MCAT General Chemistry Review - Steven A. Leduc 2015
Glossary
absolute zero
The temperature (—273.15°C or 0 K) at which the volume and pressure of an ideal gas extrapolate to zero according to the ideal gas law. [Section 6.4]
acid
See Arrhenius acid, Brønsted-Lowry acid, and Lewis acid. [Section 11.1]
acid-base indicator
A weak acid and conjugate base pair, such as litmus or phenolphthalein, that changes color with changes in solution pH.
acid-dissociation constant (Ka)
A measure of the extent of dissociation of an acid, HA. Ka is defined as Ka = [H3O+][A−]/[HA]. [Section 11.3]
activation energy, Ea
The energy required by the reactants in a chemical reaction to reach a transition state from which the products of the reaction can form. [Section 6.6]
alkali metal
One of the metal elements in Group IA (Li, Na, K, Rb, Cs, and Fr). [Section 4.7]
alkaline earth metal
One of the metal elements in Group IIA (Be, Mg, Ca, Sr, Ba, and Ra). [Section 4.7]
allotropes
Forms of a pure element with different structures and therefore different chemical and physical properties, such as O2 and O3 or carbon(diamond) and carbon(graphite). [Section 6.6]
alloy
A blend of two or more metallic elements. Bronze, for example, is an alloy of copper and tin. [Section 10.4]
alpha (a) particle
A positively charged particle consisting of two protons and two neutrons emitted during alpha decay. An alpha particle is identical to a helium-4 nucleus. [Section 4.4]
amphoteric
An ion or molecule, such as H2O or HCO3−, that can act as either a Brønsted-Lowry acid or base. [Section 11.3]
anhydrous
A substance that is devoid of water. Used, for example, to differentiate between liquid (anhydrous) ammonia at temperatures below its boiling point (—33°C) and solutions of ammonia dissolved in water.
anion
A negatively charged ion, such as F−. [Section 4.3]
anode
The site of oxidation in a galvanic or electrolytic cell. Also, the electrode towards which anions flow through a salt bridge. [Section 12.2]
aqueous
Solutions of substances where water is the primary solvent. [Section 10.4]
Arrhenius acid
A compound that dissociates when it dissolves in water to give the H+ ion. [Section 11.1]
Arrhenius base
A compound that dissociates when it dissolves in water to give the OH− ion. [Section 11.1]
Arrhenius equation
The equation for the rate constant (k) of a chemical reaction, in terms of the temperature (T), activation energy (Ea), and collision frequency/steric orientation factor (A): k = Ae—Ea/RT. [Section 9.4]
atomic mass unit (amu, u)
The unit of mass most convenient for individual subatomic particles, atoms, and molecules. [Section 3.5]
atomic number (Z)
The number of protons in the nucleus of an atom. [Section 4.1]
atomic orbital
A region in space where electrons of an atom are most probably found. [Section 5.2]
atomic radius
The size of an atom equal to the volume of space carved out by the outermost (valence) electrons. [Section 4.8]
atomic weight
The weighted average of the atomic masses of the different isotopes of an element. A single 12C atom, for example, has a mass of 12 amu, but naturally occurring carbon also contains 1.1% 13C. The atomic weight of carbon is therefore 12.011 amu. [Section 3.5]
Aufbau principle
The principle that atomic orbitals are filled one at a time, starting with the orbital that has the lowest energy and then filling upwards. [Section 4.6]
Avogadro’s number
The number of items that make up one mole, approximately 6.02 × 1023. [Section 3.6]
Avogadro’s hypothesis
The hypothesis that equal volumes of different gases at the same temperature and pressure contain the same number of particles. [Section 8.2]
base
See Arrhenius base, Brønsted-Lowry base, and Lewis base. [Section 11.1]
base-dissociation constant (Kb)
A measure of the extent of dissociation of a base, B: Kb = [HB+][OH−]/[B]. [Section 11.3]
battery
A set of electrochemical cells connected in series or parallel. [Section 12.6]
beta (β) decay
A nuclear reaction in which an electron (e−) or positron (b+) is absorbed by or emitted from the nucleus of an atom. [Section 4.4]
beta (β) particle
An electron or a positron. [Section 4.4]
bimolecular
A step in a chemical reaction in which two molecules collide to form products.
