Chemistry Glossary

134 core terms carry a worked explanation, an example, and cross-links; the rest give a one-line definition.

A

Accuracy

How close a measured value is to the true or accepted value. Accuracy is about correctness, not consistency — a set of measurements can be precise (tightly clustered) yet inaccurate if a systematic error pushes them all the same way.

e.g. Massing a 100.00 g standard four times and getting 97.8, 97.9, 97.8, 97.9 g is precise but not accurate; the balance needs calibrating.

Related: Precision Significant figures Systematic errorDeveloped in: Module 0 Module 1

Acid

In the Brønsted–Lowry model, a proton (H⁺) donor. In water an acid raises the hydrogen-ion concentration, lowering the pH below 7. Strong acids donate their protons essentially completely.

e.g. HCl is a strong acid: HCl(aq) → H⁺(aq) + Cl⁻(aq). Acetic acid is a weak acid and only partly ionises.

Acid dissociation constant

The equilibrium constant, Ka, for a weak acid's partial ionization in water (HA + H₂O ⇌ H₃O⁺ + A⁻). Unlike a strong acid, which is treated as 100% ionized, a weak acid's Ka fixes exactly how far that ionization actually goes — a larger Ka means a stronger, more fully ionized weak acid and, for the same starting concentration, a lower pH.

e.g. Acetic acid's Ka (about 1.8 × 10⁻⁵) is far larger than a very weak acid's but nowhere near a strong acid's effectively infinite ionization — which is exactly why vinegar is mildly acidic rather than dangerously corrosive.

Activation energy

The minimum energy colliding particles need for a reaction to actually proceed — the height of the energy barrier between reactants and products on a reaction energy diagram, regardless of whether the overall reaction releases or absorbs energy.

e.g. A reaction can be strongly exothermic (ΔH very negative) and still be slow at room temperature if its activation energy is high — striking a match releases plenty of energy, but the match still needs the initial heat of friction to get over that barrier.

Related: Catalyst Enthalpy Exothermic EndothermicDeveloped in: Module 9

Actual yield

Product amount measured experimentally.

Addition reaction

An organic reaction in which atoms add across a carbon-carbon double or triple bond, converting it to a single bond — the defining reactivity of alkenes and alkynes, since that extra bond is exactly the site an addition reaction opens up.

e.g. Hydrogenating vegetable oil — adding H₂ across the double bonds in unsaturated fat molecules to make them solid at room temperature (margarine) — is an addition reaction at industrial scale.

Alcohol

An organic compound containing a hydroxyl (-OH) group bonded to a carbon chain; the functional group behind everything from rubbing alcohol to the ethanol in fermented beverages.

Aldehyde

An organic compound containing a carbonyl group (C=O) bonded to at least one hydrogen atom, typically at the end of a carbon chain.

Alkane

A hydrocarbon containing only single carbon–carbon bonds — the simplest and least reactive hydrocarbon class, since a single bond has no extra electron density exposed the way a double or triple bond does.

e.g. Methane (CH₄), the main component of natural gas, and octane (C₈H₁₈), a major component of gasoline, are both alkanes despite their very different uses.

Related: Hydrocarbon Alkene Alkyne

Alkene

A hydrocarbon containing at least one carbon–carbon double bond. That extra electron density makes an alkene meaningfully more reactive than an alkane — it's a site where addition reactions can occur, breaking the double bond to attach new atoms.

e.g. Ethylene (the simplest alkene) is polymerized into polyethylene, the most common plastic, by repeatedly opening its double bond to link thousands of monomers into a chain.

Alkyne

A hydrocarbon containing at least one carbon–carbon triple bond — the most reactive of the three hydrocarbon classes, for the same reason an alkene beats an alkane: more exposed electron density, more available reaction sites.

e.g. Acetylene (the simplest alkyne) burns hot enough for oxy-acetylene welding torches precisely because its triple bond stores more energy than the double or single bonds in an equivalent alkene or alkane.

Related: Alkene Alkane Hydrocarbon

Alpha decay

Radioactive decay in which an unstable nucleus emits an alpha particle — two protons and two neutrons bound together, identical to a helium-4 nucleus — lowering the parent's mass number by 4 and atomic number by 2.

e.g. Uranium-238 alpha-decays to thorium-234: ²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂He.

Amine

An organic compound built around a nitrogen atom bonded to one or more carbon groups, conceptually ammonia (NH₃) with hydrogens swapped for carbon chains. Amines are typically weak bases — the nitrogen's lone electron pair, the same one that makes ammonia a base, is still available to accept a proton.

e.g. The amine group is what makes an amino acid's other end basic, balancing the carboxylic acid group's acidity on the same molecule.

Amino acid

An organic molecule containing both an amine functional group and a carboxylic acid functional group attached to the same carbon — the monomer that links, one after another, into the polymer chains that fold into proteins. The same functional-group logic that organizes all of organic chemistry, applied to life's own building blocks.

e.g. The twenty amino acids that build human proteins differ only in one side-chain group each — everything from a protein's shape to an enzyme's function traces back to how those twenty pieces are ordered and folded.

Amphoteric

Able to act as either a Brønsted–Lowry acid or base depending on the other species present in the reaction. Water is the most familiar example — it's why a single water molecule can either donate or accept a proton, the basis of its own self-ionization.

e.g. Water acts as a base when reacting with HCl (accepting a proton to form H₃O⁺) and as an acid when reacting with NH₃ (donating a proton to form OH⁻).

Anion

A negatively charged ion, formed when an atom or molecule gains one or more electrons. Non-metals typically form anions, and many important anions are polyatomic — a covalently bonded group carrying an overall charge.

e.g. Cl⁻, O²⁻, and the polyatomic SO₄²⁻ are anions.

Related: Cation Ion Ionic bond Polyatomic ionDeveloped in: Module 2 Module 12

Anode

The electrode where oxidation occurs in any electrochemical cell. Its electrical charge depends on the cell type — negative in a galvanic (spontaneous, power-generating) cell, positive in an electrolytic (externally driven) cell — which is a common point of confusion since the electrode's role (oxidation) never changes even though its sign does.

e.g. In a standard zinc–copper galvanic cell, zinc metal is the anode: it oxidizes (Zn → Zn²⁺ + 2e⁻), releasing electrons into the external circuit.

Aqueous

Describing a solution in which water is the solvent, written with the state symbol (aq). Many acid–base and precipitation reactions are studied in aqueous solution because water supports the mobile ions they need.

e.g. NaCl(aq) means sodium and chloride ions dispersed in water, not solid salt.

Related: Solvent Solute Electrolyte PrecipitateDeveloped in: Module 6

Arrhenius equation

An equation, k = Ae^(−Ea/RT), relating a reaction's rate constant to its activation energy and the absolute temperature. It quantifies why reaction rates are so sensitive to temperature — a relatively small temperature increase can multiply the rate several times over, because temperature sits in the exponent.

e.g. The rule of thumb that 'reaction rate roughly doubles for every 10°C rise' is a rough real-world consequence of the Arrhenius equation's exponential temperature dependence.

Atom

The smallest particle of an element that keeps that element's chemical identity. An atom has a dense central nucleus of protons and neutrons surrounded by electrons; the number of protons (the atomic number) defines which element it is.

e.g. Every atom with 6 protons is carbon, whether it has 6, 7, or 8 neutrons.

Related: Atomic number Isotope Element Electron ProtonDeveloped in: Module 2 Module 10

Atomic mass

The mass of an atom, or the abundance-weighted average mass of an element's naturally occurring isotopes, expressed on the unified atomic mass scale (u). The periodic table value is this weighted average.

e.g. Chlorine's atomic mass of 35.45 u reflects roughly 76% ³⁵Cl (34.97 u) and 24% ³⁷Cl (36.97 u).

Related: Isotope Molar mass Atomic number MoleDeveloped in: Module 2 Module 3

Atomic number

The number of protons in an atom's nucleus, written as a subscript Z. It fixes the element's identity and its position in the periodic table, and in a neutral atom it also equals the number of electrons.

e.g. Z = 26 is iron in every case; changing Z changes the element, while changing the neutron count only changes the isotope.

