Chemistry

Planck's constant

6.626 x 10^-34 J x s

As electrons go from a lower energy level to a higher energy level, they get AHED

Absorb light Higher potential Excited Distant (from the nucleus)

Group 2

Alkaline earth metals. Slightly higher effective nuclear charges and slightly smaller atomic radii than group 1. 2 electrons in their valence shell, both of which are easily removed to form a divalent cation.

Pauli Exclusion Principle

An atomic orbital may describe at most two electrons, each with opposite spin direction. No two electrons in a given atom can possess the same set of four quantum numbers.

Arrangement of the periodic table

Arranges the elements into periods (rows) and groups/families (columns), based on atomic number. Seven periods represent the principal quantum numbers, n=1 to n=7 for the s and b-block elements. Each period is filled sequentially, and each element in a given period has one more proton and one more electron than the element to its left (in their neutral states). Groups contain elements that have the same electronic configuration in their valence shell, and share similar chemical properties. The Roman numeral above each group represents the number of valence electrons that elements in that group have in their neutral state. The Roman numeral is combined with the letter A or B to separate elements with their valence electrons in S and P subshells (A) or elements with their valence electrons in s and d subshells (transition metals) or s and f subshells (lanthanide and actinide series)--these are nonrepresentative elements (B)

Atomic Emission Spectra

At room temperature, the majority of atoms in a sample are in the ground state. Heat or other energy forms can push atoms into excited states. Because the lifetime of an excited state is brief, the electrons will return rapidly to the ground state, resulting in the emission of discrete amount of energy in the form of photons. The electrons in an atom can be excited to different energy levels. When those electrons return to their ground states, each will emit a photon with a wavelength characteristic of the specific energy transition it undergoes. These energy transitions DO NOT form a continuum, but are quantized at certain values. As such, the spectrum is composed of light at specified frequencies--sometimes called the line spectrum--where each line on the emission spectrum corresponds to a specific electron transition. Because each element can have it electrons excited to a different set of distinct energy levels, each possesses a unique atomic emission spectrum, which can be used as a fingerprint for the element.

Isotopes

Atoms that share an atomic number but have different mass numbers.

Group 6

Chalgogens. Nonmetals and metalloids. Not as reactive as halogens, but crucial for normal biological functions. Each have six electrons in their valence electron shells, and generally have smaller atomic radii and larger ionic radii. Oxygen is the most important in the group--one of the primary constituents of water, carbs, and other bio molecules. sulfur is an important nutrient for microorganisms and has a role in protection from oxidative stress. The rest of the elements are primarily metallic and generally toxic to living organisms. At high concentrations, many of these elements, no matter how biologically useful, can be toxic or damaging.

Ionic Radius Trend

Decreases across a period and increases down a group. Metals lose electrons and become positive; nonmetals gain electrons and become negative. Metalloids can go in either direction, but tend to follow the trend based on which side of the staircase they fall on. Nonmetals close to the metalloid line possess a larger ionic radius than metals closer to group A. Can gain more electrons. Metals have a smaller ionic radius than the nonmetals near to the metalloid line. Less electrons to lose and so less reduction in radius during ionization. **Tellurium behaves as a nonmetal and Boron behaves as a metal. Under varying conditions, these metalloids can have opposite behavior.

Spin Quantum Number

Denoted by ms. Can have a +1/2 or -1/2 spin. Whenever two electrons are in the same orbital, they must have opposite spins (often referred to as being paired); electrons in different orbitals with the same ms have parallel spins.

Energy of an electron (Bohr Model)

E=-(RH)/n^2 RH=Rydberg unit of energy (2.18 x 10^-18 J/electron). like angular momentum, the energy of the electron changes in discrete amounts with respect to the quantum number. An electron in any of its quantized states in the atom will have an attractive force toward the proton (represented by the negative sign in the above equation). **While the magnitude of the fraction is getting smaller, the actual value it represents is getting larger (becoming less negative) E is directly proportional to n--the negative sign causes the values to approach zero from a more negative value as n increases (which means that energy is increasing).