Bohr model
A model of the distribution of electrons in an atom based on the assumption that the electron in a hydrogen atom is in one of a limited number of circular orbits. [Section 4.5]
Bohr atom
An atom or ion that has just one electron such as H, He+, Li2+, etc. [Section 4.5]
boiling point
The temperature at which the vapor pressure of a liquid is equal to the external or atmospheric pressure. [Section 5.8]
bond-dissociation energy/enthalpy
The energy needed to homolytically break an X—Y bond to give X and Y atoms in the gas phase. [Section 5.2]
bonding electrons
A pair of electrons, always the outermost (valence) electrons, used to form a covalent bond between adjacent atoms. [Section 5.1]
Boyle’s law
A statement of the relationship between the pressure and volume of a constant amount of gas at constant temperatures: P ∝ 1/V. [Section 8.2]
Brønsted-Lowry acid
Any molecule or ion that can donate an H+ (proton) to a base. [Section 11.1]
Brønsted-Lowry base
Any molecule or ion that can accept an H+ (proton) from an acid. [Section 11.1]
buffer
A mixture of a weak acid (HA) and its conjugate base (A−), which should be present in roughly equal amounts. Buffers resist a change in the pH of a solution according to Le Châtelier’s principle when small amounts of acid or base are added. [Section 11.8]
buffer capacity
The amount of acid or base a buffer solution can absorb without significant changes in pH. [Section 11.8]
calorie
The heat needed to raise the temperature of 1 gram of water by 1°C. [Section 7.2]
calorimeter
An insulated apparatus used to measure the heat absorbed/released in a chemical reaction.
catalyst
A substance that increases the rate of a chemical reaction without being consumed in the reaction. A substance that lowers the activation energy for a chemical reaction by providing an alternate pathway for the reaction. [Section 9.3]
cathode
The site of reduction in a galvanic or electrolytic cell. Also, the electrode in an electrochemical cell towards which cations flow through a salt bridge. [Section 12.2]
cation
A positively charged ion, such as Na+. [Section 4.3]
cell potential
A measure of the driving force behind an electrochemical reaction, reported in volts. [Section 12.3]
Charles’s law
A statement of the relationship between the temperature and volume of a constant amount of gas at constant pressure: V ∝ T. [Section 8.3]
collision theory model
A model used to explain the rates of chemical reactions, which assumes that molecules must collide in order to react. [Section 9.2]
common-ion effect
The decrease in the solubility of a salt that occurs when the salt is dissolved in a solution that contains another source of one of its ions. Just another form of Le Châtelier’s principle. [Section 10.6]
complex ion
An ion in which a ligand is bound to a metal via a coordinate covalent bond. An ion formed when a Lewis acid such as the Co2+ ion reacts with a Lewis base such as NH3 to form an acid-base complex such as the Co(NH3)42+ ion. [Section 10.7]
compound
A substance with a constant composition that contains two or more elements.
concentration
A measure of the ratio of the amount of solute in a solution to the amount of either solvent or solution. Frequently expressed as molarity (units of moles of solute per liter of solution). [Section 3.8]
concentration cell
A type of electrochemical cell that has identical reactants in each half reaction, but at different concentrations, thus driving a weak electrical current. [Section 12.4]
conjugate acid-base pair
Two substances related by the gain or loss of a proton. An acid (such as HBr) and its conjugate base (Br−), or a base (such as NH3) and its conjugate acid (NH4+) are examples of conjugate acid-base pairs. [Section 11.3]
coordinate covalent bond
A covalent bond formed as a result of a Lewis acid-base reaction, most often formed between a metal atom and a nonmetal atom. [Section 5.3]
coordination compound
A compound in which one or more ligands are coordinated to a metal atom. [Section 5.3]
corrosion
A process in which a metal is destroyed by a chemical reaction. When the metal is iron, the process is called rusting.
covalent bond
A bond between two atoms formed by the sharing of at least one pair of electrons. [Section 5.3]
covalent compound
A compound, such as water (H2O), composed of neutral molecules in which the atoms are held together by covalent bonds. [Section 5.3]
critical point
The temperature and pressure at which two phases of a substance that are in equilibrium (usually the gas and liquid phases) become identical and form a single phase. [Section 7.4]
crystal
A three-dimensional solid formed by regular repetition of the packing of atoms, ions, or molecules. [Section 5.8]
Dalton’s law of partial pressures
A statement of the relationship between the total pressure of a mixture of gases and the partial pressures of the individual components: Ptotal = p1 + p2 + p3 +…. [Section 8.4]
daughter nucleus
The product nucleus after a nuclear reaction. [Section 4.4]
density
The mass of a sample divided by its volume. [Section 3.2]
deposition
A process in which a gas goes directly to the solid state without passing through an intermediate liquid state. [Section 7.1]
diamagnetic
A substance in which the electrons are all paired. [Section 4.6]
dilution
The process by which more solvent is added to decrease the concentration of a solution.