Related: Atom Isotope Proton Periodic table ElementDeveloped in: Module 2 Module 10 Module 11

Aufbau principle

Ground-state filling pattern in which available lower-energy orbitals are occupied before higher-energy orbitals, subject to known exceptions.

Average atomic mass

Abundance-weighted mean isotope mass for a specified natural composition.

Avogadro constant

The number of elementary entities in one mole, 6.02214076 × 10²³ mol⁻¹, now fixed by definition. Multiplying an amount in moles by this constant gives a count of particles.

e.g. 0.25 mol of neon contains 0.25 × 6.022 × 10²³ = 1.51 × 10²³ atoms.

Related: Mole Molar massDeveloped in: Module 3

B

Base

In the Brønsted–Lowry model, a proton (H⁺) acceptor. In water a base raises the hydroxide-ion concentration, raising the pH above 7. Strong bases are essentially fully dissociated metal hydroxides.

e.g. NaOH is a strong base; ammonia is a weak base that accepts a proton from water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻.

Related: Acid pH Neutralization ElectrolyteDeveloped in: Module 6 Module 7

Base dissociation constant

The equilibrium constant, Kb, for a weak base's partial reaction with water (B + H₂O ⇌ BH⁺ + OH⁻), the direct base-side counterpart to a weak acid's Ka. A weak acid's Ka and its conjugate base's Kb are linked through water's own equilibrium constant, Kw = Ka × Kb — so the stronger a conjugate base actually is, the weaker its parent acid must have been.

e.g. Ammonia's modest Kb makes it a weak base, consistent with household ammonia solution being mildly basic rather than caustic the way a strong base like sodium hydroxide is.

Beta decay

Radioactive decay in which a neutron converts to a proton while emitting a fast electron (β⁻ decay, the common case), or a proton converts to a neutron while emitting a positron (β⁺ decay) — either way the mass number is unchanged while the atomic number shifts by one, transmuting the element.

e.g. Carbon-14 beta-decays to nitrogen-14 (¹⁴₆C → ¹⁴₇N + e⁻), the basis of radiocarbon dating.

Binding energy

The energy that would be needed to completely separate a nucleus into individual protons and neutrons — equivalently, the energy released when those particles first bind together into that nucleus. This is why a nucleus's mass is always slightly less than the sum of its separated protons and neutrons: that missing mass was converted to binding energy (via E = mc²) when the nucleus formed.

e.g. Iron-56 has the highest binding energy per nucleon of any isotope, which is why both fission (of heavier elements) and fusion (of lighter elements) release energy — each moves toward iron's more stable arrangement.

Related: Nuclear fission Nuclear fusion

Bond angle

The angle between two bonds that meet at the same atom, set by the electron-domain geometry.

Bond dipole

Vector representation of unequal electron distribution in a bond.

Boyle's law

At constant temperature and amount, gas pressure and volume are inversely related: P₁V₁ = P₂V₂.

Bronsted-Lowry theory

The definition of an acid as a proton donor and a base as a proton acceptor — the model this course's own Acid and Base definitions already use for strong-acid/strong-base neutralization. The broader Brønsted–Lowry picture goes further, into weak acids and bases that only partially donate or accept protons, and the equilibria that partial reaction sets up.

e.g. Acetic acid (vinegar) only partially donates its proton in water, establishing an equilibrium between acetic acid and its conjugate base (acetate) rather than reacting to completion the way a strong acid does.

Buffer

A solution that resists pH change when a small amount of acid or base is added, typically a weak acid paired with its conjugate base; the equilibrium calculations behind buffers go beyond this course's own module sequence.

Buffer capacity

The amount of added acid or base a buffer solution can absorb before its pH starts changing significantly — determined by how much of the weak acid and its conjugate base are actually present, not by the pH itself. A buffer can hold a stable pH yet still have a small buffer capacity if there simply isn't much of the acid-base pair to work with.

e.g. A dilute buffer and a concentrated buffer can have the exact same starting pH, but the concentrated one absorbs a far larger acid or base spike before that pH actually shifts.

C

Calorimetry

The measurement of heat flow in a physical or chemical change, using the temperature change of a known mass and its specific heat capacity: q = mcΔT.

e.g. If 100.0 g of water rises 4.2 °C, it absorbed q = (100.0)(4.18)(4.2) ≈ 1.8 kJ.

Related: Enthalpy Exothermic Endothermic EnergyDeveloped in: Module 8

Carbohydrate

A biological molecule built from carbon, hydrogen, and oxygen in roughly a 1:2:1 ratio, ranging from simple sugars to large polymers like starch and cellulose, used primarily for energy storage and structure.

Carboxylic acid

An organic compound built around a carboxyl group (-COOH) — a carbon double-bonded to one oxygen and single-bonded to a hydroxyl group. The O-H bond is weakened by the neighboring carbonyl, which is why a carboxylic acid can donate that hydrogen as a proton and behave as a weak Brønsted–Lowry acid.

e.g. Acetic acid (the acid in vinegar) and citric acid (in citrus fruit) are both carboxylic acids — their mild sourness is exactly this functional group donating protons.

Catalyst

A substance that speeds up a reaction by opening a lower-activation-energy pathway, without itself being consumed — it appears in the mechanism but is regenerated by the end of the reaction, so it changes how fast a reaction reaches equilibrium, never the reactants' or products' actual energy values.

e.g. On an energy diagram, a catalyst lowers the hump between reactants and products but leaves both endpoints — and therefore ΔH — exactly where they were.

Related: Activation energy EnthalpyDeveloped in: Module 9

Cathode

The electrode where reduction occurs in any electrochemical cell — positive in a galvanic cell, negative in an electrolytic cell. Pairing this with anode's oxidation is what 'redox' (reduction–oxidation) actually names.

e.g. In the zinc–copper cell, copper is the cathode: Cu²⁺ ions from solution gain electrons and deposit as solid copper (Cu²⁺ + 2e⁻ → Cu).

Cation

A positively charged ion, formed when an atom or molecule loses one or more electrons. Metals typically form cations, and in electrolysis cations migrate toward the negative electrode (cathode).

e.g. Ca²⁺, NH₄⁺, and Al³⁺ are cations; Ca²⁺ has lost the two electrons a neutral calcium atom holds in its outer shell.

Related: Anion Ion Ionic bondDeveloped in: Module 2 Module 12

Cell potential

The voltage a galvanic cell produces, calculated as the cathode's standard reduction potential minus the anode's. A positive cell potential confirms the reaction runs spontaneously as written; a negative value means the reverse reaction is the spontaneous one.

e.g. Cu²⁺/Cu (+0.34 V) as cathode minus Zn²⁺/Zn (−0.76 V) as anode gives a cell potential of +1.10 V — matching the actual voltage a zinc–copper battery produces.

Charles's law

At constant pressure and amount, gas volume and absolute temperature are directly related: V₁/T₁ = V₂/T₂.

Chemical change

Process that changes chemical identities.

Chemical equation

A symbolic representation of a reaction, with reactant formulas on the left, product formulas on the right, and coefficients chosen so that every element and the total charge are conserved.

e.g. CH₄ + 2 O₂ → CO₂ + 2 H₂O shows one methane molecule reacting with two oxygen molecules.

Chemical equilibrium

The dynamic state a reversible reaction settles into when its forward and reverse rates become equal, so concentrations stop changing even though both reactions are still actively happening. It is a kinetic balance, not a stopped reaction — disturb it (Le Chatelier's principle) and it shifts to a new balance.

e.g. In a sealed bottle of soda, CO₂ escaping the liquid and CO₂ re-dissolving happen at equal rates once the bottle has sat closed long enough — the fizz level looks static, but molecules are crossing that boundary in both directions constantly.

