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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.
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