General Chemistry

General Chemistry Glossary

Every term used across all 42 General Chemistry tracks, defined and alphabetized. Search it, or jump by letter.

218 terms

A

Absolute zero
The lowest possible temperature (0 K, −273.15 °C), at which particle motion is minimized and gas volume theoretically approaches zero.
Accuracy
How close a measured value is to the true or accepted value (contrast with precision).
Acid (Arrhenius)
A substance that increases the concentration of H⁺ (H₃O⁺) ions when dissolved in water.
Acid (Brønsted–Lowry)
A proton (H⁺) donor.
Acid (Lewis)
An electron-pair acceptor.
Acid dissociation constant (Ka)
The equilibrium constant for the ionization of a weak acid; larger Ka means a stronger acid.
Activated complex (transition state)
The highest-energy, unstable arrangement of atoms at the peak of a reaction's energy diagram.
Activation energy (Ea)
The minimum energy colliding particles must possess for a reaction to occur.
Actual yield
The amount of product actually obtained from a reaction experimentally.
Alkali metals
Group 1 elements; highly reactive metals whose reactivity increases down the group.
Alkaline earth metals
Group 2 elements.
Alpha decay
Nuclear decay emitting an alpha particle (helium-4 nucleus); decreases atomic number by 2 and mass number by 4.
Amphoteric (amphiprotic)
A substance that can act as either an acid or a base (e.g., water).
Amphoteric oxide
An oxide that can react with both acids and bases.
Anode
The electrode where oxidation occurs.
Aqueous solution
A solution in which water is the solvent (abbreviated aq).
Arrhenius equation
Relates the rate constant to temperature and activation energy; explains why rate increases with temperature.
Atom
The smallest unit of an element that retains its chemical identity.
Atomic mass (atomic weight)
The weighted average mass of an element's isotopes, in atomic mass units (amu).
Atomic number (Z)
The number of protons in an atom's nucleus; defines the element.
Aufbau principle
Electrons fill atomic orbitals from lowest to highest energy.
Autoionization of water
The self-ionization of water into H⁺ and OH⁻, described by Kw.
Avogadro's law
At constant temperature and pressure, gas volume is directly proportional to the number of moles.
Avogadro's number
6.022 × 10²³, the number of particles in one mole.

B

Band theory
Model explaining electrical conductivity: metals conduct because valence and conduction bands overlap or are partially filled; insulators have a large band gap.
Base (Arrhenius)
A substance that increases OH⁻ concentration in water.
Base (Brønsted–Lowry)
A proton acceptor.
Base (Lewis)
An electron-pair donor.
Base dissociation constant (Kb)
The equilibrium constant for the ionization of a weak base.
Battery
A galvanic cell (or series of cells) that stores and delivers electrical energy via redox reactions.
Beta decay (β⁻)
Nuclear decay in which a neutron converts to a proton, emitting an electron; increases atomic number by 1.
Binding energy (nuclear)
The energy that holds a nucleus together; released when nucleons combine.
Boiling point
The temperature at which a liquid's vapor pressure equals the external pressure.
Bond energy
The energy required to break one mole of a particular bond in the gas phase.
Bond order
Half the difference between bonding and antibonding electrons (MO theory); indicates bond strength/stability.
Boyle's law
At constant temperature, gas pressure and volume are inversely proportional (P₁V₁ = P₂V₂).
Buffer
A solution of a weak acid and its conjugate base (or weak base and conjugate acid) that resists changes in pH.
Buffer capacity
The amount of acid or base a buffer can neutralize before pH changes significantly; greatest near the pKa.

C

Calorimetry
The measurement of heat flow in a chemical or physical process.
Catalyst
A substance that increases reaction rate by providing an alternative pathway with lower activation energy; not consumed and does not change ΔH or equilibrium position.
Cathode
The electrode where reduction occurs.
Cathodic protection
Preventing corrosion by connecting a metal to a more easily oxidized sacrificial anode.
Cell potential (Ecell)
The voltage produced by an electrochemical cell; positive for spontaneous (galvanic) reactions.
Charles's law
At constant pressure, gas volume is directly proportional to absolute temperature.
Chemical equilibrium
The state in which forward and reverse reaction rates are equal and concentrations remain constant.
Coefficient
The number placed before a formula in a balanced equation to indicate relative amounts.
Colligative properties
Properties that depend on the number of solute particles, not their identity (e.g., boiling point elevation, freezing point depression).
Collision theory
Reactions occur when particles collide with sufficient energy and correct orientation.
Combustion reaction
A reaction of a substance with oxygen, releasing heat and light.
Common ion effect
The shift in an equilibrium caused by adding an ion already present, decreasing solubility or ionization.
Conjugate acid
The species formed when a base gains a proton.
Conjugate base
The species formed when an acid loses a proton.
Coordination compound
A compound containing a central metal ion bonded to surrounding ligands.
Coordination number
The number of ligand donor atoms bonded to a central metal ion.
Corrosion
The redox degradation of a metal (e.g., rusting of iron).
Covalent bond
A bond formed by the sharing of electron pairs between atoms.
Covalent network solid
A solid held together by a continuous network of covalent bonds (e.g., diamond, quartz); very hard with high melting points.
Crystal field theory
Model explaining color and magnetism of transition metal complexes through d-orbital splitting by ligands.