Chemical formula

A notation using element symbols and subscripts to show the kind and number of atoms in a substance. A molecular formula gives the actual atom count in a molecule; an empirical formula gives only the smallest whole-number ratio.

e.g. Hydrogen peroxide's molecular formula is H₂O₂; its empirical formula is HO. Glucose is C₆H₁₂O₆ (empirical CH₂O).

Chirality

The property of a molecule that cannot be superimposed on its own mirror image, the same relationship a left hand has to a right hand — hence 'chiral,' from the Greek word for hand. Chirality matters enormously in biology, where enzymes are themselves chiral and usually only recognize one mirror-image form of a molecule.

e.g. One mirror-image form of the drug thalidomide was an effective sedative; the other caused severe birth defects — a stark real-world case of why chirality isn't a minor technicality.

Related: Stereoisomer Isomer

Coefficient

The whole number placed in front of a formula in a balanced equation, giving the relative number of moles (or molecules) of that species. Coefficients — not subscripts — are adjusted to balance an equation.

e.g. In 2 H₂ + O₂ → 2 H₂O the coefficients 2, 1, 2 give the mole ratio; changing a subscript would change the substance.

Related: Chemical equation Stoichiometry MoleDeveloped in: Module 4 Module 5

Collision theory

The model explaining reaction rate as a function of how often particles collide, whether they collide with enough energy to react (activation energy), and whether they collide with the correct orientation.

Combustion

A rapid reaction of a fuel with oxygen that releases heat and light; complete combustion of a hydrocarbon produces carbon dioxide and water.

Compound

A pure substance made of two or more elements chemically bonded in a fixed whole-number ratio. A compound has properties different from its constituent elements and can be separated into them only by chemical reaction.

e.g. Water is always 2 hydrogen atoms to 1 oxygen atom (11.2% H by mass); it is nothing like the gases it forms from.

Related: Element Chemical formula Mixture Chemical changeDeveloped in: Module 1 Module 2

Concentration

A measure of how much solute is present in a given amount of solution or solvent. Common expressions are molarity, mass percent, parts per million, and mole fraction.

e.g. Seawater is about 0.6 M in NaCl, or roughly 3.5% salt by mass.

Related: Molarity Solute Solvent DilutionDeveloped in: Module 6

Conjugate acid–base pair

Two species that differ by one proton, such as NH₄⁺ and NH₃ or HCO₃⁻ and CO₃²⁻.

Conversion factor

A ratio equal to 1, built from two equivalent quantities expressed in different units. Multiplying by a conversion factor changes the units a quantity is expressed in without changing the quantity itself — the mechanism dimensional analysis runs on.

e.g. Because 1 mol contains exactly 6.02214076 × 10²³ particles, (6.022 × 10²³ particles / 1 mol) and its reciprocal (1 mol / 6.022 × 10²³ particles) are both valid conversion factors.

Related: Dimensional analysis MoleDeveloped in: Module 0

Covalent bond

A bond formed by a shared pair of electrons between two atoms, typical of non-metal–non-metal combinations. Sharing can be equal (non-polar) or unequal (polar), depending on the electronegativity difference.

e.g. The bond in Cl₂ is non-polar covalent; each bond in HCl is polar covalent.

Critical point

The temperature and pressure above which the liquid and gas phases of a substance stop being distinguishable — no amount of added pressure will condense a substance held above its critical temperature into an ordinary liquid.

e.g. Above its critical point (31.1 °C, 72.9 atm), CO₂ becomes a supercritical fluid, used to strip caffeine from coffee beans without the toxic solvents older methods relied on.

Related: Phase diagram Vapor pressure Ideal gasDeveloped in: Module 14

D

Dalton's law

The total pressure of a gas mixture equals the sum of each component's partial pressure.

Decay constant

The proportionality constant, λ, relating how fast a radioactive sample decays to how many undecayed nuclei it currently has — a fixed property of a given isotope, unaffected by temperature, pressure, or chemical environment, unlike almost every rate constant in ordinary chemical kinetics. It sets the sample's half-life through t½ = ln2 / λ, the exact same relationship a chemical rate constant has to its own first-order half-life.

e.g. Two samples of the same isotope, one twice the mass of the other, decay at proportionally different rates but share the identical decay constant and the identical half-life — only the total decay events per second scale with sample size, not the underlying constant.

Degree of unsaturation

A count, calculated directly from a molecular formula without drawing any structure, of how many rings and pi bonds (double or triple bonds) an organic molecule must contain. Each ring or pi bond costs the molecule exactly two hydrogens relative to a fully saturated, ring-free hydrocarbon with the same carbon count — so working backward from a formula's actual hydrogen count against that saturated maximum reveals how much unsaturation has to be hiding somewhere in the structure.

e.g. Benzene, C6H6, has a degree of unsaturation of 4 — one ring plus three double bonds in its usual Lewis structure — exactly matching the four sites a fully saturated six-carbon ring-free hydrocarbon (C6H14) would need to lose to reach C6H6's lower hydrogen count.

Denaturation

The loss of a protein's three-dimensional folded structure — and with it, its function — caused by heat, extreme pH, or other disruptions, without breaking the peptide bonds that link its amino acids together. The amino acid sequence survives; only the fold that made it functional is destroyed.

e.g. Cooking an egg white denatures its proteins irreversibly: heat unfolds them, they tangle together in their new shape, and no amount of cooling refolds them back into raw egg white.

Related: Protein Enzyme Peptide bond

Density

Mass per volume under specified conditions.

Dilution

Lowering the concentration of a solution by adding solvent. Because the amount of solute does not change, C₁V₁ = C₂V₂ relates the concentrations and volumes before and after.

e.g. To make 100 mL of 1.5 M HCl from 12 M stock: V₁ = C₂V₂ / C₁ = (1.5)(100) / 12 = 12.5 mL of stock, diluted to 100 mL.

Related: Molarity Concentration Solvent SoluteDeveloped in: Module 6

Dimensional analysis

Solving a numeric problem by chaining conversion factors so that unwanted units cancel algebraically, leaving only the units the answer requires. It is less a single technique than the organizing method behind nearly every calculation in this course — stoichiometry, gas laws, and dilutions are all dimensional analysis wearing a different topic's units.

e.g. Converting 3.5 days to seconds: 3.5 d × (24 h / 1 d) × (60 min / 1 h) × (60 s / 1 min) = 302,400 s — each factor is a ratio equal to 1, and the day/hour/minute units cancel in sequence, leaving only seconds.

Related: Conversion factor Significant figuresDeveloped in: Module 0

Dipole–dipole attraction

Intermolecular attraction among permanent molecular dipoles.

Dispersion force

Attraction arising from correlated instantaneous electron fluctuations; present in all atoms and molecules.

Dissociation

The separation of a compound into ions as it dissolves or melts.

E

Effective nuclear charge

The net positive pull an electron actually feels from the nucleus, after inner electrons partially cancel out (shield) the full nuclear charge. It is what actually drives periodic trends — not the raw number of protons, which only ever increases — because effective nuclear charge rises across a period (shielding stays roughly constant while protons increase) but stays comparatively flat down a group.

e.g. Sodium (Z = 11) and chlorine (Z = 17) both have core electrons shielding their valence shell, but chlorine's seven valence electrons feel a much higher effective nuclear charge, which is why chlorine is smaller and far more electronegative.

Electrode

A conductor where a half-reaction occurs in an electrochemical cell — the anode (oxidation) or the cathode (reduction).

Electrolysis

Using an external electrical current to force a non-spontaneous redox reaction to occur, the reverse relationship to a galvanic cell.

Electrolyte

A substance that produces mobile ions when dissolved or melted, allowing the solution to conduct electricity. Strong electrolytes ionise essentially completely; weak electrolytes only partly.

e.g. NaCl and HCl are strong electrolytes; acetic acid is a weak electrolyte; sugar is a non-electrolyte.

Related: Ion Aqueous Acid Base DissociationDeveloped in: Module 6

Electrolytic cell

An electrochemical cell that uses an external power source to force a non-spontaneous redox reaction to occur — the mirror image of a galvanic cell, which generates electricity from a spontaneous one. Electrolysis is the process; an electrolytic cell is the device that carries it out.

e.g. Electroplating a metal object with a thin layer of gold uses an electrolytic cell: left alone, the reaction wouldn't happen, so an external current drives it.

