Chemistry MCAT Review
Chapter 1.2
Atomic number (Z) = number of protons
Mass number (A) = number of protons + number of neutrons
Mass number is equivalent to atomic mass
Avogadro’s number = 6.02 × 10²³
Atomic weight = a weighted average of naturally occurring isotopes of that element
Chapter 1.3
Energy of the quantum (Planck relation): E = uf
KE = ½ mv² ← mentioned Bohr Model
If one could transfer an amount of energy exactly equal to the difference b/w one orbit and another, this could result in the electron “jumping” from one orbit to a higher-energy one
Ground state: The atom is at the state of lowest energy
Excited state: when at least one electron has moved to a higher subshell than normal energy
tend to happen when exposed to extremely high temperatures or irradiation
As electrons go from a lower energy level to a higher energy level, they get AHED:
Absorb light
higher potential
excited state
distant (from the nucleus)
E = hc/wavelength
c = speed of light
Energy is inversely proportional to wavelength
delta E = absorption or emission between any 2 energy levels; energy of the photon of light absorbed/emitted
we see the color of light that’s not absorbed by the compound
Chapter 1.4
orbitals: regions of space around the nucleus that electrons move in
4 Quantum numbers: n, l, mi, ms
Pauli exclusion principle: no 2 electrons in a given atom possess the same set of 4 quantum numbers
max number of electrons within a shell = 2n²
a larger integer number/value for the principal quantum number indicates a larger radius and higher energy
principal quantum number: n
azimuthal (angular momentum) quantum number: l; second quantum number; refers to the shape and number of subshells within a given principal energy level (shell)
spectroscopic notation: refers to the shorthand representation of the principal and azimuthal quantum numbers
l = 0 → s; l = 2 → p; l = 3 → f
magnetic quantum number (mi): third quantum number
spin quantum number (ms): 4th quantum number
paired electrons: whenever 2 electrons are in the same orbital, they must have opposite spins
parallel spins: electrons in different orbitals with the same ms values
Hund’s rule: orbitals are filled such that they are max # of half-filled orbitals with spins
paramagnetic: when materials composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field, and they will be weakly attracted to the magnetic field
means that a magnetic field will cause an attraction
diamagnetic: when materials consisting of electrons that will be slightly repelled by a magnetic field
Chapter 2.1
periods → rows
groups/families → columns
valence electrons → farthest from the nucleus and have the greatest amount of potential energy
A elements: representative elements → have valence electrons in the orbitals of either s or p subshells
B elements: nonrepresentative elements → include transitional elements (valence electrons in s and d subshells) and the lanthanide and actinide series (valence electrons in s and d subshells) and the lanthanide/actinide series (valence electrons in s and f subshells)
many transitional metals (group B elements) have 2 or more oxidation states
Chapter 2.2
metals: left-middle of Periodic Table; malleability
nonmetals: the right of the periodic table; brittle, poor conductivity
metalloids: stair group; semimetals, share both characteristics
Chapter 2.3
effective nuclear charge (Zeff): a measure of the net positive charge experienced by the outermost electron
highest at the upper right of the PT
atomic radius
highest at the lower left of the PT
atomic radius is the opposite of most trends
ionization energy: the energy required to remove an electron from a gaseous species
highest at the upper right of the PT
first ionization energy” the energy necessary to remove the first electron
second ionization energy: the energy necessary to remove the 2nd electron from teh univalent cation (X+) to form the divalent cation (X²^+)
2nd ionization energy is always greater than 1st ionization energy
endothermic process: when removing an electron from an atom, it requires an input of heat
the greater the atom’s Zeff or the closer the valence electrons are to the nucleus, the more tightly bound they are
active metals: low ionization energies/elements (elements in Group 1 and 2)
Noble or inert gas → least likely to give up electrons
exothermic process: expels process in the form of heat
electron affinity: refers to the energy dissipated by a gaseous species when it gains an electron
usually reports as a positive number
highest in the upper right of the PT
electronegativity: a measure of the attractive force that an atom will exert on an electron in a chemical bond
the greater electronegativity, the more it attracts electrons within a bond
Chapter 2.4
alkali group (group 1): possess most physical properties, easily lose one electron, react readily with nonmetals/halogens
chalcogens: an eclectic group of nonmetals and metalloids, 6 electrons
halogens: highly reactive nonmetals with 7 valence electrons
halides → diatomic molecules
Noble gases: inert gases → have minimal chemical reactivity due to their filled valence shells, high ionization, little/no tendencies to gain/lose electrons, no measurable electronegativity
Transitional metals: groups 3-12; low electron affinities, low ionization energies, low electronegativity
act as cofactors for enzymes
if an object absorbs a color of light and reflects all others, our brain mixes these subtraction frequencies and we perceive the complementary color of the frequency that was absorbed
we see reflected colors
Chapter 3.1
Octet Rule: states that an atom tends to bond with other atoms so that it has eight electrons
Incomplete octet: any element in period 3 or greater than 8 electrons in their valence shell