dimer
A compound (such as B2H6) produced by combining two smaller identical molecules (such as BH3).
dipole
Anything with two equal but opposite electrical charges, such as the positive and negative ends of a polar bond or molecule. [Section 5.6]
diprotic acid
An acid, such as H2SO4, that can lose two H+.
diprotic base
A base, such as the O2— ion, that can accept two H+.
dissolution
The process in which a bulk solid or liquid breaks up into individual molecules or ions and diffuses throughout a solvent. [Section 10.4]
ductile
Capable of being drawn into thin sheets or wires without breaking; a property of metals. [Section 5.8]
effusion
The process by which a gas escapes through a pinhole into a region of lower pressure. [Section 8.5]
elastic collision
A collision in which no kinetic energy is lost. [Section 8.1]
electrolyte
A substance that increases the electrical conductivity of water by dissociating into ions. [Section 10.1]
electrolysis
A process in which an electric current is used to drive a nonspontaneous chemical reaction. [Section 12.6]
electrolytic cell
A nonspontaneous electric cell in which electrolysis is done. [Section 12.6]
electron
A subatomic particle with a mass of only about 0.0005 amu and a charge of —1e that surrounds the nucleus of an atom. [Section 4.1]
electron affinity
The energy given off when a neutral atom in the gas phase picks up an electron to form a negatively charged ion. [Section 4.8]
electron capture
A type of beta decay where the nucleus of an atom captures an electron and converts a nuclear proton into a neutron. [Section 4.4]
electron configuration
The arrangement of electrons in atomic orbitals; for example, 1s22s22p3. [Section 4.6]
electronegativity
The tendency of an atom to draw or polarize bonding electrons toward itself. [Section 4.8]
element
A substance that cannot be decomposed into a simpler substance by a chemical reaction. A substance composed of only one kind of atom. [Section 3.5]
empirical formula
The formula for a compound in which the number of atoms of each element in the compound are represented by the lowest whole number ratio.[Section 3.4]
endergonic
A process that leads to an increase in the free energy of a system and is therefore not spontaneous: ∆Grxn is positive. [Section 6.6]
endothermic
A chemical reaction that absorbs heat from the surroundings: ∆Hrxn is positive. [Section 6.2]
endpoint
The point at which the indicator of an acid-base titration changes color.
enthalpy (H)
The total potential energy in a substance due to intermolecular forces and covalent bonds. [Section 6.2]
enthalpy of reaction (∆Hrxn)
The change in the enthalpy that occurs during a chemical reaction. The difference between the sum of the enthalpies of the products and the reactants. [Section 6.2]
entropy (S)
A measure of the disorder in a system. Increasing disorder yields a positive ∆S. [Section 6.4]
equilibrium (dynamic)
The point at which there is no longer a change in the concentrations of the reactants and the products of a chemical reaction. The point at which the rates of the forward and reverse reactions are equal: ∆Grxn = 0. [Section 10.4]
equilibrium constant (Keq)
The product of the concentrations (or partial pressures) of the products of a reaction at equilibrium divided by the product of the concentrations (or partial pressures) of the reactants. [Section 10.1]
equilibrium expression
The expression used to calculate the equilibrium constant for a reaction that takes the form [products]/[reactants]. [Section 10.1]
equivalence point
The point in an acid-base titration at which the number of moles of H3O+ in solution equals the number of moles of OH− in solution. [Section 11.10]
excited state configuration
One of an infinite number of electron configurations of an energized atom where at least one electron occupies an orbital of higher energy than that dictated by Hund’s rule and/or the Aufbau principle. [Section 4.5]
exergonic
A process that leads to a decrease in the free energy of the system and is therefore spontaneous: ∆Grxn is negative. [Section 6.6]
exothermic
A chemical reaction that releases energy to the surroundings: ∆Hrxn is negative. [Section 6.2]
family
A vertical column of elements in the periodic table, such as the elements Li, Na, K, etc. [Section 4.6]
Faraday’s law of electrolysis
A statement of the relationship between the amount of electric current that passes through an electrolytic cell and the amount of product formed during electrolysis. The amount of chemical change is proportional to the amount of electric current that flows through the cell. [Section 12.7]