Electron

Negatively charged subatomic particle.

Electron affinity

Energy change associated with adding an electron to a gaseous atom or ion under a defined convention.

Electron configuration

The distribution of an atom's or ion's electrons among orbitals, written as a string like 1s²2s²2p⁶ — built up using the Aufbau principle (fill lowest energy first), the Pauli exclusion principle (at most two electrons per orbital, opposite spins), and Hund's rule (spread out across degenerate orbitals before pairing). The pattern it produces is what drives the periodic table's own structure.

e.g. Oxygen (Z = 8) has the ground-state configuration 1s²2s²2p⁴ — the 2p⁴ tail is why it commonly forms a 2− ion, gaining two electrons to reach neon's 2p⁶.

Electronegativity

A relative measure of how strongly a bonded atom attracts shared electrons. It rises across a period and falls down a group; large differences give ionic bonds, small differences give non-polar covalent bonds.

e.g. The O–H bonds in water are polar because oxygen (3.44) pulls electron density away from hydrogen (2.20).

Element

A pure substance made of only one kind of atom, defined by a fixed number of protons. The 118 known elements are the entries of the periodic table and cannot be broken into simpler substances by chemical means.

e.g. Oxygen (8 protons) and gold (79 protons) are elements; water and brass are not.

Related: Atom Atomic number Compound Periodic tableDeveloped in: Module 1 Module 2 Module 11

Empirical formula

The formula that gives the simplest whole-number ratio of atoms in a compound. It is found from mass or percent composition data and becomes the molecular formula once multiplied by the ratio of measured molar mass to empirical formula mass.

e.g. A compound that is 40.0% C, 6.7% H, 53.3% O by mass has the empirical formula CH₂O; if its molar mass is 180 g/mol, the molecular formula is C₆H₁₂O₆.

Related: Molecular formula Chemical formula Molar mass MoleDeveloped in: Module 3

Endothermic

Describing a process that absorbs heat from its surroundings, so the surroundings cool and the system's enthalpy increases (ΔH > 0).

e.g. Dissolving ammonium nitrate in water is endothermic — the basis of instant cold packs.

Related: Exothermic Enthalpy CalorimetryDeveloped in: Module 8

Endpoint

Observable titration signal used to estimate the equivalence point.

Energy

Capacity to transfer heat or do work.

Enthalpy

A thermodynamic quantity, H = E + PV, whose change at constant pressure equals the heat absorbed or released by a reaction. A negative ΔH is exothermic; a positive ΔH is endothermic.

e.g. Combustion of methane has ΔH°rxn ≈ −890 kJ/mol — strongly exothermic.

Related: Exothermic Endothermic Calorimetry Hess's law EnergyDeveloped in: Module 8 Module 9

Entropy

A measure of how many equivalent ways a system's particles and energy can be arranged — informally, a measure of disorder or dispersal. The second law of thermodynamics states that the total entropy of an isolated system never decreases, which is the deeper reason heat flows from hot to cold and not the other way around.

e.g. An ice cube melting increases entropy: a rigid crystal lattice (few possible arrangements) becomes a liquid (vastly more possible arrangements of the same molecules).

Enzyme

A protein that acts as a biological catalyst — lowering a specific reaction's activation energy, without being consumed, by binding its target molecule (the substrate) at a precisely shaped active site. Because the active site's shape comes from the protein's fold, denaturing an enzyme destroys its catalytic function even though its chemical composition is unchanged.

e.g. The enzyme amylase in saliva breaks down starch into simpler sugars — which is why chewing a plain cracker for a while starts to taste sweet.

Equilibrium constant

The ratio K = [products]/[reactants] (each concentration raised to its balanced-equation coefficient) at equilibrium — a single number that summarizes how far a reversible reaction proceeds before forward and reverse rates equalize. K is fixed for a given reaction at a given temperature, regardless of the starting amounts.

e.g. A reaction with K = 10⁻⁵ barely proceeds — at equilibrium it is overwhelmingly reactants; K = 10⁵ means the reverse is true.

Equivalence point

Titration condition at which reactants have combined in the stoichiometric ratio.

Ester

An organic compound formed when a carboxylic acid's -OH is replaced by an -OR group from an alcohol, in a reaction (esterification) that releases water. The carboxylic acid's sharp acidity disappears once it becomes an ester, replaced by what's often a pleasant smell.

e.g. Ethyl acetate, an ester formed from acetic acid and ethanol, is the fruity-smelling solvent in many nail-polish removers — a small structural change from a sour-smelling acid to a fragrant ester.

Esterification

The reaction between a carboxylic acid and an alcohol that produces an ester and water, the reaction that builds an ester functional group.

Excess reactant

In a reaction with more than one reactant, whichever one is not fully consumed when the limiting reactant runs out; some of it remains unreacted at the end.

Exothermic

Describing a process that releases heat to its surroundings, so the surroundings warm up and the system's enthalpy decreases (ΔH < 0).

e.g. Dissolving calcium chloride in water is exothermic — the container gets noticeably warm.

Related: Endothermic Enthalpy CalorimetryDeveloped in: Module 8

F

First law of thermodynamics

Energy is conserved: it is never created or destroyed, only transferred as heat or work, or converted between forms. Formally, ΔE = q + w — a system's change in internal energy equals the heat added to it (q) plus the work done on it (w). Every calorimetry and enthalpy calculation in this course is an application of this one law to a specific, simplified case.

e.g. In a sealed, rigid, insulated container, a reaction can convert chemical potential energy entirely into heat with no work done (w = 0) — so ΔE = q exactly, which is the assumption behind a bomb calorimeter.

Related: Internal energy Energy Calorimetry EnthalpyDeveloped in: Module 8

Formal charge

Electron-bookkeeping value assigned from a Lewis structure.

Formula mass

The sum of atomic masses of every atom in a chemical formula, in atomic mass units; called molecular mass for a molecular compound and used interchangeably with molar mass when expressed in grams per mole.

Formula unit

Lowest formula-ratio entity used for an ionic compound.

Frequency

Number of wave cycles passing a point per unit time.

Functional group

A specific, recognizable arrangement of atoms within a larger organic molecule — a hydroxyl (–OH), carbonyl (C=O), carboxyl (–COOH), amine (–NH₂), and dozens of others — that reacts in a predictable, characteristic way essentially independent of what the rest of the molecule looks like. Organic chemistry is organized almost entirely around functional groups rather than around individual compounds.

e.g. Ethanol (an alcohol, –OH) and acetic acid (a carboxylic acid, –COOH) both start from a two-carbon backbone, but their functional groups make them react in completely different, predictable families of reactions.

Related: Hydrocarbon Isomer Polymer

G

Galvanic cell

An electrochemical cell that converts the energy of a spontaneous redox reaction directly into electrical current — oxidation happens at the anode, reduction at the cathode, and electrons flow through an external wire between them (ions flow through a salt bridge to keep both half-cells electrically neutral).

e.g. A standard alkaline battery is a galvanic cell; once its spontaneous redox reaction reaches equilibrium, it stops producing usable voltage — the battery is 'dead.'

Gamma radiation

High-energy electromagnetic radiation emitted from an unstable nucleus, usually accompanying alpha or beta decay, with no change to mass number or atomic number.

Gas

State of matter that expands to fill a container and is readily compressible.

Gibbs free energy

A single quantity, ΔG = ΔH − TΔS, that combines a reaction's enthalpy change and entropy change to determine whether it's spontaneous at a given temperature. Negative ΔG means spontaneous as written; positive means the reverse direction is spontaneous instead; zero means the system is at equilibrium.

e.g. Ice melting above 0°C has both a positive ΔH (absorbs heat) and a positive ΔS (more disorder) — above 0°C, the TΔS term wins out, making ΔG negative and melting spontaneous.