D

Dalton's law of partial pressures
The total pressure of a gas mixture equals the sum of the partial pressures of each component.
Decomposition reaction
A reaction in which one compound breaks into two or more simpler substances.
Delocalized electrons
Electrons not confined to a single atom or bond (e.g., the electron sea in metals).
Diagonal relationship
Similarity in properties between an element and the element diagonally down and to the right (e.g., Li/Mg, Be/Al).
Diamagnetic
Having all electrons paired; weakly repelled by a magnetic field.
Diffusion
The spreading of particles from high to low concentration.
Dipole
A separation of partial positive and negative charge in a molecule.
Dipole–dipole force
An attraction between polar molecules.
Dissociation
The separation of a compound into ions in solution.
Dynamic equilibrium
A state where opposing processes occur at equal rates.

E

Effusion
The escape of gas molecules through a tiny opening.
Electrolysis
Using electrical energy to drive a nonspontaneous redox reaction.
Electrolyte
A substance that conducts electricity when dissolved or molten due to ion formation.
Electrolytic cell
An electrochemical cell that uses external electrical energy to drive a nonspontaneous reaction.
Electron
A negatively charged subatomic particle occupying orbitals around the nucleus.
Electron affinity
The energy change when an atom gains an electron.
Electron capture
Nuclear decay in which the nucleus captures an inner electron, converting a proton to a neutron.
Electron configuration
The arrangement of electrons among an atom's orbitals.
Electronegativity
An atom's ability to attract shared electrons in a bond; increases across a period and up a group.
Electroplating
Depositing a metal coating onto a surface using electrolysis.
Elementary step
A single molecular event in a reaction mechanism, whose rate law follows directly from its molecularity.
Empirical formula
The simplest whole-number ratio of atoms in a compound.
Endothermic
A process that absorbs heat (positive ΔH).
Endpoint
The point in a titration where the indicator changes color, chosen to approximate the equivalence point.
Enthalpy (H)
The heat content of a system at constant pressure; ΔH is the heat absorbed or released.
Enthalpy of formation (ΔH°f)
The enthalpy change when one mole of a compound forms from its elements in their standard states; zero for pure elements.
Entropy (S)
A measure of disorder or the number of accessible microstates; increases with dispersal of energy and matter.
Equilibrium constant (K)
The ratio of product to reactant concentrations (each raised to its coefficient) at equilibrium; Kc for concentrations, Kp for pressures.
Equivalence point
The point in a titration where moles of acid and base are stoichiometrically equal.
Exothermic
A process that releases heat (negative ΔH).

F

Faraday's constant
The charge of one mole of electrons, ≈ 96,500 C/mol.
Faraday's laws of electrolysis
The amount of substance produced at an electrode is proportional to the electric charge passed.
First law of thermodynamics
Energy is conserved; the internal energy change equals heat added plus work done on the system.
Fission (nuclear)
Splitting a heavy nucleus into lighter nuclei, releasing energy.
Formal charge
The charge assigned to an atom in a Lewis structure assuming equal sharing of bonding electrons.
Free energy (Gibbs, G)
ΔG = ΔH − TΔS; a negative value indicates a spontaneous process.
Fuel cell
A cell that continuously generates electricity from the redox reaction of a supplied fuel (e.g., H₂) with oxygen.
Functional group
A specific arrangement of atoms conferring characteristic reactivity (organic chemistry).
Fusion (nuclear)
Combining light nuclei into a heavier nucleus, releasing large amounts of energy (as in stars).

G

Galvanic (voltaic) cell
An electrochemical cell that converts chemical energy into electrical energy via a spontaneous redox reaction.
Gamma decay
Emission of high-energy photons (gamma rays) from a nucleus, with no change in atomic or mass number.
Gas constant (R)
The proportionality constant in the ideal gas law (0.0821 L·atm/mol·K).
Gay-Lussac's law
At constant volume, gas pressure is directly proportional to absolute temperature.
Groundstate
The lowest-energy electron configuration of an atom.