Expanded octet: Any element in period 3 or greater can hold more than 8 electrons
odd numbers of electrons: any molecule with an odd number of valence electrons can’t distribute those electrons give 8 to each atom
Ionic bonding: when one or more electrons from an atom with a low ionization are transferred to an atom with a high electron affinity (metal → nonmetal)
covalent bonding: when an electron pair is shared between 2 atoms (nonmetals) that have relatively similar values of electronegativity
the degree to which the pair of electrons is shared equally or unequally between the 2 atoms determines the degree of polarity in the covalent bond
electron pair is shared equally = nonpolar
electron pair is shared unequally = polar
coordinate covalent: a bond when both of the shared electrons are contributed by only one of the 2 atoms
Chapter 3.2
Ionic bonds form between atoms that have significantly different electron activities
electrons are not shared
atom that loses electrons becomes a cation (+); atom that gains electrons becomes an anion (-)
Characteristic physical properties of ionic compounds:
have very high melting and boiling points
compounds dissolve readily in water
polar solvents are good conductors of electricity
Chapter 3.3
when the difference in electronegativity between is great, the “stronger” atom wins all of the electrons and becomes the anion
2 atoms sharing one, two, or three pairs of electrons are said to be joined by a single, double, or triple covalent bond
bond order: the # of shared electron pairs between 2 atoms
bond length: the average distance between the 2 nuclei of atoms in a bond
as the number shared of electron pairs increases, the 2 atoms are pulled closer together, resulting in a decrease in bond length
inverse relationship → bond length and strength
single bond is long; triple bond is short
bond energy: the energy required to break a bond by separating its components into their isolated, gaseous atomic states
the greater the number of pairs of electrons shared between the atomic nuclei, the more energy is required to break the bonds holding the atoms together
polarity: when 2 atoms have a relative difference in electronegativities
nonpolar covalent bond: when atoms that have identical or nearly identical electronegativities share electron pairs, they do so with an equal distribution of electrons
diatomic molecules: H2, N2, O2, F2, Cl2, Br2, l2 (form the number seven and there are seven elements)
polar covalent bond: when atoms that differ moderately in their electronegativities will share electrons unevenly
partial negative charge: the result when the more electronegativity element acquired a smaller portion of the electron density
partial positive charge: the result when the less electronegative element acquires a smaller portion of the electron density
polar molecule: a molecule that has such a separation of positive and negative charges
In a coordinate covalent bond, both of the shared electrons originated on the same atom
a lone pair of one atom attacked another atom with an unhybridized p-orbital to form a bond
Lewis acid → accept a lone pair of electrons
Lewis base → donate a pair of electrons - form covalent bond)
Bonding electrons → only in electrons involved in a covalent bond that are in the valence shell
nonbonding electrons → only electrons in the valence shell that are not involved in covalent bonds
Lewis structure (Lewis dot diagram): the chemical symbol of an element surrounded by dots, each representing one of the s or p valence electrons of the atom
least electronegativity atom is the central atom
total valence electrons is the amount of bonds/electrons in Lewis structure
central atom must try to have an octet
Chapter 4.1
molecule: a combination of 2 or more atoms held together by covalent bonds
ionic compounds form from combinations of elements with large electronegativity differences
formaula weight of an ionic compound → add up the atomic weights of ions according to its empirical formula
mole: a quantity of any substance equal to the number of particles that are found in 12 grams of carbon-12
molar mass: the mass of one mole of a compound (g/mol)
n → number of particles of interest produced/consumed per molecule of the compound in the reaction
normality: a measure of concentration (equivalents/L)
Molarity = Normality/n
n → number of protons, hydroxide ions, electrons, or ions produced or consumed by solute
Chapter 4.2
Law of constant composition: states that any pure sample of a given compound will contain the same elements in an identical mass ratio
empirical formula: simplest whole number ratio of the elements in the compound
2 ways to the formula of a compound: empirical and molecular formula
molecular formula: gives the exact number of atoms of each element in the compound and is a multiple of the empirical formula
percent composition: the percent of a specific compound that is made up of a given element
percent composition = mass of element in formula/molar mass x 100
Chapter 4.3
combination reaction → has 2 or more reactants forming one product
A + B → C
Decomposition reaction: the opposite of a composition reaction → a single reactant breaks down into 2 or more products
A → B + C
Combustion reaction: a special type of reaction that involves a fuel-usually a hydrocarbon-and an oxidant (normally oxygen)
combustion involves oxidation
single-displacement reaction: occurs when an atom or ion in a compound is replaced by an atom or ion of another element
double-displacement reaction (metathesis): elements from 2 different compounds swap places with each other to form 2 new compounds
Neutralization reactions: a specific type of double displacement reaction in which an acid reacts with a base to produce salt
acid + base → salts > reaction is not always visible