Ground state

The lowest-energy, most stable electron arrangement available to an atom or ion under the Aufbau ordering. Any arrangement with an electron in a higher-energy orbital than the ground state requires — one it wasn't required to occupy — is an excited state, reached only by absorbing energy.

e.g. Hydrogen's ground state has its one electron in the 1s orbital; absorbing a photon can excite it to 2s or 2p, and it emits a photon of the same energy falling back.

Related: Electron configuration Aufbau principle PhotonDeveloped in: Module 10

H

Half-life

The time required for half of a reactant (or a radioactive sample) to be consumed or decay; constant for a first-order or radioactive process regardless of starting amount.

Half-reaction

One of the two halves of a redox reaction, written to show only the oxidation or only the reduction occurring, each balanced separately for atoms and charge before being combined (with electrons cancelling) into the overall equation.

e.g. Zn → Zn²⁺ + 2e⁻ (oxidation) and Cu²⁺ + 2e⁻ → Cu (reduction) combine to give the overall reaction Zn + Cu²⁺ → Zn²⁺ + Cu.

Heat

Energy transferred because of a temperature difference.

Heat capacity

Energy required per temperature change for an object or system.

Heat of fusion

The energy required to melt one mole of a solid into a liquid at its melting point, with the temperature held constant throughout — all of that energy goes into overcoming enough intermolecular attraction to let particles move past each other, none of it into faster particle motion (which is why the thermometer doesn't move during melting).

e.g. Melting an ice cube takes real, measurable time at exactly 0 °C because the added heat is being spent on water's 6.01 kJ/mol heat of fusion, not on raising the temperature.

Heat of vaporization

The energy required to boil one mole of a liquid into a gas at its boiling point, with the temperature held constant. Always much larger than the same substance's heat of fusion — melting only loosens a solid's rigid structure enough to flow, while vaporization has to fully separate molecules against every remaining intermolecular attraction.

e.g. Water's heat of vaporization (40.7 kJ/mol) is nearly seven times its heat of fusion (6.01 kJ/mol) — which is also why a steam burn is far worse than a hot-water burn of the same temperature: condensing steam back to liquid on skin releases that entire 40.7 kJ/mol on contact.

Related: Heat of fusion Vapor pressure Intermolecular forcesDeveloped in: Module 14

Henderson–Hasselbalch equation

A rearrangement of a weak acid's equilibrium expression into pH = pKa + log([A⁻]/[HA]), giving a buffer's pH directly from the acid's pKa (= −log Ka) and the ratio of conjugate base to weak acid actually present — without solving the full equilibrium from scratch each time. When the acid and conjugate base are present in equal amounts, that ratio is 1, its logarithm is 0, and pH simply equals pKa.

e.g. A buffer built from equal moles of acetic acid and acetate sits at a pH equal to acetic acid's own pKa (about 4.74); doubling the acetate relative to the acid shifts the pH up by log(2) ≈ 0.3 units from there.

Hess's law

Enthalpy changes add when chemical equations add.

Heterogeneous mixture

Mixture with nonuniform regions at the scale considered.

Homogeneous mixture

Mixture uniform at the scale considered.

Hund's rule

Degenerate orbitals are singly occupied with parallel spins before electrons pair in the ground-state model.

Hybridization

Localized valence-bond model that combines atomic orbitals into directional bonding orbitals.

Hydrocarbon

An organic compound composed only of carbon and hydrogen; the parent class from which nearly every other class of organic molecule is built by adding functional groups.

Hydrogen bond

A strong dipole–dipole attraction between a hydrogen atom bonded to N, O, or F and a lone pair on another N, O, or F atom. It is the strongest common intermolecular force and explains water's unusual properties.

e.g. Hydrogen bonding holds the two strands of DNA together and gives ice its open, less-dense structure.

Hydrogen bonding

Strong, directional intermolecular attraction involving an H bonded to N, O, or F and a suitable lone pair in the standard introductory model.

I

Ideal gas

A model gas whose particles have negligible volume and no attractions, so it obeys PV = nRT exactly. Real gases approach ideal behaviour at low pressure and high temperature.

e.g. At STP one mole of an ideal gas occupies 22.41 L; real gases are within about 1% of this.

Related: Ideal gas law Pressure Kinetic molecular theory MoleDeveloped in: Module 14

Ideal gas law

The relationship PV = nRT, obeyed exactly by an ideal gas and closely by real gases at low pressure and high temperature.

Indicator

A substance, usually a weak acid or base itself, that changes color over a specific pH range — used to visually flag when a titration has reached its endpoint, as a practical stand-in for the true equivalence point a color change can only approximate.

e.g. Phenolphthalein is colorless below pH 8.2 and pink above pH 10 — chosen for titrations where the equivalence point falls in that range, and a poor choice for one that doesn't.

Integrated rate law

A rate law algebraically rearranged to express concentration as a function of elapsed time, rather than rate as a function of concentration. This is what actually lets you answer 'how much reactant is left after 10 minutes' instead of only 'how fast is it disappearing right now.'

e.g. For a first-order reaction, the integrated rate law is the same exponential-decay math used for radioactive half-life — both describe a quantity that loses a constant fraction of itself per unit time.

Intermolecular forces

Attractions between whole molecules — dispersion forces, dipole–dipole attraction, and hydrogen bonding — that are far weaker than bonds but set boiling point, melting point, viscosity, and solubility.

e.g. Water boils far higher than H₂S of similar size because water molecules hydrogen-bond to one another.

Internal energy

The total energy contained within a system — every particle's kinetic and potential energy summed together. It cannot be measured directly or in an absolute sense; what's tracked is always a change in internal energy (ΔE), which by the first law of thermodynamics equals the heat added to the system plus the work done on it.

e.g. Striking a match converts stored chemical potential energy (internal energy of the reactants) into heat and light released to the surroundings — the system's internal energy drops as the surroundings' rises.

Related: First law of thermodynamics Energy EnthalpyDeveloped in: Module 8

Ion

An atom or group of atoms with a net electric charge because it has gained or lost one or more electrons. Losing electrons gives a positive cation; gaining electrons gives a negative anion.

e.g. A sodium atom loses one electron to become Na⁺; a chlorine atom gains one to become Cl⁻; together they form the ionic compound NaCl.

Related: Cation Anion Ionic bond Electron ElectrolyteDeveloped in: Module 2 Module 6 Module 12

Ion–dipole attraction

Attraction between an ion and a polar molecule.

Ionic bond

The electrostatic attraction between oppositely charged ions, formed when one atom transfers electrons to another — typically a metal to a non-metal. Ionic compounds form extended lattices, not molecules.

e.g. NaCl is held together by the attraction between Na⁺ and Cl⁻ ions repeating through the crystal.

Related: Covalent bond Cation Anion Electronegativity Formula unitDeveloped in: Module 12

Ionization energy

Energy required to remove an electron from a gaseous species under defined conditions.

Isoelectronic

Having the same number of electrons.

Isoelectronic series

A set of atoms and/or ions that all have the same number of electrons, useful for isolating how nuclear charge alone (with electron count held constant) affects size and other properties.

Isomer

One of two or more compounds with the same molecular formula but a different arrangement of atoms, giving different physical or chemical properties.

Isotope

Atoms of the same element (same number of protons) that differ in the number of neutrons, and therefore in mass. Isotopes of an element behave almost identically chemically but can differ in nuclear stability.

e.g. Carbon-12 and carbon-14 both have 6 protons; carbon-14 has 8 neutrons and is radioactive, which is what makes radiocarbon dating possible.

Related: Atom Atomic mass Atomic numberDeveloped in: Module 2

Isotope notation

The standard way of writing a specific nuclide, mass number as a left superscript and atomic number as a left subscript on the element symbol, e.g. ¹⁴₆C.

K

Ketone

An organic compound containing a carbonyl group (C=O) bonded to two carbon atoms, distinguishing it from an aldehyde, where the carbonyl carbon bonds to at least one hydrogen.