H

Half-equivalence point
The point in a weak-acid titration where half the acid is neutralized, so [HA] = [A⁻] and pH = pKa.
Half-life (t½)
The time required for half of a reactant (or radioactive isotope) to be consumed; constant for first-order and radioactive decay.
Half-reaction
The oxidation or reduction portion of a redox reaction written separately.
Halogens
Group 17 elements; highly reactive nonmetals whose reactivity decreases down the group.
Heat (q)
Energy transferred due to a temperature difference.
Heat capacity
The heat required to raise a substance's temperature by 1 °C (or 1 K).
Henderson–Hasselbalch equation
pH = pKa + log([A⁻]/[HA]); relates buffer pH to the acid/conjugate-base ratio.
Hess's law
The total enthalpy change of a reaction is the same regardless of the number of steps, because enthalpy is a state function.
Heterogeneous mixture
A mixture with non-uniform composition.
Homogeneous mixture
A mixture with uniform composition (a solution).
Hund's rule
Electrons occupy degenerate orbitals singly before pairing.
Hybridization
The mixing of atomic orbitals to form equivalent hybrid orbitals (e.g., sp³, sp², sp).
Hydrogen bond
A strong dipole attraction between H bonded to N, O, or F and a lone pair on another electronegative atom.
Hydronium ion (H₃O⁺)
The species formed when a proton associates with water.

I

Ideal gas
A hypothetical gas that perfectly obeys the gas laws (PV = nRT).
Ideal gas law
PV = nRT, relating pressure, volume, moles, and temperature.
Indicator
A substance that changes color to signal the endpoint of a titration.
Intermediate
A species formed in one step of a mechanism and consumed in a later step; does not appear in the overall equation.
Intermolecular forces
Attractions between molecules (dispersion, dipole–dipole, hydrogen bonding).
Ion
An atom or group of atoms with a net electrical charge.
Ionic bond
An electrostatic attraction between oppositely charged ions.
Ionic solid
A solid composed of a lattice of ions held by strong electrostatic forces; high melting point and brittle.
Ionization energy
The energy required to remove an electron from a gaseous atom; increases across a period and up a group.
Isotopes
Atoms of the same element with different numbers of neutrons (same Z, different mass number).

K

Kinetic molecular theory
Model describing gases as tiny particles in constant random motion with negligible volume and no intermolecular forces.
Kinetics
The study of reaction rates and the factors that affect them.
Kw (ion product of water)
[H⁺][OH⁻] = 1 × 10⁻¹⁴ at 25 °C.

L

Lattice energy
The energy released when gaseous ions combine to form an ionic solid.
Le Châtelier's principle
A system at equilibrium shifts to partially counteract an applied stress (concentration, pressure, temperature).
Lewis structure
A diagram showing bonding and lone-pair electrons using dots and lines.
Ligand
A Lewis base that donates an electron pair to a central metal ion in a complex.
Limiting reactant
The reactant that is completely consumed first, limiting the amount of product formed.
London dispersion forces
Weak temporary attractions arising from instantaneous dipoles; present in all molecules.

M

Mass number (A)
The total number of protons and neutrons in a nucleus.
Metallic solid
A solid of metal cations in a sea of delocalized electrons; conductive, malleable, lustrous.
Molality (m)
Moles of solute per kilogram of solvent.
Molar mass
The mass of one mole of a substance (g/mol).
Molarity (M)
Moles of solute per liter of solution.
Mole
The amount of substance containing 6.022 × 10²³ particles.
Molecular formula
The actual number of atoms of each element in a molecule.
Molecular orbital (MO) theory
Model in which atomic orbitals combine into bonding and antibonding molecular orbitals; correctly predicts O₂ paramagnetism.
Molecular solid
A solid of discrete molecules held by intermolecular forces; low melting points.
Molecularity
The number of species colliding in an elementary step.

N

Nernst equation
Relates actual cell potential to standard potential and the reaction quotient under nonstandard conditions.
Net ionic equation
An equation showing only the species that participate in the reaction, omitting spectator ions.
Neutralization
The reaction of an acid with a base to form water and a salt.
Neutron
An electrically neutral subatomic particle in the nucleus.
Noble gases
Group 18 elements; largely unreactive due to full valence shells.
Nuclear equation
An equation representing a nuclear reaction; balanced by conserving mass number and atomic number.

O

Octet rule
Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.
Orbital
A region of space where an electron is likely to be found.
Order of reaction
The exponent to which a reactant's concentration is raised in the rate law; determined experimentally.
Oxidation
The loss of electrons (increase in oxidation number).
Oxidation number (state)
A bookkeeping charge assigned to an atom to track electron transfer.
Oxidizing agent
The species that gains electrons and is itself reduced.