first ionization energy
The energy needed to remove the valence electron from a neutral atom in the gas phase. [Section 4.8]
first law of thermodynamics
The total energy in the universe is conserved: energy is neither created nor destroyed, but may change from one form to another. [Section 6.1]
first-order reaction
A reaction in which the rate is proportional to the concentration of a single reactant raised to the first power: rate = k[A]. [Section 9.4]
formal charge
The theoretical charge on an atom in a molecule, calculated by V — 
B — L, where V is the number of valence electrons, B is the number of bonding electrons, and L is the number of lone-pair electrons. [Section 5.1]
free energy, Gibbs (G)
The energy associated with a chemical reaction that can be used to do work. The change in free energy of a system is calculated by the formula ∆G = ∆H — T∆S, where G is the free energy, H is enthalpy, T is temperature (in kelvins), and S is entropy. [Section 6.5]
free radical
An atom or molecule that contains an unpaired electron. [Section 5.2]
freezing point
The temperature at which the solid and liquid phases of a substance are in equilibrium. [Section 7.1]
fusion
The melting of a solid to form a liquid. [Section 7.1]
galvanic cell
An electrochemical cell that uses a spontaneous chemical reaction to do work. Synonymous with voltaic cell. [Section 12.2]
gamma ray (γ)
A high energy, short wavelength form of electromagnetic radiation emitted by the nucleus of an atom that carries off some of the energy generated in a nuclear reaction. [Section 4.4]
Gibbs free energy
See free energy. [Section 6.5]
Graham’s law
The relationship between the rate at which a gas diffuses or effuses and its molecular weight: rate ∝ 1/(MW)½. [Section 8.5]
ground state configuration
The most stable arrangement of electrons in an atom that satisfies Hund’s rule and the Aufbau principle. [Section 4.5]
group
A vertical column, or family, of elements in the periodic table. [Section 4.7]
half-life
The time required for the amount of a decaying substance to decrease to half its initial value. [Section 4.4]
halogen
Elements of Group VIIA: F, Cl, Br, I, and At. [Section 4.7]
heat (q)
Thermal energy in transit from a hotter system to a colder one. [Section 6.1]
heat of fusion
The heat that must be absorbed to melt a unit quantity of a solid. [Section 7.2]
heat of reaction
The change in the enthalpy of the system that occurs when a reaction is run at constant pressure. Synonymous with enthalpy of reaction, ∆Hrxn. [Section 6.2]
heat of vaporization
The heat that must be absorbed to boil a unit quantity of a liquid. [Section 7.2]
heat capacity
The amount of heat required to raise the temperature of a given amount of a substance by one degree. Not to be confused with specific heat. Heat capacity is typically the product of mass and specific heat. [Section 7.3]
Hess’s law
The heat given off or absorbed in a chemical reaction does not depend on whether the reaction occurs in a single step or in many steps. [Section 6.3]
homonuclear diatomic molecule
A molecule, such as O2 or F2, that contains two atoms of the same element.
Hund’s rule
Rule for placing electrons in equal-energy orbitals, which states that electrons are added with parallel spins until each of the orbitals has one electron, before a second electron is placed in a given orbital. [Section 4.6]
hybrid orbitals
Orbitals formed by mixing two or more atomic orbitals. [Section 5.5]
hybridization
A process in which things are mixed. A resonance hybrid is a mixture, or average, of two or more Lewis structures. Hybrid orbitals are formed by mixing two or more atomic orbitals. [Section 5.5]
hydride
The species H−. [Section 5.4]
hydrogen bonding
A strong dipole-dipole interaction that occurs between a hydrogen atom covalently bonded to an F, O, or N that electrostatically interacts with a lone pair of electrons on another F, O, or N atom. [Section 5.7]
hydrophilic
To be attracted or compatible with water (for example, ions that are soluble in water).
hydrophobic
To be incompatible with water (for example, lipids insoluble in water).
ideal gas
A gas that obeys all the postulates of the kinetic-molecule theory and has properties that can be predicted by the ideal gas law. [Section 8.1]
ideal gas law
The relationship between the pressure, volume, temperature, and amount of an ideal gas: PV = nRT. [Section 8.2]
immiscible
Liquids, such as oil and water, that do not dissolve in one another.