Equation for determining electromagnetic energy of photons (Planck relation for wavelength)

E=energy h=Planck's constant (6.626 x 10^-34 J x s) c=speed of light in a vacuum (3.00 x 10^8 m/s) wavelength=wavelength of radiation **this is combined with two other equations (E=hf and c=f x lambda)

Planck Relation

E=hf E=energy, h=Planck's constant (6.626x10^-34 J x s)), f=frequency of radiation (sometimes designated by the Greek letter v).

Reading the periodic table for electron configurations (neutral atoms)

Electron configurations can be abbreviated by placing the noble gas that precedes the element of interest in brackets. E.g., the electron of any element in period four (starting with potassium), can be abbreviated by starting with [Ar]

Reading the periodic table for electron configurations (ions)

Electron configurations can be abbreviated by placing the noble gas that precedes the element of interest in brackets. E.g., the electron of any element in period four (starting with potassium), can be abbreviated by starting with [Ar]. For anions, add 1 to the last subscript. For cations, start with the highest neutral atom and remove electrons from the subshells with the highest n value, then electrons are removed from the subshell with the highest l value among these.

Valence electrons

Electrons on the outermost energy level of an atom. Easily removed and available for bonding. The active electrons of an atom; dominate the chemical behavior of the atom. For elements in Groups 1 and 2, only the highest s subshell electrons are valence electrons For elements in groups 13-18, the highest s and p subshell electrons are valence electrons. For transition elements, the valence electrons are those in the highest s and d subshells For the lanthanide and actinide series, the valance electrons are those in the highest s and f subshells all elements in period 3, starting with sodium and below, may accept electrons into their d subshell. this Allows them to violate the octet rule.

Active metals

Elements in Groups 1 and 2; have low ionization energies. Do not exist naturally in their neutral forms, always found in ionic compounds, minerals, or ores. The loss of 1 electron from Group 1 or 2 electrons from Group 2 result in the formation of a stable, filled valence shell (octet). Elements in Group 7, in contrast, typically do not give up their electrons (they just want one more to reach octet). Very low electron affinities--prefer to give up electrons to achieve the octet configuration.

Flouresence

Emissions from electrons dropping from an excited state to a ground state give rise to flourescence. What we see is the color of the emitted light.

More on ionization energies

First is always smaller than second which is always smaller than third, however, the degree to which an IE increases provides clues about the identity of an atom. if losing a certain number of electrons gives an element a noble gas-like configuration, then removing a subsequent electron will cost MUCH more energy. The values for second ionization energies are MUCH larger for group 1 but not that much larger for group 2.

Protons

Found in the nucleus of an atom. Each proton has an amount of charge equal to the fundamental unit of charge (e=1.6x10-19 C--denoted as +1). Have a mass of approximately one atomic mass unit (amu)

Metals

Found on left side and middle of the periodic table. Include active metals, transition metals, and lanthanide and actinide series. Lustrous (shiny) solids, except for mercury, which is a liquid under standard conditions. Generally have high melting points and densities (exceptions include lithium, with half the density of water). Ability to be deformed without breaking. At the atomic level, defined by a low effective nuclear charge, low electronegativity (high electropositivity), large atomic radius, small ionic radius, low ionization energy, and low electron affinity. Easily give up electrons!

Periodic trends

From left to right across a period, electrons and protons are added one at a time. As the positivity of the nucleus increases, the electrons surrounding nucleus, including those in the valence shell, experience a stronger electrostatic pull toward the center of the atom. As such, atomic radius DECREASES From left to right across a period, effective nuclear charge (attraction between valence shell electrons and nucleus) (Zeff) increases Moving top to bottom in a group, the principal quantum number increases by one each time. This means that the valence electrons are increasingly separated from the nucleus by a greater number of filled principal energy levels (inner shells). This increased separation results in a reduction in the electrostatic attraction between the valence electrons and the positively charged nucleus--the outermost electrons are held less tightly as the principal quantum number increases. Zeff is still constant. Elements can gain or loss electrons in order to achieve a stable octet formation representative of the noble gases (group 8A). Elements (especially ones that have biological roles, tend to be most stable with eight electrons in their valence shell).

valence electrons

Furthest out electrons that have the strongest attractions with the surrounding environment and weakest interactions with the nucleus. Much more likely to be involved in bonds with other atoms, because they experience less electrostatic pull from their own nucleus. Determine the reactivity of an atom.

