AP Chemistry Ultimate Guide (copy)

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.

Percent Composition

The percentage by mass of each element in a compound.

Empirical Formula

The simplest ratio of elements in a compound.

Electron Configuration

Describes the distribution of electrons in an atom's energy levels.

Coulomb's Law

Describes the electrostatic force between charges in an atom.

Ionic Charges

Result from elements gaining or losing electrons to achieve stability.

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 needed to remove a second electron from an atom.

Electronegativity Trend

Increases from left to right and decreases from top to bottom in the periodic table.

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 bonding between atoms using valence electrons.

Molecular Geometry

Arrangement of atoms in a molecule based on VSEPR theory.

Intermolecular Forces

Forces between molecules affecting their physical properties.

Boiling

Adding heat to break intermolecular forces (IMFs) for a substance to change from liquid to gas.

Vaporization

Transition from liquid to gas without additional heat.

Solution Separation

Process of separating substances based on different IMFs and Coulombic attractions.

Solutes and Solvents

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

Paper Chromatography

Separating mixtures by substance polarity using a medium like paper.

Retention Factor (Rf)

Measure of substance movement in chromatography.

Kinetic Molecular Theory

Describes behavior of ideal gases based on motion and collisions of gas molecules.

Effusion

Rate of gas escaping through microscopic holes based on speed, temperature, and molar mass.

Ideal Gas Equation

PV=nRT; relates pressure, volume, moles, gas constant, and temperature.

Dalton's Law

Total pressure of a gas mixture is the sum of partial pressures of individual gases.

Deviations From Ideal Behavior

Changes in gas behavior under extreme conditions or with strong IMFs.

Density

Mass of gas per unit volume; D=m/V.

Electromagnetic Spectrum

Energy change in electrons determined by frequency and Planck's constant.

Beer’s Law

Relates absorbance, molar absorptivity, pathlength, and concentration in spectrometry.

Types of Reactions

Synthesis, decomposition, acid-base, redox, hydrocarbon combustion, precipitation.

Stoichiometry

Balancing chemical equations and calculations based on moles and reactants.

Percent Error

Measure of accuracy in experimental values compared to expected values.

Oxidation States

Assigning charges to atoms in compounds based on rules.

Redox Reactions

Electron transfer between reactants leading to changes in oxidation states.

Acids and Bases

Substances donating or accepting protons in reactions.

Rate Law

Equation describing the rate of a reaction based on reactant concentrations.

Zero-Order Rate Laws

Rate independent of reactant concentration.

First-Order Rate Laws

Rate proportional to the concentration of one reactant.

Half-Life

The time taken for half of a substance to react, showing the percentage remaining after each half-life.

Collision Theory

States that reactions occur when molecules collide with sufficient energy, affected by concentration, surface area, and correct orientation.

Reaction Mechanisms

Reactions occurring in multiple steps, with intermediates, and the slowest step determining the rate law for the entire reaction.

Catalysts

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

Enthalpy

Measure of energy released or absorbed during bond formation or breaking, determining if a reaction is exothermic or endothermic.

Bond Energy

Energy required to break a bond, calculated by the difference between energy to break and form bonds in a reaction.

Hess’s Law

States that the sum of enthalpy changes of individual steps in a reaction equals the overall enthalpy change.

Equilibrium Constant (Keq)

Expression showing the relationship between reactant and product concentrations at equilibrium, with different types of K values for various reactions.

Le Chatelier’s Principle

When a system is stressed, the reaction shifts to counteract the stress, affecting equilibrium by changes in concentration, pressure, or temperature.

Le Chatelier's Principle

When the pressure of a container is altered, the equilibrium of a reaction will shift to counteract the change.

Equilibrium Constant

A value that remains constant when concentration or pressure changes, indicating the ratio of products to reactants at equilibrium.

Reaction Quotient (Q)

A value used to determine the direction a reaction will shift at any point during a reaction, compared to the equilibrium constant.

Solubility

The ability of a salt to dissolve in a solvent, often influenced by temperature changes.

Common Ion Effect

The impact of adding a common ion to a solution on the solubility of a slightly soluble salt.

pH

A measure of the acidity or basicity of a solution, determined by the concentration of hydrogen ions.

Acid Strengths

Categorized into strong acids that completely dissociate and weak acids that partially dissociate in water.

Polyprotic Acids

Acids that can donate more than one hydrogen ion in a solution.

Equilibrium Constant of Water (Kw)

The product of hydrogen ion and hydroxide ion concentrations in water at equilibrium.

Buffers

Solutions that resist changes in pH when small amounts of acid or base are added, often composed of a weak acid and its conjugate base.

Free Energy Change

The change in energy that occurs during a chemical reaction, calculated as the difference between the sum of free energy of products and the sum of free energy of reactants.

Thermodynamically Favored Process (TFP)

If the free energy change (∆G) is negative, the process is considered thermodynamically favored.

Equilibrium

When the free energy change (∆G) of a reaction is 0, the system is at equilibrium.

Gibbs Free Energy

The energy in a system available to do work, usually expressed in kJ/mol*K.

Standard Free Energy Change

The free energy change under standard conditions, calculated using the equation ∆G° = ∆H° - T∆S°.

Equilibrium Constant (K)

A measure of the extent of a reaction at equilibrium, related to the Gibbs free energy change through ∆G° = -RT(ln K).

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.

Reduction Potentials

The electric potentials associated with half-reactions, where positive values indicate favored redox reactions.

Faraday’s Constant (F)

A constant representing the charge of one mole of electrons, equal to 96,500 coulombs/mol.