indicator
See acid-base indicator. [Section 11.9]
induced dipole
A short-lived separation of charge, or dipole, of a nonpolar atom or molecule caused by the electrostatic influence of a nearby polar atom or ion. [Section 5.7]
inert
Unreactive. Used to describe compounds that do not undergo chemical reactions. [Section 4.7]
insoluble
Used to describe a substance that does not noticeably dissolve in a solvent. [Section 10.4]
intermolecular forces
Attractive electrostatic forces, the strength of which determine a compound’s phase, vapor pressure, melting point, boiling point, solubility, and viscosity. From strongest to weakest, the main categories of intermolecular forces are: ionic, dipole-dipole (with H-bonds the strongest), and London dispersion forces. [Section 5.7]
intramolecular bonds
Synonym for covalent bonds. There are three primary types: normal covalent bonds, metallic covalent bonds, and coordinate covalent bonds. [Section 7.1]
ion product (Qsp)
The product of the concentrations of the ions in a solution at any moment. [Section 10.5]
ion-product constant of water (Kw)
The product of the equilibrium concentration of the H3O+ and OH− ions in an aqueous solution at 25°C: Kw = 1.0 × 10−14. [Section 11.4]
ionic bond
Misappropriation of the term bond. Simply the strong electrostatic attraction between two oppositely charged ions; there is no electron sharing. [Section 5.3]
ionic compound
A compound made up of ions (synonymous with salt). [Section 3.5]
ionizability factor (i)
The number of individual particles formed when an individual solute dissolves. Synonymous with van’t Hoff factor. [Section 10.4]
ionization
A process in which an ion is created from a neutral atom or molecule by adding or removing one or more electrons. [Section 4.8]
ionization energy
See first ionization energy. [Section 4.8]
isoelectronic
Atoms or ions that have the same number of electrons and therefore the same electron configuration, such as O2—, F−, Ne, and Na+. [Section 4.6]
isotopes
Nuclides of the same element, but with differing numbers of neutrons, such as 12C, 13C, and 14C. Isotopes have nearly identical chemical properties. [Section 4.2]
joule
A unit of measurement for both heat and work in the SI system. 1 J = 4.184 cal. [Section 4.2]
kinetic energy
The energy associated with motion. The kinetic energy of an object is equal to one-half the product of its mass and the square of its speed: KE = mv2. [Section 5.8]
kinetic-molecular theory
The theory that states heat is associated with the thermal motion of particles, taking into account the important assumptions that individual gas molecules take up no volume and collisions between gas molecules are perfectly elastic. [Section 8.1]
Le Châtelier’s principle
A principle that describes the effect of changes in the temperature, pressure, or concentration of one of the reactants or products of a reaction at equilibrium. It states that when a system at equilibrium is subjected to a stress, it will shift in the direction that minimizes the effect of this stress. [Section 10.3]
Lewis acid
An atom or molecule that accepts a pair of electrons to form a new coordinate covalent bond. Almost always a metal atom, positively charged ion, or both. [Section 5.3]
Lewis base
An atom or molecule that donates a pair of electrons to form a new coordinate covalent bond. Almost always a nonmetal with a pair of nonbonding electrons. Synonymous with ligand and chelator. [Section 5.3]
ligand
See Lewis base. [Section 5.3]
limiting reagent
The reactant in a chemical reaction that is exhausted first, thus limiting the amount of product that can be formed. [Section 3.11]
London dispersion forces
Intermolecular forces that arise from interactions between an instantaneous dipole/induced dipole pair. Typically, these are the weakest of all intermolecular forces. However LDFs are additive, and nonpolar molecules with large, flat surface areas can experience moderate LDFs. [Section 5.7]
malleable
Something that can be hammered, pounded, or pressed into different shapes without breaking (a common property of metals). [Section 5.8]
mass number (A)
The total number of protons and neutrons in the nucleus of an atom. [Section 4.1]
melting point
The temperature at which the solid and liquid phases of a substance are in equilibrium at a particular external pressure. [Section 5.7]
metal
An element that is solid, has a metallic luster, is malleable and ductile, and conducts both heat and electricity. [Section 3.9]
metalloid
An element with properties that fall between the extremes of metals and nonmetals. [Section 4.7]
mixture
A substance that contains two or more elements or compounds that retain their chemical identities and can be separated by a physical process. For example, the mixture of N2 and O2 in the atmosphere. [Section 3.8]
molarity (M)
The number of moles of a solute in a solution divided by the volume of the solution in liters. [Section 3.1]
mole
6.02 × 1023 of anything. [Section 3.6]
mole fraction (X)
The fraction of the total number of moles in a mixture due to one component of the mixture. The mole fraction of a solute, for example, is the number of moles of solute divided by the total number of moles of solute plus solvent. [Section 3.8]
mole ratio
The ratio of the moles of one reactant or product to the moles of another reactant of product in the balanced equation for a chemical reaction.