Kinetic molecular theory

Particle model that explains ideal gas behavior through motion, collision, and negligible attractions/volume.

L

Lattice energy

Energy change associated with separating or forming an ionic lattice under a defined convention; always check sign definition.

Law of conservation of mass

Mass is neither created nor destroyed in a chemical reaction; a balanced equation expresses this atom by atom.

Le Chatelier's principle

If a system at equilibrium is disturbed — a reactant added or removed, volume or pressure changed, temperature changed — the equilibrium position shifts in whichever direction partially counteracts the disturbance, reaching a new balance rather than returning to the old one.

e.g. Adding more N₂ to the Haber process (N₂ + 3H₂ ⇌ 2NH₃ at equilibrium) shifts the reaction toward more NH₃ production, consuming some of the added N₂ but not all of it.

Lewis structure

A diagram that shows every valence electron in a molecule or ion as bonding pairs and lone pairs, drawn to give each atom (usually) an octet. Lewis structures are the starting point for predicting shape and polarity.

e.g. The Lewis structure of CO₂ shows two C=O double bonds and two lone pairs on each oxygen.

Related: Covalent bond VSEPR Formal charge Lone pairDeveloped in: Module 12 Module 13

Limiting reactant

The reactant that runs out first and therefore sets the maximum amount of product. It is found by converting each reactant to moles, dividing by its coefficient, and taking the smallest result; the others are in excess.

e.g. Burning 4 mol H₂ with 1 mol O₂ (needs 2 mol) makes O₂ the limiting reactant; only 2 mol H₂O forms and 2 mol H₂ is left over.

Lipid

A broad class of biological molecules, including fats and oils, that are largely nonpolar and therefore insoluble in water; used for long-term energy storage, cell membranes, and signaling.

London dispersion

See dispersion force.

Lone pair

A pair of valence electrons localised on one atom and not shared in a bond.

M

Mass number

Whole-number total of protons and neutrons in a nuclide.

Matter

Anything that has mass and occupies space.

Mixture

A physical combination of two or more substances whose proportions can vary and which keep their own properties. Mixtures are separated by physical methods such as filtration, distillation, or chromatography.

e.g. Salt water is a homogeneous mixture; its salinity can be anything from a trace to saturation, and evaporation recovers the salt unchanged.

Related: Compound Element Physical changeDeveloped in: Module 1

Molar mass

The mass of one mole of a substance, in grams per mole, numerically equal to its atomic or formula mass in unified atomic mass units. It is the conversion factor between the mass of a sample and the amount of substance in it.

e.g. The molar mass of CO₂ is 12.01 + 2(16.00) = 44.01 g/mol, so 88.02 g of CO₂ is 2.000 mol.

Related: Mole Atomic mass Chemical formula StoichiometryDeveloped in: Module 3 Module 5

Molarity

A concentration unit: moles of solute per litre of solution, symbol M. Molarity lets a volume of solution be treated as a source of a known number of moles for stoichiometry.

e.g. Dissolving 0.50 mol NaCl and making up to 250 mL gives 2.0 M NaCl; 50 mL of it delivers 0.10 mol.

Related: Solute Solvent Dilution Concentration StoichiometryDeveloped in: Module 6

Mole

The SI unit for amount of substance: exactly 6.02214076 × 10²³ elementary entities (the Avogadro constant). The mole is the bridge between the number of particles, which is uncountable directly, and mass, which can be weighed.

e.g. One mole of water is 6.022 × 10²³ molecules and has a mass of 18.02 g. Half a mole of NaCl contains 3.011 × 10²³ formula units.

Molecular formula

The formula that states the actual number of atoms of each element in one molecule of a compound. It is always a whole-number multiple of the empirical formula.

e.g. Benzene's molecular formula is C₆H₆; its empirical formula is CH.

Related: Empirical formula Chemical formula Molar massDeveloped in: Module 3

Molecular geometry

The three-dimensional arrangement of the atoms in a molecule, predicted from the electron-domain count by VSEPR.

Molecule

Electrically neutral group of atoms joined by covalent bonding in the usual introductory definition.

Monomer

A comparatively small molecule that reacts with many identical or chemically complementary copies of itself, linking end to end into the long repeating chain of a polymer. The same monomer-to-polymer relationship describes a synthetic plastic built from a single repeating unit and a protein built from twenty different amino acid monomers strung together in a specific sequence.

e.g. Ethylene gas (C2H4) is the monomer that links into polyethylene, the polymer used in plastic bags and bottles; amino acids are the monomers that link into every protein, sequence determining function instead of just repeating identically.

N

Nernst equation

An equation that corrects a galvanic cell's potential away from its standard value when the reacting concentrations are not the standard 1 M (or 1 atm for a gas) that a standard reduction potential assumes. It's the reason a real battery's voltage isn't fixed — it drifts downward as the cell discharges and its reactant concentrations fall toward its product concentrations, reaching zero exactly when the cell reaches equilibrium and Q equals K.

e.g. A freshly made battery briefly reads slightly above its standard cell potential if started from reactant-favoring concentrations, then reads progressively lower as it discharges — the same Q-versus-K logic as any other reaction, just measured as a voltage instead of tracked as a concentration.

Net ionic equation

An equation that shows only the species that actually change during a reaction in solution, with spectator ions removed. It captures the chemistry common to many reactions that share the same essential step.

e.g. Every strong-acid/strong-base neutralisation has the same net ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l).

Related: Spectator ion Precipitate Acid BaseDeveloped in: Module 6

Neutralization

The reaction of an acid with a base to produce water and, usually, a dissolved salt. A strong-acid/strong-base neutralisation goes essentially to completion and shares the same net ionic equation regardless of which acid or base is used.

e.g. HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l); the net ionic equation is just H⁺(aq) + OH⁻(aq) → H₂O(l).

Related: Acid Base pH Net ionic equationDeveloped in: Module 4 Module 6 Module 7

Neutron

Electrically neutral nucleon.

Nonelectrolyte

A substance that dissolves in water without forming ions, so its solution does not conduct electricity — most molecular compounds, such as sugar, behave this way.

Nuclear fission

The splitting of a large, unstable nucleus into two smaller nuclei, releasing a large amount of energy along with a few additional neutrons. Those released neutrons can go on to split further nuclei, which is what makes a self-sustaining chain reaction — and a nuclear reactor — possible.

e.g. A nuclear power plant sustains a controlled fission chain reaction in uranium-235, using the released heat to boil water and turn a turbine — the same basic principle as a coal plant, with a very different heat source.

Nuclear fusion

The combining of two small nuclei into one larger nucleus, releasing far more energy per reaction than fission does. Fusion requires extreme temperature and pressure to force two positively charged nuclei close enough to overcome their mutual electrical repulsion — conditions found naturally in stars, and only briefly and partially reproduced on Earth.

e.g. The sun's energy comes from fusing hydrogen nuclei into helium — the same basic reaction researchers are trying to sustain and harness in experimental fusion reactors on Earth.

Related: Nuclear fission Binding energy

Nucleic acid

A biological polymer, either DNA or RNA, built from nucleotide monomers, that stores and transmits genetic information.

O

Octet rule

The tendency of main-group atoms to gain, lose, or share electrons until surrounded by eight valence electrons — the same stable configuration as the nearest noble gas. It is a strong guide for drawing Lewis structures and predicting ionic charges, not an absolute law: hydrogen and helium fill with two electrons, boron compounds are commonly electron-deficient, and elements in period 3 and below can expand past an octet.

e.g. In CO₂, both C=O double bonds give carbon and each oxygen eight electrons; in BF₃, boron has only six and does not violate the rule so much as sit outside it.

Orbital

Quantum-mechanical state/spatial probability description associated with an electron in an atom.

Oxidation

The loss of electrons by a species, shown by an increase in its oxidation number. Oxidation always accompanies reduction of another species; together they make a redox reaction.

e.g. In Zn + Cu²⁺ → Zn²⁺ + Cu, zinc is oxidised (0 → +2) as it hands two electrons to the copper ion.