Group 7

Halogens. Highly reactive nonmetals with seven valence electrons. Desperate to complete their octets by gaining one additional electron--highly electronegative! physical properties are variable--at standard conditions, they can be gaseous (F2, Cl2); liquid (Br2), and solid (I2). Especially reactive towards alkali and earth metals. Fluorine the most electronegative among all the elements. Naturally found as ions or diatomic molecules. **Frequently tested on the MCAT

Electronegativity

Increases across a period from left to right. and decreases from top to bottom. Measure of the attractive force that an atom will exert on an electron in a chemical bond. The greater the electronegativity of an atom, the more it attracts electrons within a bond. Related to ionization energies--the lower the ionization energy, the lower the electronegativity (and vice versa). Exception--the first three noble gases, which have high ionization energies, but negligible electronegativity, because they do not often form bonds.

Equation for Angular momentum of an electron (Bohr Model)

L=(nh)/2pi n=principal quantum number (any positive integer) h=Planck's constant (6.626x10^-34 J x s) Because the only variable is the principal quantum number, the angular momentum of an electron changes only in discrete amounts with respect to the principal quantum number

Transition metals as cofactors

Many transition metals act as cofactors for enzymes.

Types of elements

Metals, nonmetals, metalloids

Neutrons

Neutral--have no charge. Slightly larger mass than a proton. Number can be variable for a given element.

Group 8

Noble gases (also known as inert gases). Have minimal chemical reactivity due to their filled valence shells. have high ionization energies, little or no tendency to gain or lose electrons, and zero electronegativity. Extremely low boiling points and exist as gases at room temp. Used in neon lights.

Pauling electronegativity scale

Ranges from 0.7 for cesium, the least electronegative (most electropositive element) to 4.0 for Flourine, the most electronegative element. Cs=largest, least electronegative F=smallest, most electronegative

Azimuthal quantum number

Second quantum number, designated l. Refers to the shape and number of subshells within a given principal energy level (shell). Has important implications for chemical bonding and bond angles. The value of n limits the value of n: for any given value of n, the range of possible values for l is 0 to (n-1). Eg, within the first principal energy level, n=1 is the only possible values for l is 0; for n=2, the possible values for l are 0 and 1. **you can remember this because n tells you the number of subshells (n=1=1 subshell, n=2=2 subshells). The energies of the subshell increase with increasing value of l, however, the energies of subshells from different principal energy levels may overlap (e..g the 4s subshell will have a lower energy than the 3d subshell).

Metalloids

Separate the metals and nonmetals. Share characteristics with both and electronegativies and ionization energies lie between metals and nonmetals. Wide variance in densities, melting points, and boiling points. Reactivities are dependent on the elements with which they are reacting. Form a staircase on the periodic table.

Mass number

Sum of the protons in the atom's nucleus.

Equation for energy of electron transition

The equation says that the energy of the emitted photon corresponds to the difference in energy between the higher-energy initial state and the lower energy final state.

Electron affinity

The greed an element has for acquiring electrons--halogens are the most greedy. Need just one more to reach octet. Increases across a period from left to right and decreases from top to bottom. The energy dissipated by a gaseous species when it gains an electron. Exothermic process--Delta Hrxn is negative, however electron affinity is reported as a positive number. The stronger the electrostatic pull between the nucleus and valence shell electrons, the greater the energy release will be when the atom gains the electron. Group 8 elements have an electron affinity of zero because they already have a stable octet and cannot readily accept an electron.

Lyman series

The group of hydrogen emission lines corresponding to energy levels of n is greater than or equal to 2 to n=1. Larger energy transitions than the Balmer series and therefore has shorter photon wavelengths in the UV region of the electromagnetic spectrum.