molecular formula
The formula representing the number and type of constituent atoms in a compound. [Section 3.3]
molecular geometry
The arrangement of atoms surrounding a central atom of a small molecule. Molecular geometry (or the shape of a molecule) and orbital geometry are identical only when the central atom possesses no lone pairs of electrons. [Section 5.4]
molecular weight
The weight of the molecular formula, calculated from a table of atomic weights. Note that atomic weights are a weighted average of masses of isotopes as they occur in nature. [Section 5.7]
molecule
The smallest particle that has any of the chemical or physical properties of a compound. [Section 3.3]
monoprotic acid (HA)
An acid, such as HF or HOCl, that can lose only one H+.
negative electrode
The electrode in an electrochemical cell that carries a negative charge. In a galvanic cell, it is the anode; in an electrolytic cell, it is the cathode.
Nernst equation
Used to calculate or track the voltage of an electrochemical cell under nonstandard conditions: E = E° — (0.06/n) log Q. [Section 12.4]
network solid
A solid, such as diamond, in which every atom is covalently bonded to its nearest neighbors to form an extended array of atoms rather than individual molecules. [Section 5.8]
neutron
A subatomic particle with a mass of about 1 amu and no charge. [Section 4.1]
noble gases
The elements in the last column of the periodic table that are chemically unreactive. [Section 4.6]
nonbonding electrons
Electrons in the valence shell of an atom that are not used to form covalent bonds. [Section 5.1]
nonmetal
An element that lacks the properties generally associated with metals. These elements are found in the upper right of the periodic table. [Section 4.7]
nonpolar
Used to describe a compound that has a homogenous electron distribution and thus does not carry a permanent dipole moment. [Section 5.3]
nonspontaneous
A reaction in which the products are not favored, implying that the reverse reaction would be favored: ∆Grxn is positive (and Ecell is negative). [Section 2.6]
nucleon
Generic term for a proton or neutron. [Section 4.1]
nuclide
The generic term for any particular isotope of an element, such as the 13C nuclide.
octet rule
The tendency of main-group elements to react in order to possess eight valence-shell electrons in their compounds.
orbital geometry
The arrangement of electron clouds surrounding a central atom of a small molecule. Orbital geometry is a consequence of hybridization. Not to be confused with shape, or molecular geometry. [Section 5.4]
orbitals
Regions in space where electrons have a high probability of existing. [Section 4.5]
order
Used to describe the relationship between the rate of a step in a chemical reaction and the concentration of one of the reactants consumed in that step. Essentially just the value of the exponent found in a reactant term in the rate law. [Section 9.4]
oxidation
A process in which an atom, ion, or molecule loses one or more electrons. [Section 3.13]
oxidation number
Synonymous with oxidation state. It is the hypothetical charge that would be present on each atom if a molecule was shattered into its individual constituent atoms with bonding electrons ending up with the atom of the bond having the higher electronegativity. [Section 3.13]
oxidation-reduction reaction
A chemical reaction involving the exchange of electrons such that oxidation numbers of reactants change. [Section 12.1]
oxidizing agent / oxidant
An atom, ion, or molecule that undergoes reduction by gaining electrons, thereby oxidizing something else. [Section 4.7]
pH
A measure of acidity ranging from about —1.5 to 15.5 in aqueous media, defined as —log [H3O+]. [Section 11.5]
pOH
The complement of pH, defined as —log [OH−]. [Section 11.5]
paramagnetic
A compound that contains one or more unpaired electrons and is attracted into a magnetic field. [Section 4.6]
parent nucleus
The initial nucleus prior to a nuclear reaction. [Section 4.4]
partial pressure
The fraction of the total pressure of a mixture of gases that is due to one component of the mixture. As molarity is the primary way of expressing the amount of solute in a solution, so too is partial pressure the primary way to report the quantity of a gas in a mixture of gases. [Section 8.4]
Pauli exclusion principle
The maximum number of electrons in any given orbital is two, and they must have the opposite spin. [Section 4.6]
period
A horizontal row in the periodic table. [Section 4.6]
polar covalent bond
A covalent bond between atoms with differing electronegativities such that electrons spend more time in the vicinity of one atom than the other.
polar
Used to describe a molecule that has a dipole moment because it consists of a positive pole and a negative pole. [Section 5.3]
polyatomic ion
An ion that contains more than one atom, such as CO32— or SO42—.