Oxidation number

A formal charge assigned to an atom in a compound or ion by a fixed set of rules, used to track electron transfer even in covalent bonds where charge isn't really localised. Rising oxidation numbers signal oxidation; falling ones signal reduction.

e.g. In SO₄²⁻: oxygen is assigned −2 (four of them, −8 total), and since the ion's overall charge is −2, sulfur must be +6 for the sum to balance.

Oxidizing agent

The species that causes another species to be oxidised by accepting its electrons — and is itself reduced in the process. Strong oxidizing agents are eager electron acceptors.

e.g. In Zn + Cu²⁺ → Zn²⁺ + Cu, Cu²⁺ is the oxidizing agent: it oxidises zinc and is itself reduced to Cu.

P

Partial pressure

Pressure contribution assigned to one gas in an ideal mixture.

Pauli exclusion principle

No two electrons in an atom share the same four quantum numbers.

Peptide bond

The covalent bond linking two amino acids, formed between one amino acid's carboxylic acid group and the next one's amine group, releasing a water molecule in the process — the same condensation logic as esterification, run on amino acids instead of a carboxylic acid and an alcohol.

e.g. A protein a few hundred amino acids long is held together end-to-end by a few hundred peptide bonds, one per junction, before it ever folds into its functional shape.

Percent composition

The mass percent each element contributes to a compound's total molar mass; reversing the calculation from a sample's percent composition is how an empirical formula is found.

Percent yield

The actual yield of a reaction divided by its theoretical yield, times 100%. It measures how efficiently a reaction produced the desired product; losses, side reactions, and incomplete reaction all lower it.

e.g. Recovering 9.2 g of product when 11.0 g was predicted is a percent yield of 9.2 / 11.0 × 100% = 84%.

Periodic table

The arrangement of the elements by increasing atomic number into periods and groups whose members share chemical behaviour.

Periodic trend

A regular change in an atomic property — radius, ionisation energy, electronegativity — across a period or down a group.

A logarithmic measure of hydrogen-ion concentration: pH = −log[H⁺]. At 25 °C, pH 7 is neutral, below 7 is acidic, above 7 is basic, and pH + pOH = 14.

e.g. [H⁺] = 1.0 × 10⁻³ M gives pH 3.00 (acidic); each whole pH unit is a tenfold change in [H⁺].

Related: Acid Base Neutralization Scientific notationDeveloped in: Module 6 Module 7

Phase diagram

A map of which phase — solid, liquid, or gas — is stable at every combination of temperature and pressure. The lines are phase boundaries, where two phases coexist; they meet at the triple point and the liquid–gas line ends at the critical point.

e.g. Following water's phase diagram straight up from 1 atm and 0 °C crosses the melting line: ice becomes liquid, which is why applying pressure can melt ice at a fixed temperature.

Related: Triple point Critical point Vapor pressureDeveloped in: Module 14

Photon

Quantum of electromagnetic radiation energy.

Physical change

Change in state or form without change in chemical identity.

Pi bond

A covalent bond formed by side-by-side overlap of parallel, unhybridized p orbitals, placing electron density above and below the bond axis rather than along it. Pi bonds only occur alongside a sigma bond in double and triple bonds, and restrict rotation around that bond — the reason cis/trans isomers exist.

e.g. A C≡C triple bond is one sigma bond plus two pi bonds, oriented at right angles to each other.

Related: Sigma bond Covalent bond Valence bond modelDeveloped in: Module 13

Polar covalent bond

A covalent bond whose shared electrons are pulled toward the more electronegative atom, giving each end a partial charge.

Polar molecule

Molecule with a nonzero net dipole moment.

Polarizability

How easily an atom's or molecule's electron cloud is distorted into a temporary, induced dipole by a nearby charge or field. Larger, more electron-rich, more loosely held electron clouds are more polarizable, which strengthens dispersion forces and raises boiling point.

e.g. Iodine (I₂) is a solid at room temperature and fluorine (F₂) is a gas, even though both are nonpolar diatomics — iodine's far larger, more polarizable electron cloud gives it much stronger dispersion forces.

Polyatomic ion

Covalently connected group of atoms with a net charge.

Polymer

A large molecule built from many repeating smaller units (monomers) linked by covalent bonds into a chain or network. The concept spans synthetic materials (polyethylene, nylon) and biological macromolecules (DNA is a polymer of nucleotides; proteins are polymers of amino acids) — the same bonding logic underlies both.

e.g. Polyethylene terephthalate (PET), the plastic in most beverage bottles, is a polymer built by repeatedly linking a diol and a diacid monomer through ester bonds.

Precipitate

An insoluble solid that forms and separates when two solutions are mixed. Whether a precipitate forms is predicted from solubility rules applied to the possible ion combinations.

e.g. Mixing AgNO₃(aq) and NaCl(aq) forms a white AgCl precipitate: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

Precision

How reproducible repeated measurements of the same quantity are — how tightly they cluster, regardless of whether they are correct. Precision is limited by random error and by the finest division of the instrument.

e.g. Readings of 4.512, 4.514, 4.513 mL are precise to about ±0.001 mL. Whether they are accurate depends on the burette's calibration.

Related: Accuracy Significant figures Random errorDeveloped in: Module 0 Module 1

Pressure

Force per area.

Product

Species on the product side of a represented reaction.

Protein

A biological polymer built from amino acid monomers linked end-to-end by peptide bonds, then folded into a specific three-dimensional shape. That folded shape — not just the sequence of amino acids — is what determines a protein's function, which is why denaturation (unfolding without breaking bonds) destroys function even though every atom is still present.

e.g. Hemoglobin's folded shape creates a pocket that binds oxygen; unfold it (denature it) and the amino acid sequence is unchanged, but it can no longer carry oxygen.

Proton

Positively charged nucleon; also H+ in acid–base language depending on context.

Q

Quantum

Discrete amount associated with quantized energy exchange.

Quantum number

One of a set of integers (or half-integers, for spin) that together specify the state of an electron in an atom. The principal quantum number n sets shell size and energy; the angular momentum number l sets orbital shape (s, p, d, f); the magnetic number sets orbital orientation; the spin number distinguishes the two electrons that can share one orbital.

e.g. An electron in a 2p orbital has n = 2 and l = 1; the Pauli exclusion principle requires its orbital-mate to have the opposite spin quantum number.

R

Random error

Unpredictable scatter in repeated measurements; it lowers precision but tends to average out over many trials.

Rate constant

The proportionality constant k in a rate law (rate = k[A]^m[B]^n), specific to one reaction at one temperature. Unlike the rate itself — which falls as reactants are consumed — k stays fixed throughout a reaction at constant temperature; it only changes if temperature (or a catalyst) changes.

e.g. Doubling a reactant's concentration might double the rate, but it never changes k — k only shifts if you heat the reaction up, cool it down, or add a catalyst.

Rate law

An experimentally measured equation, rate = k[A]ᵐ[B]ⁿ, relating reaction rate to reactant concentrations. The exponents (reaction orders) must be determined by experiment — they are not simply read off the balanced equation's coefficients, which is one of the most common wrong turns when this topic is first introduced.

e.g. For a reaction found experimentally to be first-order in A and zero-order in B, doubling [A] doubles the rate, but doubling [B] does nothing to it — even if B's coefficient in the balanced equation is not 1.

Related: Reaction rate Collision theory

Reactant

Starting species in a represented reaction.

Reaction order

The exponent on a given reactant's concentration in an experimentally measured rate law. A reaction can be zero-order (rate independent of that reactant), first-order (rate directly proportional), second-order, and so on — and the order almost never matches that reactant's coefficient in the balanced equation, which is the single most common wrong assumption when this idea is first introduced.

e.g. A reaction found to be second-order in A means quadrupling [A] quadruples the rate squared — i.e., increases it 4x — even if A's coefficient in the balanced equation is 1.