Atomic radius

The size of a neutral element; equal to one half the distance between the centers of two atoms of an element that are briefly in contact with each other. Akin to diameter. Cannot be measured by examining a single atom, because electrons are constantly moving around, making it impossible to mark the outer boundary of the electron cloud. Across the periodic table, increases going down and to the left. Within each group, the largest atom will be at the bottom and within each period, the largest atom will be in Group 1A. Cesium has the largest atomic radius; helium has the smallest.

Atomic Absorption Spectra

When an electron is excited to a higher level, it must absorb exactly the right amount of energy to make the transition. This means that exciting the electrons of a particular element result in energy absorption at specific wavelengths. Every element possesses a characteristic absorption spectrum--the wavelengths of absorption correspond exactly to the wavelengths of emission, because the difference in energy between levels in energy remains unchanged. **identification of elements in the gas phase requires absorption spectra.

Relationship between complementary colors

When we perceive an object as a particular color, is it because that color is not absorbed, but rather reflected, by the object. If an object absorbs a given color of light, and reflects all others, our brain mixes these subtraction frequencies and we perceive the complement of the color that was absorbed.

Paragmagnetic

a magnetic field will cause parallel spins in unpaired electrons and therefore cause an attraction.

Atomic mass unit

a unit of mass that describes the mass of an atom or molecule. 1.66x10^-24

Malleability

ability of metal to be hammered into shapes

paired v. unpaired electrons

affects the chemical and magnetic properties of an atom or molecule. Materials composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field, and the material will be weakly attracted to the magnetic field. These materials are considered paramagnetic. paired electrons will be slightly repelled by a magnetic field and are diamagnetic. Given sufficiently strong magnetic fields beneath an object, any diamagnetic substance can levitate.

Group 1

alkali metals. Possess most of the classical properties of metals, except their densities are lower than those of other metals. Have only one loosely bound electron in their outermost shells. Very low Zeff values, which means they have the largest atomic radii. The low Zeff value also explains the low ionization energies, low electron affinities, and low electronegativities. Easily lose one electron to form univalent cations and react readily with nonmetals--especially halogens. Due to their high reactivity with water and air, most of these metals are stored in mineral oil.

Electrostatic force of attraction

because subatomic masses are so small, the electrostatic force of attraction between the unlike charges of the proton and electron is far greater than the gravitational force of attraction based on their masses

Alkali and alkaline earth metals

both metallic in nature because they easily lose electrons from the s subshell in their valence shells.

Application of atomic emission spectroscopy

can be used to analyze stars and planets--the light from a star can be resolved into its component wavelengths, which are then matched to the known line spectra of elements.

Periodic table

creates a visual representation of periodic law.

magnetic quantum number

designated as ml. specifies the particular orbital within a subshell where an electron is most likely to be found at a given moment in time. Each orbital can hold a max of two electrons. The possible value of ml are the integers between -l and +l, including 0. The s subshell, with l=0 limits the possible ml values to 0 and because there is a single value of ml, there is only one orbital in the s subshell. The p subshell, with l=1, limits the possible ml values to -1, 0, and +1, and because there are three values for ml, there are three orbitals in the p subshell. The d subshell has 5 orbitals (-2 to +2) and the f subshell has 7 orbitals (-3 to +3). The shape of the orbital, like the number of orbitals, is dependent on the subshell in which they are found. Orbitals in the s subshell are spherical, while the three orbitals in the p subshell are dumbbell shaped **For any value of l, there will be 2l+1 possible values for ml. For any n, this produces n^2 orbitals. For any value of n, there will be a maximum number of 2n^2 electrons (two per orbital). 2p=3 orbitals with 2 electrons each=6 elements s=2 elements d=ten elements f=fourteen elements

Aufbau principal

electrons fill from lower to higher energy subshells, and each subshell will fill completely before electrons begin to enter the next one. n+l rule: the lower the sum of the values of the first and second quantum numbers, n+l, the lower the energy of the subshell. So, if two subshells possess the same n+l value, the subshell with the lower n value has a lower energy and will fill with electrons first

Nonmetals

found predominately on the upper right side of the periodic table. Generally brittle in the solid state and show little or no metallic luster. Have high ionization energies, electron affinities, and electronegativities, as well as small atomic radii and large ionic radii. Usually poor conductors of heat and electricity. Do NOT want to give up electrons.

