have a dipole moment, which is measured in Debye units and defined as the product of the charge magnitude and the distance between two partial charges. u=qr

single displacement reactions

-occur when atom or ion of one compound is replaced by atom of another element. -example: Zn metal will displace Cu ion in CuSO4 solution to form ZnSO4. -single displacements are further classified as redox reactions.

dipole-dipole interaction

-polar molecules tend to orient themselves such that the positive region of one molecule is close to the negative region of another molecule. this arrangement is energetically favorable because an attractive dipole force is formed between the two molecules. -present in solid and liquid phases, not gas because these molecules are typically farther out. -polar species have higher boiling points than nonpolar species of comparable weights.

covalent bond notation

-shared valence electrons are called bonding electrons. the valence electrons not included in covalent bond are called nonbonding electrons, or lone pairs.

metal elements

-shiny solids at room temperature (except mercury) and generally have high melting points and densities. -are malleable. -have ductility -large atomic radius, low IE, low EN due to the fact that few electrons in valence shell of a metal atom can easily be removed. -because valence electrons can move easily, metals are good conductors of heat and electricity. -groups IA and IIA are the most reactive metals, and the transition elements are metals that have partially filled d orbitals.

factors affecting solubility

-solubility of a substance varies on temperature of solution, solvent, and in the case of gas-phase solute, the pressure. -solubility is also affected by the addition of other substances to the solution.

neutralization reactions

-specific type of double displacement that occurs when an acid reacts with base to produce a solution of salt and water. -example: Hcl and NaOH will react to form NaCl and H2O.

molecular mass

-the molecular mass is the sum total of atomic masses of the atoms in the molecule. -formula mass of an ionic compound is found by adding up the atomic masses of each constituent atom to the empirical formula of the substance.

dispersion forces

-weakest -at any particular point in time, the charges are located randomly throughout orbital. this permits unequal sharing of electrons, causing rapid polarization and counterpolarization of the electron cloud and forming short lived dipole moments. these dipoles interact with electron clouds of neighboring molecules, inducing formation of more dipoles. -they don't extend over long distances and are most important when they are close together the strength of these interactions within a given substance depends directly on how easily the electrons in the molecules can move. -large molecules in which the electrons are far from nucleus are easy to polarize and therefore possess greater dispersion forces. if not for dispersion forces, noble gases wouldn't liquefy at any temperature since no other intermolecular forces exist between noble gas atoms. the low temperature at which noble gases liquefy is to some extent indicative of the magnitude of dispersion forces between the atoms.

limiting reactant

-when reactions are mixed they are seldom added in the same exact stoichiometric proportions as shown in the balanced equation. in most reactants, one will be consumed first. -this is the limiting reagent because it limits amount of product formed. -reactant that remains after all of the limiting reactant is used up first is called the excess reactant.

electron affinity

EA is the energy change that occurs when an electron is added to a gaseous atom, and it represents the ease with which the atom can accept an electron. the higher the Zeff, the greater the EA will be. -positive electron affinity value represents the energy release when an electron is added to an atom. -group IIA, alkaline earth metals, have low EA values. these elements are relatively stable because their s subshell is filled. -group VIIA elements have high EA because the addition of an electron to an atom results in noble gas configuration. -achieving the stable octet involves a release of energy, and the strong attraction of the nucleus for the electron leads to a high change in energy. -group VIIIA have 0 EA because they already possess full shells and can't readily accept electrons. -elements of other groups have relatively low EA.

polarity of molecules

a molecule of two atoms bonded by a polar bond must have net dipole moment and therefore be polar. the two equal and opposite partial charges are localized at ends of each molecule on the two atoms. a molecule consisting of more than two atoms bound with with polar bonds may be either polar or nonpolar depending on direction of dipole moment.

nature of solutions

a solution consists of solute dispersed in a solvent. -solvent is component of solution whose phase remains the same after mixing. if two substances are already in same phase, the solvent is the component presented in greater quantity. -solute molecules move about freely in the solvent and can interact with other molecules or ions. consequently, chemical reactions occur easily in solution.

