AP Chemistry: Atomic Structure, Bonding, and Thermodynamics Review

Comprehensive flashcards covering atomic structure, periodic trends, chemical bonding, inter-molecular forces, kinetics, thermodynamics, and equilibrium based on lecture notes.

Mass number

The sum of protons and neutrons in an atom ($P + N$).

Isotopes

Atoms of an element that have different numbers of neutrons.

Avogadro’s number

A constant represented by $6.022 imes 10^{23}$.

STP (Standard Temperature and Pressure)

Conditions defined as $1\,\text{atm}$ and $273\,\text{K}$, where the molar volume of a gas is $22.4\,\text{L/mol}$.

Molarity ($M$)

A measure of concentration defined as $\text{moles}/\text{L}$.

Empirical formula

The simplest ratio of elements in a chemical substance.

Coulomb’s Law

The electrostatic force given by the equation $F = \frac{k q_1 q_2}{r^2}$.

Photoelectron Spectroscopy

A technique measuring energy in electronvolts ($\text{eV}$) where the incoming radiation energy equals the binding energy plus the kinetic energy of the ejected electron.

Aufbau principle

The rule stating that electrons fill the lowest energy subshells available first.

Pauli exclusion principle

The rule stating that two electrons in the same orbital cannot have the same spin.

Hund’s rule

The rule stating that electrons occupy empty subshells first.

Ionization energy

The amount of energy required to remove an electron from an atom.

Electronegativity

How strongly the nucleus of an atom attracts electrons of other atoms within a bond.

Electron affinity

The energy change that occurs when an electron is added to an atom in the gas state, usually appearing as an exothermic process.

Ionic bonds

Bonds formed when a cation gives up electrons completely to an anion, resulting in electrostatic attractions in a lattice structure.

Metallic bonds

A sea of electrons model where a positively charged core remains stationary while valence electrons are highly mobile.

Single bond

A bond consisting of one sigma bond (order 1), characterized by having the longest length and the least bond energy.

Triple bond

A bond consisting of one sigma and two pi bonds (order 3), characterized by being the shortest and having the greatest bond energy.

Resonance

The process of averaging together all possible bond orders for bond order calculations of a specific bond.

Formal charge

The number of valence electrons minus the assigned electrons (including one electron for each shared bond).

Dipole-dipole forces

Attractions between polar molecules where the positive end of one molecule is attracted to the negative end of another.

Hydrogen bonds

A strong type of dipole-dipole attraction occurring when hydrogen is attracted to an extremely electronegative element, specifically $F$, $O$, or $N$.

London dispersion forces (LDF)

Weak attractions present in all molecules due to the random motion of electrons creating instantaneous polarity.

Vapor pressure

The pressure exerted by molecules that escape IMF at the surface of a liquid; it is directly proportional to temperature and inversely proportional to IMF strength.

Beer’s Law

Defined by $A = abc$, showing a direct relationship between the concentration ($c$) and absorbance ($A$) of a solution.

Oxidation-reduction (redox) reaction

A reaction that involves changes in the oxidation state of the chemical species involved.

Rate law

The equation $\text{Rate} = k [A]^x [B]^y [C]^z$, where the constant $k$ is dependent only on temperature.

Activation energy ($E_a$)

The sufficient energy required for reactants to collide and successfully undergo a chemical reaction.

Reaction intermediates

Species that are produced during a reaction mechanism but are fully consumed and do not appear in the final balanced equation.

Catalyst

A substance that increases the reaction rate without being consumed by providing a pathway with a lower activation energy.

First Law of Thermodynamics

The law stating that energy can be neither created nor destroyed.

Enthalpy of formation ($\Delta H_f$)

The change in energy when one mole of a compound is formed from its pure component elements in their standard states.

Hess’s Law

A method of finding the overall $\Delta H$ by adding the enthalpy changes of individual reaction steps together.

Specific heat

The amount of heat required to raise the temperature of one gram of a substance by one degree $C$ or $K$, used in the equation $q = mc\Delta T$.

Le Chatelier’s principle

The principle stating that a system at equilibrium will shift to favor products or reactants to counteract changes in concentration, pressure, or temperature.

Reaction quotient ($Q$)

A value calculated using initial molar concentrations or partial pressures, used to determine which direction a reaction must shift to reach equilibrium ($K$).

Solubility product ($K_{sp}$)

The equilibrium constant for the dissociation of a solid salt into its aqueous ions.

Strong acids

Acids that dissociate completely in water, leaving no tendency for a reverse reaction to occur.

Buffer

A solution of a weak acid or base and its conjugate salt that resists changes in $\text{pH}$ when additional acid or base is added.

Entropy ($S$)

A measurement of the randomness or dispersion within a system.

Gibbs free energy ($\Delta G$)

Determined by $\Delta G = \Delta H - T \times \Delta S$; a negative value indicates the reaction is spontaneous or thermodynamically favored.

Galvanic cells

Also called voltaic cells, these use a favored redox reaction to generate a flow of current, with oxidation at the negative anode and reduction at the positive cathode.

Faraday’s Constant ($F$)

A value of $96500\,\text{C/mol}$ of electrons used in electrochemistry calculations.