AP Chemistry Ultimate Guide

Periodic Table

Periodic Table

Provides basic information about each element, including symbol, atomic number, and molar mass.

Atomic Number

Indicates the number of protons and neutrons in an element, as well as the electrons in a neutral atom.

Isotopes

Atoms of an element with different numbers of neutrons but the same number of protons.

Moles

A unit connecting different quantities in chemical equations, where 1 mole equals 6.022x10^23 particles.

Molarity

Expresses the concentration of a solution in terms of volume, used in various chemical calculations.

Electron Configuration

Describes the distribution of electrons in different subshells (s, p, d, f) based on energy levels.

Ionization Energy

Energy required to remove an electron from an atom, increases from left to right in the periodic table.

Coulomb's Law

Describes the electrostatic force between charges, influencing the attraction between electrons and the nucleus.

Atomic Radius

Approximate distance from the nucleus to the valence electrons, decreases from left to right in a period.

Ionic Charges

Elements gain or lose electrons to achieve stability, forming ions with positive (cations) or negative (anions) charges.

Ionization Energy

Energy required to remove an electron from an atom.

Electronegativity

Measure of an element's ability to attract electrons.

Second Ionization Energy

Energy required to remove a second electron from an atom.

Fluorine

The most electronegative element due to its small size and need for one electron to complete its shell.

Ionic Bonds

Bonds between a metal and a nonmetal where electrons are transferred.

Metallic Bonds

Bonds where electrons move freely among metal atoms.

Covalent Bonds

Bonds where electrons are shared between nonmetal atoms.

Lewis Dot Structures

Diagrams showing the arrangement of valence electrons in a molecule.

Molecular Polarity

Unequal sharing of electrons leading to partial charges in a molecule.

Intermolecular Forces

Forces between molecules affecting phase changes.

Dipole-Dipole Forces

Forces between polar molecules with partial charges.

Hydrogen Bonds

Strong dipole-dipole forces between hydrogen and electronegative atoms.

London Dispersion Forces

Weak forces between all molecules due to random electron movements.

IMF Strength

Ranking of intermolecular forces from strongest to weakest.

Vapor Pressure

Pressure exerted by a vapor in equilibrium with its liquid phase.

Boiling vs

Boiling requires added heat to break IMFs, while vaporization occurs without added heat.

Solution Separation

Substances can be separated based on different IMFs and Coulombic attractions of molecules.

Solutes and Solvents

Like dissolves like - polar substances dissolve in polar solvents, nonpolar in nonpolar solvents.

Paper Chromatography

Technique to separate mixtures based on substance polarity using a solvent on paper.

Retention Factor (Rf)

Measure in chromatography, Rf = (distance traveled by solute) / (distance traveled by solvent front).

Kinetic Molecular Theory

Describes ideal gases' behavior, stating KE is proportional to temperature, and gas molecules have no forces of attraction.

Maxwell-Boltzmann Diagrams

Show the range of velocities for gas molecules at different temperatures.

Effusion

Rate at which a gas escapes from a container through microscopic holes based on molecule speed, temperature, and molar mass.

Ideal Gas Equation

PV = nRT, where P is pressure, V is volume, n is moles, R is the gas constant, and T is temperature.

Dalton's Law

States that the total pressure of a gas mixture is the sum of the partial pressures of individual gases.

Deviations From Ideal Behavior

Factors like low temperature, high pressure, or strong IMFs can cause gases to deviate from ideal behavior.

Density

Density of a gas can be calculated using D = m/V, and molar mass from density using MM = DRT/P.

Electromagnetic Spectrum

Describes the relationship between energy change, frequency, and wavelength of electromagnetic radiation.

Beer’s Law

Relates absorbance of a solution to its molar absorptivity, path length, and concentration.

Types of Reactions

Includes synthesis, decomposition, acid-base, oxidation-reduction, hydrocarbon combustion, and precipitation reactions.