atomic emission spectra

at room temperature, the majority of atoms in a sample are in ground state. electrons can be excited to higher energy levels by heat or other energy to yield the excited state of the atom. the lifetime of the excited state is brief, and the electrons will return rapidly to the ground state while emitting energy in the form of photons. the electromagnetic energy of these photons may be determined using the following equation: E=hc/wavelength. -the different electrons in an atom will be excited to different energy levels. when these electrons return to their ground state, each will emit a photon with wavelength characteristics of the specific transition it undergoes. the spectrum is composed of light at specific frequencies and is known as line spectrum, where each line on the emission spectrum corresponds to specific electronic transition. each element possesses a unique atomic emission spectrum, which can be used as a fingerprint. -hydrogen has the simplest emission spectrum. the group of hydrogen lines corresponding to transitions from upper levels n>2 to n=2 is known as the Balmer series (four wavelengths in the visible region), while the group corresponding to transitions between upper levels n>1 to n=1 is known as the Lyman series (higher energy transitions in UV). -the energy associated with a change in the quantum number from an initial value ni to a final value nf is equal to the energy of Planck's emitted photon, and the energy of the emitted photon corresponds to the precise difference in E between the higher-energy initial, and the lower-energy final state.

quantum mechanical model of atoms

bohr's model did not take into consideration the repulsion between multiple electrons surrounding one nucleus. -the most important difference between the bohr model and modern quantum mechanical models is that bohr's assumption that electrons follow a circular orbit at a fixed distance is wrong. instead, electrons are described as being in a state of rapid motion within regions of space around the nucleus called orbitals. an orbital is a representation of the probability of finding an electron within a given region. -pinpointing the exact location AND momentum of an electron at any given point is impossible. this idea is best described by the Heisenberg uncertainty principle.

balancing redox reactions

by assigning oxidation numbers to reactants and products, one can determine how many moles of each species is required for conservation of mass and charge, which is necessary to balance equation. both the net charge and number of atoms must equal on both sides. -half reaction method where it is separated into two reactions, oxidation part and reduction part. 1. separate the two half reactions. 2. balance atoms of each half reactions. first balance all atoms except H and O. next, add H2O to balance O atoms, and H+ to balance H atoms. 3. balance the electrons of each half reaction. each half reaction must have same net charge on left and right sides, and only species that can be used to balance charges are electrons, each with -1 charge. additionally, the reduction half reaction must consume the same number of electrons as supplied by oxidation half. 4. combine half reactions. to find the final equation, cancel out the electrons and anything that appears on both sides. 5. confirm that mass and charge are balanced.

electrons

carry a charge equal in magnitude but opposite in sign to protons. it has a very small mass, appx. 1/1837 the mass of proton or neutron, which is negligible for purposes of PCAT. -the electrons in the electron shell farthest from the nucleus are known as valence electrons. the farther the valence electrons are from the nucleus, the weaker the attractive force of the positively charged nucleus and the more likely the valence electrons are to be influenced by other atoms. -the valence electrons and their activity determine the reactivity of the atom. -in a neutral atom, the protons and electrons are equal. Therefore, the atomic number indicates the number of electrons present in the atom. A positive or negative charge on atom is due to to loss or gain of electron. -note a positive charge is loss of electrons, and a negative charge is gain of electrons.

protons

carry a single positive charge and have a mass of appx. one unified atomic mass unit or amu, which is equal to one dalton (Da). -the atomic number Z of an element is equal to the number of protons found in an atom of that element. all atoms of a given element have the same atomic number. -these carry the same quantity of charge as an electron, but they have a mass that appx. 1,840x greater than an electron. -although the mass of the nucleus of an atom therefore comprises almost the entire weight of the atom, it occupies only 10^-13 of the volume of the atom.

molecule

combination of two or more atoms held together by covalent bonds and is the smallest unit of a compound that still displays the properties of the compound. molecules may contain two of the same element, or be comprised of two or more different atoms. -molecules are usually discussed in terms of molecular mass and molecules. -ionic compounds don't form true molecules, in the solid state they can be considered to be nearly infinite, 3D array of charged particles of which the compound is composed. because no molecule actually exists, molecular mass becomes meaningless and the term formula mass is used instead.