Transition metals

groups 3-12. Considered metals and have low electron affinities, low ionization energies, and low boiling points. Generally malleable and good conductors due to the loosely held electrons that progressively fill the d-orbitals in their valence shells. Many can have different oxidation states, because they are capable of losing different numbers of electrons from their s and d orbitals in their valence shells. E.g., copper can exist in either the +1 or +2 oxidation state. As such, they can form many different ionic compounds. Tend to associate in solution with either molecules of water or with nonmetals--their ability to form complexes contributes to the variable solubility (e.g. some aren't soluble in water but are soluble in ammonia). The formation of complexes causes d-orbitals to split into two energy levels, which enables many complexes to absorb certain frequencies of light--those containing the precise amount of energy required to raise electrons from lower to higher energy d-orbitals. The frequencies not absorbed (known as subtraction frequencies), give the complexes their characteristic colors.

Transition metals (group B elements)

have two or more oxidation states (charges when forming bonds with other atoms); because the valence electrons of all metals are only loosely held to their atoms, they are free to move, which makes good metal conductors of heat and electricity. Their valence electrons are found in the s and d subshells and the lanthanide and actinide series elements are in the s and f subshells. Some transition metals, including copper, nickel, silver, gold, palladium, and platinum--are relatively nonreactive, which makes the ideal for the production of coins and jewelry.

Excited state

in Bohr's model, when the electron is promoted to an orbit with a larger radius (higher energy). In general, an atom is in an excited state when at least one electron has moved to a subshell of higher than normal energy

Hund's rule

in subshells with more than 1 orbital, such as the 2p subshell with 3 orbitals, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins. Electrons want their own seat rather than being forced to double up! Half filled and fully filled orbitals have lower stabilities (higher energies) than other states. TWO exceptions often tested on MCAT--chromium and copper. Copper has a half shell configuration of [Ar]4s^1 3d^10 rather than [Ar]4s^2 3d^9 Chromium has a half shell configuration of [Ar]4s^1 3d^5, not [Ar] 4s^2 3d^4

Heisenberg uncertainty principle

it is impossible to know exactly both the velocity and the position of a particle at the same time. If we want to assess the position, the electron has to stop (removing momentum); if we want to assess the momentum, the electron has to be moving (changing its position).

Periodic trends in sum

left to right: increasing electronegativity increasing ionization energy decreasing atomic radius increasing electron affinity Bottom to top increasing electronegativity increasing ionization energy decreasing atomic radius increasing electron affinity

quantum numbers

modern atomic theory postulates that any electron in an atom can be completely described by four quantum numbers: n, l, ml, and ms. The quantum numbers become more specific from n to ms One lives in a particular state (n), in a particular city (l), on a particular street (ml), at a particular house (ms)

Electrons

move through the space surrounding the nucleus and are associated with varying levels of energy. Each electron has a charge equal and opposite to a proton (-1 e). The mass of an electron is 1/2000 of a proton. Move around the nucleus at varying distances, which correspond to varying levels of electrical potential energy. Electrons closer to the nucleus are at lower energy levels and those further out in higher electron shells have higher energy.

Anion

negatively charged atom

mole

number of things--atoms, ions, molecules--equal to Avogadro's Number (6.02=10^23). E.g., the atomic weight of carbon is 12.0 amu, which means that the avg. carbon atom has a mass of 12.0 amu, and 6.02x10^23 carbon atoms have a combined mass of 12.0 grams.

quantized energy

only certain energy levels are allowed--stairs, not a ramp.