polyprotic acid
An acid, such as H2SO4 or H3PO4, that can lose more than one H+. [Section 11.4]
polyprotic base
A base, such as the PO43— ion, that can accept more than one H+.
positive electrode
The electrode in an electrochemical cell that carries a positive charge. In a galvanic cell, it is the cathode; in an electrolytic cell, it is the anode.
positron (β+)
The antiparticle of the electron. A positron has the same mass as an electron but is positively charged. Contact between an electron and positron results in instant annihilation and emission of two high energy gamma rays. [Section 4.4]
positron emission
A mode of beta decay where a positron is emitted as a consequence of the conversion of a nuclear proton to a neutron. [Section 4.4]
potential
A measure of the driving force behind an electrochemical reaction that is reported in units of volts. [Section 12.3]
precipitation
A process where dissolved ions combine to form a solid salt in solution. [Section 10.4]
precision
A measure of the extent to which individual measurements of the same phenomenon agree.
pressure
The force exerted perpendicular to a surface divided by the area of the surface. [Section 5.7]
proton
A subatomic particle that has a charge of +1e and a mass of about 1 amu. (Synonymous with H+.) [Section 4.1]
quantized
A property or quality that appears only in certain discrete amounts, such as electric charge. [Section 4.5]
radioactivity
The spontaneous disintegration of an unstable nuclide by a first-order rate law. Synonymous with nuclear decay. [Section 4.4]
rate of reaction
The change in the concentration of a compound divided by the amount of time necessary for this change to occur: rate = ∆[X]/∆t. [Section 9.1]
rate constant (k)
The proportionality constant in the equation that describes the relationship between the rate of a step in a chemical reaction and the product of the concentrations of the reactants consumed in that step. [Section 9.4]
rate law
An equation that describes how the rate of a chemical reaction depends on the concentrations of the reactants consumed in that reaction, along with the rate constant that takes into account temperature, activation energy, and collision frequency / steric effects. [Section 9.4]
rate-determining step
The slowest step in a chemical reaction. [Section 9.1]
reaction quotient (Q)
The quotient obtained when the concentrations (or partial pressures) of the products of a reaction are multiplied and the result is divided by the product of the concentrations (or partial pressures) of the reactants. Basically, putting nonequilibrium values into an equilibrium expression yields a reaction quotient instead of the equilibrium constant. [Section 10.2]
real gas
A gas that deviates from the behavior predicted by the ideal gas law. Real gases differ from the expected behavior of an ideal gas (e.g., lower V and P) for two reasons: (1) the forces of attraction between the particles in a gas are not zero and (2) the volume of the particles in a gas is not zero. [Section 8.3]
redox
An abbreviation for oxidation-reduction. [Section 12.1]
reducing agent / reductant
An atom, ion, or molecule that is oxidized by giving up electrons, thereby reducing something else. [Section 12.1]
reduction
A process in which an atom, ion, or molecule gains one or more electrons. [Section 12.1]
resonance structures
Two or more Lewis dot structures that differ only by the placement of electrons in the molecule. Taken together as an average (a resonance hybrid), they best approximate the electron distribution and types of bonds in the molecule better than any one structure can alone. [Section 5.1]
salt
Synonymous with ionic compound. [Section 5.8]
salt bridge
An ion-rich junction between the anodic and cathodic chambers of an electrochemical cell that prevents charge separation that would otherwise stop the cell from functioning. Anions always migrate toward the anode, and cations always migrate toward the cathode of any cell. [Section 12.2]
saturated solution
A solution that contains as much solute as possible. [Section 10.4]
second ionization energy
The energy needed to remove an electron from a +1 cation in the gas phase. [Section 4.8]
second law of thermodynamics
Processes that increase the entropy in the universe are spontaneous. [Section 6.4]
second-order reaction
A reaction in which rate is proportional to the concentration of a single reactant raised to the second power: rate = k[A]2, or two reactants each raised to the first power: rate = k[A][B].
shielding
The masking and weakening of the electrostatic attraction between the nucleus and outer electrons by inner electrons. [Section 4.8]
solubility
The ratio of the maximum amount of solute to the volume of solvent in which this solute can dissolve. Often expressed in units of grams of solute per 100 g of water, or in moles of solid per liter of solution. [Section 10.4]
solubility equilibria
Equilibria that exist in a saturated solution, in which additional solid dissolves at the same rate that particles of solution come together to precipitate more solid.