mole

defined as amount of substance that contains the same number of particles found in a 12.000 g of C-12. this quantity, Avogadro's number is equal to 6.022x10^23. -once mole of a compound has a mass in g equal tot he molecular mass of that compound in amu, and contains 6.022x10^23 molecules of that compound.

spin quantum number

denoted by ms. the spin of a particle is its intrinsic angular momentum and is characteristic of a particle. the two spin orientations are either +1/2 or -1/2. whenever two electrons are in the same orbital, they must have different spins. electrons in different orbitals with same spins are said to have parallel spins. electrons with opposite spins in the same orbital are referred to as paired spins.

azimuthal quantum number

designated by letter l. number tells us the shape of the orbitals and refers to the subshells or sublevels that occur within each principle energy level. for any given n, the value of l can be any integer in the range of 0 to n-1. -the four subshells corresponding to l= 0,1,2,3 are known as the sharp, principal, diffuse, and fundamental subshells or s,p,d,f. -the maximum number of electrons that can exist within a subshell is given by the equation 4l+2. -the greater the value of l, the greater the energy of the subshell. -energies of subshells from different principal energy levels can overlap. -for example, the 4s subshell will have a lower energy than the 3d subshell because its average distance from the nucleus is smaller.

magnetic quantum number

designated by ml. this describes the orientation of the orbital in space. an orbital is a specific region within a subshell that may contain no more than two electrons. this number specifies the particular orbital within a subshell where an electron is highly likely to be found at a given point in time. -the possible values of ml are all integers from l to -l, including 0. therefore the s subshell, l=0 has only one possible value of ml, and will contain only one orbital. the p subshell will contain 3 ml values, the d subshell will have 5 values, and the f subshell will have 7. -the shape and energy of each orbital are dependent upon the subshell in which the orbital is found.

intermolecular forces

dipole-ion > H bonding > dipole-dipole > LDF -forces not due to interaction of ions, or H bonding are known as Van Der Waals forces. -stronger intermolecular forces hold molecules together more tightly, so more energy (generally represented by higher temperature) is required to weaken these bonds to allow for phase changes.

valence electrons

electrons that are in its outer energy shell or that are available for chemical bonding. -for elements in groups IA and IIA, only the outermost s electrons are valence. -for elements in groups IIIA through VIIIA, the outermost s and p electrons in the highest energy shell are valence electrons. -for transition elements, the valence electrons are those in the outermost s subshell and in the d subshell of the next-to-outermost energy shell. -for the inner transition elements, the valence electrons are those in the s subshell of the outermost energy shell, the d subshell of the next-to-outermost energy shell and the f subshell of the energy of the energy shell two levels below the outermost shell. -groups IIIA-VIIA elements beyond period 2 might under some circumstances, accept electrons into their empty d subshells, which gives them more than 8 valence electrons.

ionization energy

energy required to remove an electron completely from gaseous atom or ion. this always requires an input of energy and is endothermic. the closer and more tightly bound an electron is the nucleus, the more difficult it will be to remove and the higher the IE. the first IE is the energy required to remove 1 valence e-, the second IE is the energy needed to remove a second valence e- from the univalent ion to make a divalent ion and so on. successive ionization energies grow increasingly larger. -IE increases from left to right across the PT as Zeff increases. IE decreases top to bottom as Zeff decreases. -Group IA have low ionization energies because the loss of an electron results in the formation of a stable noble-gas configuration.

mole fraction (X)

equal to number of moles of compound divided by the total number of moles of all species within the system. the sum of all the mole fractions in a system will always equal 1. Xb = moles of B / sum of moles of all components

isotopes

for a given element, multiple species of atoms with the same number of protons but different number of neutrons (different mass numbers, A) exist. since isotopes have the same number of protons and electrons, they generally exhibit the same chemical properties. -in nature, almost all elements exist as a collection of two or more isotopes, and these are usually present in the same proportions in any sample of a naturally occurring element. -standard atomic weight is a weighted average of all isotopes of an element found naturally on earth. -i.e. N has two stable isotopes, 14N and 15N, but 14N is much more common (99.6%) so the weighted average of the two is 14.007.