Chapter 4.4
the mass of the reactants consumed must equal the mass of products generated
stoichiometric coefficients: the number placed in front of each compound → are used to indicate the relative # of moles of a given species involved in the reaction
Chapter 4.5
Limiting reagent: a reagent will be used up or consumed first; it limits the amount of product that can be formed in the reaction
excess reagents: reactants that remain after all the limiting reagent is used up
yield of reaction → refers to either the amount of product predicted or actually obtained when a reaction is carried out
Theoretical yield: the maximum amount of product that can be generated as predicted from the balanced equation; the amount of product predicted
actual yield: the amount of product one actually obtains during the reaction
percent yield = actual yield / theoretical yield x 100
Chapter 4.6
ionic bonds → holds ionic compounds together
rely on the force of electrostatic attraction between oppositely charged particles
Nomenclature of ionic compounds:
1. Elements (usually metals) can form more than one positive ion, the charge is indicated by a Roman numeral in parentheses following the name of the element
ex. Fe2+ → Iron (II) ; Cu2+ → copper(I)
2. An older, less commonly used method is to add the endings -ous or -ic to the root of the Latin name of the element to represent the ions with lesser and greater charge
ex. Fe2+ → ferrous ; Cu+ → cuprous
3. Monatomic anions are named by dropping the ending of the name of the element s and adding -ide
ex. H- → Hydride ; S2- → sulfide
4. Many polyatomic anions contain oxygen and are called oxyanions. When an element forms 2 oxyanions, the name of the one with less oxygen ends in -ite, and the one with more oxygen ends in -ate
ex. NO2- → Nitrite ; NO3- → nitrate
5. In an extended series of oxyanions, prefixes are also used. Hypo- and hyper, written as per-, are used to indicate less oxygen and more oxygen
ex. CIO- → hypochlorite; CIO2- → Chlorite
6. Polyatomic anions often gain one or more H+ ions to form anions of lower charge. The resulting ions are named by adding the word hydrogen or dihydrogen to the front of the anion’s hydrogen ion
ex. HCO3- → hydrogen carbonate/ bicarbonate
the color of a solution can be indicated of the oxidation state of a given element in the solution
solid ionic compounds tend to be poor conductors of electricity
electrolytes’ solutes that enable solutions to carry currents
ionic compounds make good electrolytes because they dissolve most readily
nonpolar covalent compounds are the weakest because they don’t form current-carrying ions
weak-electrolyte → ionizes or hydrolyzes incompletely in aqueous solution, and only some solute is dissolved into its constituents
Chapter 5.1
reactions can be spontaneous or nonspontaneous → the change in Gibbs free energy (delta G) determines whether or not a reaction will occur itself without outside assistance
rate-determining step: the slowest step in any proposed mechanism
the rate of the whole reaction is only as fast as the rate-determining step
collision theory of chemical kinetics: states that the rate of a reaction is proportional to the number of collisions per second between the reacting molecules
suggests that not all collisions result in a chemical reaction
an effective collision (one that leads to the formation of products) occurs only if the molecules collide with each other in the correct orientation and with sufficient energy to break their existing bonds and form new ones
activation energy/ energy barrier (Ea): the minimum energy of collision necessary for a reaction to take place
rate = Z x f
Z → total number of collisions occurring per second
f → the fraction of collisions that are effective
Arrhenius equation: k = -Ae^(Ea/RT)
k → rate constant
A → frequency factor
Ea → activation energy
R → ideal gas constant
T → temp (K)
As the frequency factor of the reaction increases, the rate constant of the reaction also increases
low activation energy and high temp make the negative exponent of the Arrhenius equation smaller in magnitude and thus increase the rate constant k
the frequency factor can be increased by increasing the number of molecules in a vessel → when there are more molecules, the opportunities for collision are increased
when molecules collide with energy equal to or greater than the activation energy, they form a transition state in which the old bonds are weakened and the new bonds begin to form
the transition state dissociates into products, fully forming the new bonds
Transition state (activated complex) → has greater energy than both the reactants and products
have the highest energy; the peak of the energy diagram
activation energy: the energy required to reach the transition state
free energy change of the reactions (delta Grxn): the difference between the free energy of the products and the free energy of the reactants
exergonic reaction: (energy is given off) → negative free energy change
endergonic reaction: (energy is absorbed) → positive free energy change
the greater the concentrations of the reactants, the greater the number of effective collisions per unit time
reaction rate will increase as the temperature increases
If the temp is too high, a catalyst may denature → reaction rate plummets
catalysts: substances that increase reaction rate without themselves being consumed in the reaction (interact with reactants)
homogenous catalyst → the catalysts is in the same phase (solid, liquid, and gas) as the reactants
heterogenous → catalysts is in a distinct phase
equilibrium: a dynamic process that seeks to find balance in all systems
5.2
to show a standard rate of reaction in which the rates with respect to all reaction species are equal, the rate of concentration