P

Paramagnetic
Having unpaired electrons; attracted to a magnetic field.
Partial pressure
The pressure a gas in a mixture would exert if it alone occupied the volume.
Pauli exclusion principle
No two electrons in an atom can have the same set of four quantum numbers; an orbital holds at most two electrons of opposite spin.
Percent yield
(Actual yield ÷ theoretical yield) × 100%.
pH
−log[H⁺]; a measure of acidity (below 7 acidic, 7 neutral, above 7 basic at 25 °C).
Phase change
A physical transition between solid, liquid, and gas states.
pKa
−log(Ka); a lower pKa indicates a stronger acid.
pOH
−log[OH⁻]; pH + pOH = 14 at 25 °C.
Polar covalent bond
A covalent bond with unequal electron sharing due to an electronegativity difference.
Polyatomic ion
A charged group of covalently bonded atoms (e.g., sulfate, nitrate).
Polyprotic acid
An acid that can donate more than one proton (e.g., H₂SO₄, H₃PO₄), ionizing in successive steps.
Positron emission
Nuclear decay in which a proton converts to a neutron, emitting a positron; decreases atomic number by 1.
Precipitate
An insoluble solid that forms from solution when Q exceeds Ksp.
Precision
How reproducible repeated measurements are (contrast with accuracy).
Pressure
Force per unit area exerted by gas particles.
Product
A substance formed in a chemical reaction.
Proton
A positively charged subatomic particle in the nucleus.

Q

Quantum numbers
Four numbers (n, l, mₗ, mₛ) describing an electron's energy, orbital shape, orientation, and spin.

R

Rate constant (k)
The proportionality constant in a rate law; its units depend on the overall reaction order.
Rate law
An equation relating reaction rate to reactant concentrations raised to their orders.
Rate-determining step
The slowest step in a mechanism, which limits the overall reaction rate.
Reactant
A starting substance consumed in a chemical reaction.
Reaction mechanism
The sequence of elementary steps by which a reaction occurs.
Reaction quotient (Q)
Has the same form as K but uses current concentrations; comparing Q to K predicts the direction of shift.
Reaction rate
The change in concentration of reactant or product per unit time.
Redox reaction
A reaction involving the transfer of electrons (simultaneous oxidation and reduction).
Reducing agent
The species that loses electrons and is itself oxidized.
Reduction
The gain of electrons (decrease in oxidation number).
Resonance
Delocalization of electrons represented by two or more valid Lewis structures.

S

Sacrificial anode
A more easily oxidized metal used to protect another metal from corrosion.
Salt
An ionic compound formed from the neutralization of an acid and a base.
Salt bridge
A connection allowing ion flow between the half-cells of a galvanic cell to maintain electrical neutrality.
Second law of thermodynamics
The total entropy of the universe increases in any spontaneous process.
Solubility product constant (Ksp)
The equilibrium constant for the dissolution of a sparingly soluble ionic solid.
Solute
The substance dissolved in a solution.
Solvent
The dissolving medium in a solution.
Specific heat
The heat required to raise 1 gram of a substance by 1 °C.
Spectator ion
An ion that appears unchanged on both sides of an ionic equation.
Spontaneous process
A process that occurs without continuous external input; ΔG < 0.
Standard cell potential (E°cell)
The cell potential under standard conditions; E°cathode − E°anode.
Standard reduction potential
The tendency of a species to be reduced under standard conditions; more positive means a stronger oxidizing agent.
State function
A property depending only on the current state, not the path taken (e.g., enthalpy, entropy, free energy).
Stoichiometry
The quantitative relationships between reactants and products in a reaction.
Strong acid/base
An acid or base that ionizes completely in water.
Sublimation
The direct transition from solid to gas.
Surroundings
Everything outside the system under study.
System
The specific part of the universe being studied.

T

Theoretical yield
The maximum product predicted from the limiting reactant.
Thermochemistry
The study of heat changes in chemical reactions.
Third law of thermodynamics
The entropy of a perfect crystal at absolute zero is zero.
Titration
A technique for determining concentration by reacting a solution with a standard solution of known concentration.
Transition metals
The d-block elements, often forming colored, paramagnetic complexes with variable oxidation states.
Transition state
See activated complex.
Triple point
The temperature and pressure at which solid, liquid, and gas coexist in equilibrium.

U

Unit cell
The smallest repeating unit of a crystal lattice.

V

Valence electrons
The outermost electrons involved in bonding.
van der Waals forces
Collective term for weak intermolecular attractions.
Vapor pressure
The pressure of a vapor in equilibrium with its liquid.
VSEPR theory
Valence Shell Electron Pair Repulsion theory; predicts molecular geometry by minimizing electron-pair repulsion.

W

Weak acid/base
An acid or base that only partially ionizes in water, establishing an equilibrium.
Work (w)
Energy transferred when a force moves an object; for gases, often pressure–volume work.