resonance

for some molecules, two or more non-identical Lewis structures can be drawn, these are resonance structures. the molecule doesn't exist as one or the other, but a hybrid. -often nonequivalent resonance structures may exist, the more stable the structure is, the more it contributes to the actual structure. -formal charges are often used to determine the stability of a particular resonance structure --a Lewis structure with small or no formal charge is preferred over Lewis structure with large formal charges. --a Lewis structure in which negative formal charges are placed on more EN atoms is more stable than those with formal charges on low EN atoms.

bonding

many molecules contain atoms bonded according to the octet rule, which stats than an atom tends to bond with other atoms until it has 8 valence electrons in outermost shell, thereby forming noble gas configuration. -exceptions: H & He (can only have 2) Li (2) B (6) Be (4), P and S can have more than 8 electrons by incorporating d orbitals. -ionic bonding: one or more electrons from an atom with smaller ionization energy are transferred to an atom with greater electron affinity, and the resulting ions are held together by electrostatic forces. -covalent bonding: electron pair is shared between two atoms. in many cases, the bond is partially covalent and partially ionic, these are polar covalent bonds.

percent concentration by mass

mass of solute / mass of solution (solute + solvent) x 100%

solubility

maximum amount of that substance that can be dissolved in a particular solvent at a particular temperature. when this maximum amount of solute has been added, the solution is saturated; if more solute is added it won't dissolve. -example: at 18ºC, a maximum of 83 g of glucose will dissolve in 100 mL of H2O. the solubility of glucose is 83 g/100 mL. if more glucose is added, it will remain in solid form -a solution in which the proportion of solute/solvent is small is said to be dilute, and one where the proportion is large is said to be concentrated. -when a dissolved solute comes out of solution and forms crystals, this is called crystallization. -some substances can form supersaturated solutions which contain more solute than found in saturated solution. these are formed by manipulating temperature or pressure. the addition of more solute will cause xs solute to separate, and a saturated solution with precipitate will form.

ion charges

metals, found in the left part of PT, generally form positive ions. nonmetals, found in the right part of the PT, generally form negative ions. -all elements in a given group tend to form monoatomic ions with the same charge. alkali metals form +1, alkaline earth metals forms +2, and halogens form -1.

normality (N)

number of GEW of solute / L of solution. -GEW = measure of reactive capacity of a molecule. -to calculate normality, we must known for what purpose the solution is being used because it is the concentration of the reactive species with which we are concerned. it is reaction dependent. -you can calculate normality by multiplying molarity of solution by number of equivalents per mol: N = molarity x equivalent/mol = [ mol/L ] x [ equivalents/mol ] = equivalents/L

formal charges

number of electrons officially assigned to an atom in Lewis structure doesn't always equal the number of valence electrons of the free atom. the difference between the two numbers is the formal charge. formal charge = V - (1/2Nbonding) - (Nnonbonding) or V - (# of sticks - # of dots) -the formal charge is equal to the sum of the formal charges of the individual atoms comprising it.

ionic radius

radius of cation or anion. the ionic radius will affect the physical and chemical properties of an ionic compound. in most situations, cations will be smaller than corresponding neutral atoms since possessing fewer electrons leads to less repulsion among the remaining electrons. most anions will be larger because there is greater repulsion.

aq solutions

soluble salts in water: --all salts of alkali metal ions. --all salts of ammonium ion --all salts with Cl-, Br-, I-, with exceptions of salts containing Ag+, Pb 2+, and Hg2 2+. --all salts of sulfate ion, with exception of those containing Ca 2+, Sr 2+, Ba 2+, and Pb 2+. insoluble salts: -all metal oxides, with exception of CaO-, SrO-, BaO, which hydrolyze to form solutions of the corresponding metal hydroxides. -all hydroxides, with exception of alkali metal hydroxides, Ca(OH)2, Sr(OH)2, and Ba(OH)2. -all salts with carbonates (CO3 2-), phosphates (PO4 3-), sulfides (S 2-), and sulfites (SO3 2-), with exception of those that contain alkali metals or ammonium.

dilution

solution is diluted when solvent is added to solution of higher concentration to produce solution of lower concentration. the concentration of solution after dilution can be determined using M1V1 = M2V2.