Half-Life

The time taken for half of a substance to react or decay.

Collision Theory

States that reactions occur when molecules collide with sufficient energy and correct orientation.

Reaction Mechanisms

Reactions occurring in multiple steps, with intermediates, and the slowest step determines the rate law.

Catalysts

Substances that speed up reactions without being consumed, present in the beginning and end of elementary steps.

Enthalpy

Measure of energy released or absorbed during bond formation or breaking in a reaction.

Bond Energy

Energy required to break a bond, always a positive value.

Hess’s Law

States that the total enthalpy change of a reaction is the sum of enthalpy changes of individual steps.

Equilibrium Constant (Keq)

Expression showing the relationship between reactant and product concentrations at equilibrium.

Le Chatelier’s Principle

When a system is stressed, the reaction shifts to counteract the stress and reach a new equilibrium.

Le Chatelier's Principle

When the pressure of a container is increased, the reaction shifts to the side with fewer moles of gaseous particles, and vice versa when the pressure is decreased.

Temperature and Equilibrium

Changes in temperature affect the equilibrium position of a reaction; increasing temperature shifts the reaction towards the reactants, while decreasing temperature shifts it towards the products.

Reaction Quotient (Q)

Q can be used to determine how close a reaction is to equilibrium; if Q < K, the reaction shifts right, if Q > K, it shifts left, and if Q = K, the reaction is at equilibrium.

Solubility Product (Ksp)

Ksp measures the solubility of a salt in a solution; a higher Ksp value indicates higher solubility of the salt.

pH

pH is a measure of the concentration of hydrogen ions in a solution; a pH < 7 indicates an acidic solution, pH = 7 is neutral, and pH > 7 is basic.

Strong Acids

Strong acids completely dissociate in water and do not reach equilibrium; examples include HCl, HNO3, and H2SO4.

Weak Acids

Weak acids partially dissociate in water, and their dissociation constants are represented by Ka; a higher Ka value indicates a stronger acid.

Polyprotic Acids

Polyprotic acids can donate more than one hydrogen ion in a solution, with each successive dissociation having a different Ka value.

Equilibrium Constant of Water (Kw)

Kw represents the ionization of water into H+ and OH- ions; Kw = 1x10^-14 at 25 degrees Celsius.

Buffers

Buffers are solutions that resist changes in pH when an acid or base is added; they are made of a weak acid and its conjugate base, maintaining a stable pH.

Free Energy Change

The change in Gibbs free energy (∆G) during a chemical reaction, calculated as the difference between the sum of the free energy of products and the sum of the free energy of reactants.

Thermodynamically Favored Process (TFP)

If ∆G is negative, the process is thermodynamically favored, indicating spontaneity.

Equilibrium

When ∆G is 0, the system is at equilibrium, with no net change occurring.

Standard Free Energy Change

Calculated using the equation ∆G° = ∆H° - T∆S°, where ∆H° is the enthalpy change, ∆S° is the entropy change, and T is the temperature in Kelvin.

Equilibrium Constant (K)

The equilibrium constant can be used to calculate Gibbs free energy (∆G°) through the equation ∆G° = -RT(ln K), where R is the gas constant and T is the temperature in Kelvin.

Galvanic Cells

Cells that use favored redox reactions to generate electrical current, with oxidation occurring at the anode and reduction at the cathode.

Electrolytic Cells

Cells that use external voltage sources to drive non-spontaneous redox reactions, often in aqueous solutions, with the total cell potential always being negative.

Electroplating

The process of depositing a metal coating on a conductive object using electrolytic cells, where current (I) is related to charge (q) and time (t) by I = q/t.

Faraday’s Constant (F)

A constant value of 96,500 coulombs/mol used in calculations involving the number of moles of electrons exchanged in a redox reaction.

Standard Reaction Potential (E°)

The standard potential of a redox reaction, where a positive E° indicates a thermodynamically favored process and a negative E° indicates a non-favored process.