Diamagnetic

paired electrons, repelled by magnetic field

cation

positively charged atom

Spectroscopic notation

shorthand representation of the principal (n) and azimuthal (l) numbers Principal number remains a number Azimuthal number is designated by a letter l=0 subshell is called s; Group 1-2 plus helium, 1s^1-7s^2 l=1 subshell is called p; Group 3-8 minus helium 2p^1-7p^6 l-2 subshell is called d; middle gap 3d^1- 6d^10 l-3 subshell is called f; bottom two rows 4f^1-5f^14 s, p, d, f example; electron in shell n=4 and subshell l=2 is said to be in the 4d subshell

ground state of an atom

state of lowest energy, in which all electrons are in the lowest possible orbitals. On the MCAT, atoms of any element will generally exist in the ground state, unless subjected to extremely high temperatures or radiation

Bohr's model of a hydrogen atom

structural model in which an electron moves around the nucleus only in circular orbits, each with a specific allowed radius; the orbiting electron does not normally emit electromagnetic radiation, but does so when changing from one orbit to another. An electron can jump from one orbit to another with the transfer of the amount of energy exactly equal to the difference between one orbit and another. Orbits have increasing radii. The orbit with the smallest, lowest-energy radius is defined as the ground state (n=1). Explained the atomic emission spectrum of hydrogen--the simplest emission spectrum among all elements.

MCAT equations (pro tip)

tends to ask how changes in one variable may affect another variable. Focusing on ratios and relationships is usually enough to give the right answer

Ductility

the ability of a metal to be drawn, pulled, or extruded through a small opening to produce a wire

periodic law

the chemical and physical properties of the elements are independent, in a periodic way, upon their atomic numbers.

Ionization energy (ionization potential)

the energy required to remove an electron from a gaseous species; always requires heat. Endothermic process. The greater Zeff, the closer the valence electrons are to the nucleus the harder it is to remove an electron, the greater the ionization energy. Ionization energy increases from left to right across a period, and from bottom to top in a group. The removal of a second or third electron requires increasing amounts of energy, because the removal of more than one electron means that the electrons are being removed from an increasingly cationic species. First ionization energy=energy required to remove the first electron; second ionization energy=energy required to move second electron, etc. Group 8 are the least likely to give up electrons--they already have a stable configuration--elements with the highest ionization energies.

Balmer series

the group of hydrogen emission lines corresponding to transition levels from n is greater than or equal to 3 to n=2. Includes four wavelengths in the visible region.

Pachen series

the group of hydrogen emissions that corresponds to transition levels from n is greater than or equal to 4 to n=3.

Principal quantum number

the larger the value of n, the higher the energy and radius of the electron's shell. Within each shell, there is a capacity to hold a certain number of electrons, determined by: 2n^2 The difference in energy between the two shells decreases as the distance from the nucleus increases. E.g., the energy difference between n=3 and n=4 shells, is less than the energy difference between n=1 and n=2 shells

Effective nuclear charge (Zeff)

the net positive charge experienced by valence electrons. Relies on the principles of electrostatic forces. q1 and q2 represent the net charge of the nucleus and valence electron shell respectively. The large each charge gets (going to the right in the periodic table), the higher the value of Zeff

Atomic number

the number of protons in the nucleus of an atom. Acts a unique identifier for each element, because elements are defined by the number of protons they contain.

Electron configurations

the pattern by which subshells are filled as well as the number of electrons within each principal energy level and subshell. Spectroscopic notation us used, where the first number denotes the principal energy level, the letter designates the subshell, and the subscript gives the number of electrons in that subshell. E.g., 2p^4 indicates there are four electrons in the second (p) subshell of the second principal energy level. This also implies that the energy levels below 2p (1s and 2s) have already been filled.

Energy state

the position and energy of an electron described by its quantum numbers. The value of n limits the value of l, which in turn limits the value of ml. The values of the quantum numbers qualitatively give information about the size, shape, and orientation of the orbitals.

Probability density

the shapes of orbitals are defined in terms of probability density, the likelihood that an electron will be found in a particular region of space.

Atomic weight

weighted average of an element's different isotopes