Law of Constant Composition

states that all samples of a given compound will contain same elements in identical mass ratios. -example: every sample of water will contain two atoms of H for one atom of O and therefore one gram of H for every 8 gram of O.

Lewis structures

steps: 1. count all valence electrons of atoms. number of valence electrons of molecule is the sum of valence electrons of all atoms present. 2. write skeletal structure of compound. the least EN atom is central, H and halogens occupy end positions. draw single bonds between the atoms. 3. complete the octets of all atoms bonded to central atom, using remaining valence electrons still to be assigned. 4. place any extra electrons on central atom. if it has less than octet, make double or triple bonds.

mass number

the atomic mass number (A) of an atom is equal to the total number of nucleons (protons and neutrons). on a larger scale, the molecular weight is the weight in g/mol of a given element. -mole is a unit used to count particles and is represented by Avagadro's number, 6.02x10^23 particles/mol, which is how many atoms of C are in 12.0 g of C-12. --Avogadro's number is the conversion factor between amu and g such that, if one atom of N has a mass of 14 u, then one mole of N has a mass of 14 g.

atomic radii

the atomic radius is equal to 1/2 the distance between the centers of two atoms of that element that are just barely touching one another. -the atomic radius decreases across a period from left to right, and increases down a given group. the atoms with the largest atomic radii are those in the bottom left corner. -something that affects the size of the electron cloud will change the radius of the atom, but altering the size of the nucleus will not directly affect the size of the atom. -from left to right, electrons are added one at a time to the outer energy shell. electrons within the same shell don't shield one another from the attractive pull of protons. since the number of protons increases from left to right, the Zeff increases as well. the greater the positive charge experienced by the valence e-, the closer those electrons are pulled TOWARDS the nucleus, decreasing the atomic radius. -from top to bottom, the number of electrons and filled shells increase. although the number of valence remain the same, these electrons will be found farther from the nucleus as they are in larger shells. Zeff becomes smaller with distance, so valence in higher energy shells will feel less pull from the nucleus. with more electrons comes additional negative charges that repulse each other, so the atomic radii increases.

periodic table

the elements are arranged in periods (rows) and groups (columns). there are seven periods, representing the principal quantum numbers n=1 to n=7, and each period is filled sequentially. -groups represent 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. -there are two sets of groups A & B. A elements are representative elements which have either s or p sublevels as their outermost orbitals. the B elements are the non-representative elements including the transition elements, which have partially filled d sublevels, and the lanthanide and actinide series, which have partially filled f sublevels. -i.e. an element in group VA has a valence electron configuration of s2p3 (2+3=5 valence e-).

percent composition

the mass percent of the element in a specific compound. % composition = [ mass of X in formula/formula weight ] x 100% -can be determined using either empirical formula or molecular formula. if % composition is known, empirical formula can be derived. it is possible to know the molecular formula if both the percent composition and the molecular mass of the compound are known.

periodic properties of the elements

the properties of the elements exhibit certain trends that can be explained in terms of the position of the element in the periodic table or in terms of the electron configuration of the element. all elements seek to gain or lose valence electrons so as to achieve stable fully filled forms possessed by the inert or noble gases in group VIIIA. -from left to right across a period, protons are added one at a time and the electrons of the outermost shell experience an increasing amount of nuclear attraction becoming closer and more tightly bound to the nucleus. the net positive charge from the nucleus as felt by the electron is called the effective nuclear charge (Zeff). -from top to bottom down a given column, the outermost electrons become less tightly bound to the nucleus. this is because the number of filled principal energy levels (which shield the outermost electrons from attraction by the nucleus) increases downward within each group. -together, this shows that the top right corner of the periodic table has the highest Zeff, and the bottom left corner has the lowest.