solubility product (Ksp)
The product of the equilibrium concentrations of the ions in a saturated solution of a salt. [Section 10.1]
solute
The substance that dissolves in a solvent to form a solution. [Section 10.4]
solution
A homogeneous mixture of one or more solutes dissolved in a solvent. [Section 10.4]
solvent
The substance in which a solute dissolves. [Section 10.4]
specific heat
The amount of heat required to raise the temperature of 1 g of a substance by 1°C (or 1 K). (Do not confuse with heat capacity). [Section 7.2]
spontaneous reaction
A reaction in which the products are favored: ∆Grxn is negative (and Ecell is positive). [Section 6.5]
standard cell potential
The potential, E°cell, of a cell measured under standard-state conditions.
standard heat of formation (∆Hf°)
The change in the enthalpy that occurs during a chemical reaction that leads to the formation of one mole of a compound from its elements in their standard states at standard conditions. [Section 6.3]
standard state/condition
State in which T = 298 K (25°C), P = 1 atm, and all concentrations are 1 M. Not to be confused with STP. [Section 6.3]
standard temperature and pressure (STP)
State in which T = 273 K (0°C) and P = 1 atm. Generally used when referring to gases. Not to be confused with standard state or standard conditions. [Section 6.3]
state
1. One of the three states of matter: gas, liquid, or solid.
2. A set of physical properties that describe a system. [Section 3.12]
state function
A quantity whose value depends only on the state of the system and not its history; X is a state function, if and only if, the value of ∆X does not depend on the path used to go from the initial to the final state of the system. [Section 6.3]
stoichiometry
The study of the quantitative relationships between the reactants and the products of a balanced chemical reaction. [Section 3.10]
strong acid
An acid that dissociates completely in water. [Section 11.3]
strong base
A base that dissociates completely in water. [Section 11.3]
sublimation
The process in which a solid goes directly to the gas phase without passing through an intermediate liquid state. [Section 6.4]
supercooled liquid
A substance that is a liquid even though its temperature is below its freezing point.
supercritical fluid
A substance that displays properties of both a liquid and a gas and exists under conditions of high temperature and pressure. If a substance is in this state—where the liquid and gas phases are no longer distinct—no amount of increased pressure can force the substance back into its liquid phase. [Section 7.4]
surface tension
The perpendicular force per unit length of liquid surface that acts to reduce the surface area of a liquid, resulting from intermolecular forces below the surface.
surroundings
In thermodynamics, the part of the universe not included in the system. [Section 6.1]
system
In thermodynamics, that small portion of the universe in which we are interested at the moment. [Section 6.1]
thermal conductor
A substance or object that readily conducts heat (metals, for example).
thermal insulator
An object, such as a blanket or a fur coat, that tends to slow down the rate at which heat is transferred from one object to another.
titrant
The strong acid or base reagent added to the unknown solution in a titration experiment. [Section 11.10]
titration
A technique used to determine the concentration and/or the chemical identity of a solute in a solution. [Section 11.10]
torr
A unit of pressure equal to the pressure exerted by a column of mercury 1 millimeter tall. By definition, 1 torr = 1 mm Hg. [Section 8.1]
transition metal
Metals in the block of elements that serve as a transition between the two columns on the left side of the table, where s orbitals are filled, and the six columns on the right, where p orbitals are filled. [Section 3.13]
triple point
The unique pressure and temperature at which the three phases of a substance (gas, liquid, and solid) can all coexist in equilibrium. [Section 7.4]
unimolecular
Describes a step in a reaction mechanism in which only one reactant molecule is present. [Section 9.4]
valence electrons
Electrons in the outermost or highest-energy level or shell of an atom. The electrons that are gained, shared, or lost in a chemical reaction. [Section 3.13]
van der Waals equation
An equation that accounts for deviations from ideal behavior in gaseous systems due to interactions between gas particles and particle volume. [Section 8.3]
van’t Hoff factor (i)
See ionizability factor. [Section 10.4]
vapor pressure
The partial pressure of the gas molecules over the surface of a liquid that originate from the surface of a liquid. [Section 5.7]
voltaic cell
An electrochemical cell in which a spontaneous chemical reaction is used to create electricity. Synonymous with galvanic cell. [Section 12.2]
volatile
The physical characteristic of having a high vapor pressure at standard conditions. [Section 5.7]
VSEPR theory
A model in which the repulsion between pairs of valence electrons is used to predict the shape of a molecule. [Section 5.4]
weak acid
An acid that only partially dissociates in water. [Section 11.3]
weak base
A base that only partially dissociates in water. [Section 11.3]
zero-order reaction
A reaction in which rate is not proportional to the concentration of any of the reactants.