geometry and polarity of covalent molecules

the valence shell electron-pair repulsion (VSEPR) theory uses Lewis structures to predict the molecular geometry of covalently-bonded molecules. it states that the 3D arrangement of atoms surrounding a central atom is determined by repulsions between bonding and nonbonding electrons in the valence shell of central atom. these electron pairs arrange themselves as far apart as possible thereby minimizing repulsion. 1. draw Lewis structure 2. count total number of bonding and nonbonding electrons. 3. arrange e- pairs around central atom so they are as far apart as possible. 4. determine bond angle, accounting for additional repulsion due to nonbonding electrons, which pushes any bonding pairs slightly closer together.

yields

the yield of the reaction, which is the amount of product predicted or obtained when the reaction is carried out, can be predicted from the balanced equation. -theoretical yield: amount that can be predicted from balanced equation. -actual yield: amount of product isolated experimentally. -percent yield: expresses relationship between the actual yield and the theoretical yield percent yield = actual/theoretical x100%

principal quantum number (n)

this is the quantum number used in bohr's model that can take on any positive integer value and represents the shell where an electron is present in an atom. the maximum n that can be used to describe the electrons of an element at its ground state corresponds with that element's period in the PT. the larger the integer value of n, the higher the energy level and radius of the electron's orbit. the maximum number of electrons in an electron shell n is 2n^2. the difference in energy between shells decreases as the distance from the nucleus increases. the energy difference between the third and fourth shells is less than that between the second and third shells.

Bohr model

using Rutherford's findings, Bohr assumed that the H atom consisted of a central proton around which an electron traveled in circular orbit, and that the centripetal force acting on the electron as it revolved around the nucleus was the electrical force between the positively charged proton and the negatively charged electron. -Bohr's model used the quantum theory of Planck in conjunction with concepts from classical physics. an object, such as an electron, revolving in a circle may assume an infinite number of values for its radius and velocity. therefore, the angular momentum (L=mvr) and kinetic energy (KE=mv^2/2) can take on any value. Bohr used the quantum theory to place conditions on the value of the angular momentum, L=nh/2pi. h is Planck's constant, n is principle quantum number. since h, 2, and pi are all constants, the angular momentum changes only in discrete amounts with respect to n. -Bohr equated the allowed values of angular momentum to the energy of the electron. E=-Ry/n^2. Ry is the Rydberg energy constant equal to 2.18x10^-18 J/electron. Therefore, like angular momentum, the energy of the electron changes in discrete amounts with respect to n. a value of 0 energy was assigned to the state in which a proton and electron were separated completely, meaning that there was no attractive force between them. therefore, the electron in any of its quantized states in the atom would have a negative energy as a result of the attractive forces between the electron and the proton. this explains the negative sign in the equation.

experimental determination of rate law

values of k, x, and y in the rate law equation (rate= k[A]^x[B]^y must be determined experimentally for any reaction at a given temperature. the rate is usually measured as a function of the initial concentrations of the reactants, A and B.

atomic absorption spectra

when an electron is excited to a higher energy level, it must absorb energy. the energy absorbed as an electron jumps from orbital of low energy to one of higher energy is characteristic of that transition. this means the excitation of electron in a particular element results in energy absorptions at specific wavelengths. so every element possesses a characteristic absorption spectrum. the wavelengths of absorption spectrum correspond directly to the wavelengths of emission since the energy difference between levels remains unchanged. the absorption spectra can thus be used in the identification of elements present in gas phase sample.

ionic bonds

when two atoms with large differences in EN react, there is a complete transfer of electrons from one to the other. the atom that loses electrons becomes a cation, the other an anion. difference in EN must be greater than 1.7 on Pauling scale. in general groups IA and IIA bond ionically to elements of group VIIA. -have characteristic physical properties. they form crystal lattices consisting of arrays of positive and negative ions in which the attractive forces between ions of opposite charge are maximized, while the repulsive forces between ions of like charges are minimized. they therefore have high melting point and boiling points due to strong electrostatic forces between the ions. they also can conduct electricity in the liquid and aqueous states, though not in solid state.