Reaction quotient

The identical ratio of product to reactant concentrations as the equilibrium constant K, but evaluated from whatever concentrations a reaction happens to have right now — before, after, or at equilibrium. Comparing Q to K tells you which direction a reaction still has left to run: Q < K means it proceeds forward, Q > K means it runs in reverse, and Q = K means it has already arrived.

e.g. Mix reactants and products together and Q starts wherever those starting concentrations put it. If Q comes out smaller than the reaction's known K, the reaction shifts forward — consuming reactants and forming products — until Q rises to meet K.

Reaction rate

How quickly a reaction proceeds, measured as reactant concentration falling (or product concentration rising) per unit time — typically mol/(L·s). Distinct from whether a reaction is thermodynamically favorable (ΔH, ΔG): a reaction can release enormous energy and still be immeasurably slow without enough activation energy or a catalyst.

e.g. Diamond converting to graphite releases energy (it's thermodynamically favorable) but the reaction rate at room temperature is so close to zero that diamonds are, in practice, forever.

Redox reaction

A reaction in which electrons are transferred between species, so one is oxidised and another reduced. Combustion, corrosion, respiration, and battery chemistry are all redox processes.

e.g. 4 Fe + 3 O₂ → 2 Fe₂O₃ (rusting) transfers electrons from iron to oxygen.

Related: Oxidation Reduction Oxidation numberDeveloped in: Module 4 Module 7

Reducing agent

The species that causes another species to be reduced by donating electrons to it — and is itself oxidised in the process. Metals low on the activity series tend to be strong reducing agents.

e.g. In Zn + Cu²⁺ → Zn²⁺ + Cu, Zn is the reducing agent: it reduces Cu²⁺ and is itself oxidised to Zn²⁺.

Reduction

The gain of electrons by a species, shown by a decrease in its oxidation number. The species that causes it (by being oxidised) is the reducing agent.

e.g. In Zn + Cu²⁺ → Zn²⁺ + Cu, the copper ion is reduced (+2 → 0).

Related: Oxidation Redox reaction Oxidation numberDeveloped in: Module 4 Module 7

Resonance

Use of multiple Lewis contributors to represent delocalized electron distribution in one species.

S

Scientific notation

Writing a number as a coefficient between 1 and 10 multiplied by a power of ten. It keeps very large and very small quantities readable and makes the number of significant figures unambiguous.

e.g. The Avogadro constant is 6.022 × 10²³ mol⁻¹; a wavelength of 0.000000550 m is 5.50 × 10⁻⁷ m.

Related: Significant figures MoleDeveloped in: Module 0

Shielding

The reduction in nuclear attraction an outer electron feels because inner-shell electrons sit between it and the nucleus, partially blocking the pull. Shielding is why effective nuclear charge is always less than the full atomic number for any electron past the first, and why it changes much more slowly down a group (each new shell shields the one before) than it does across a period (new electrons enter the same shell and shield each other poorly).

e.g. A valence electron in potassium (Z = 19) is shielded by 18 core electrons, so it feels an effective nuclear charge closer to +1 than to +19.

Related: Effective nuclear charge Periodic trend OrbitalDeveloped in: Module 11

Sigma bond

A covalent bond with electron density concentrated directly along the line connecting two nuclei, formed by head-on orbital overlap. Every single, double, and triple bond contains exactly one sigma bond, which is what allows free rotation around a single bond but not around a double or triple bond.

e.g. The C–C bond in ethane (single) is one sigma bond; the C=C bond in ethylene (double) is one sigma bond plus one pi bond.

Related: Pi bond Covalent bond Valence bond modelDeveloped in: Module 13

Significant figures

The digits in a measurement that carry meaning: every certain digit plus one estimated digit. Significant figures communicate the precision of a measurement, and calculations may not report more of them than the least precise input allows.

e.g. 0.004560 has four significant figures. A density found from 12.3 g and 4.1 mL is reported as 3.0 g/mL — two significant figures, matching the volume.

Related: Precision Accuracy Scientific notationDeveloped in: Module 0

Solubility

Equilibrium extent to which a substance dissolves under specified conditions.

Solubility product

An equilibrium constant, Ksp, describing the equilibrium between a sparingly soluble ionic solid and its dissolved ions in a saturated solution. Unlike a general equilibrium constant, Ksp has no denominator terms — the solid itself doesn't appear, since a pure solid's 'concentration' doesn't meaningfully change.

e.g. Silver chloride's very small Ksp (about 1.8 × 10⁻¹⁰) is why so little of it actually dissolves in water — most of it stays as an undissolved solid at the bottom of the beaker.

Solute

The substance dissolved in a solution — usually the component present in the smaller amount. It is dispersed at the particle level throughout the solvent.

e.g. In 2 M aqueous NaCl, NaCl is the solute and water is the solvent.

Related: Solvent Concentration Molarity ElectrolyteDeveloped in: Module 6

Solution

Homogeneous mixture.

Solvent

The substance that dissolves the solute, usually the component present in the larger amount and the one that keeps its physical state. Water is the most common solvent in general chemistry.

e.g. Ethanol is the solvent in tincture of iodine; water is the solvent in every aqueous solution.

Related: Solute Concentration AqueousDeveloped in: Module 6

Specific heat capacity

Energy required per unit mass per unit temperature change.

Spectator ion

An ion that appears unchanged on both sides of an ionic equation and takes no part in the reaction. Cancelling spectator ions from the complete ionic equation leaves the net ionic equation.

e.g. In the AgCl precipitation, Na⁺ and NO₃⁻ are spectator ions.

Related: Net ionic equation Precipitate Aqueous IonDeveloped in: Module 6

Spontaneous process

A process that proceeds on its own under given conditions once started, without a continuous outside energy input — determined by Gibbs free energy, not by how fast it happens. 'Spontaneous' is a thermodynamics word, not a speed word, which is easy to mix up with reaction rate.

e.g. Diamond converting to graphite is spontaneous (negative ΔG) — but so slow at room temperature that it's not observed on any human timescale, exactly the same rate-vs-spontaneity gap reaction rate calls out for diamond.

Standard enthalpy of formation

The enthalpy change when exactly one mole of a compound forms from its constituent elements, each in its standard state, at standard conditions. Tabulated values let a reaction's enthalpy be calculated without running a calorimetry experiment: ΔH°_rxn = ΣνΔH°_f(products) − ΣνΔH°_f(reactants). By definition, an element in its own standard state has a standard enthalpy of formation of exactly zero.

e.g. ΔH°_f of O₂(g) is 0 by definition; ΔH°_f of CO₂(g) is −393.5 kJ/mol, the heat released forming one mole of CO₂ from graphite and O₂.

Related: Enthalpy Hess's law Standard state CalorimetryDeveloped in: Module 9

Standard reduction potential

A measured voltage, referenced against the hydrogen electrode's defined 0.00 V, describing how readily a given half-reaction accepts electrons (is reduced) under standard conditions. A more positive value means a stronger oxidizing agent; combining two half-reactions' potentials predicts whether a redox reaction is spontaneous and what voltage a galvanic cell built from them would produce.

e.g. Cu²⁺ + 2e⁻ → Cu has a standard reduction potential of +0.34 V; Zn²⁺ + 2e⁻ → Zn has −0.76 V — the more positive copper half-reaction runs forward while the zinc half-reaction runs in reverse, which is exactly the classic zinc–copper galvanic cell.

Related: Half-reaction Galvanic cell

Standard state

Defined reference state used for thermodynamic quantities.

State function

Property whose change depends on endpoint states, not path.

Stereoisomer

One of two or more compounds sharing the same molecular formula and the same atom-to-atom connectivity, but differing in their three-dimensional spatial arrangement — a narrower, more specific case of isomer, which only requires the same formula.

e.g. Cis and trans forms of a molecule with a double bond are stereoisomers: identical atoms and bonds, but the arrangement in space around that double bond differs, often changing physical properties like melting point.

Related: Isomer Chirality

Stoichiometry

The quantitative relationships between amounts of reactants and products in a chemical reaction, read from the coefficients of the balanced equation. A stoichiometry problem converts a known quantity to moles, applies the mole ratio, then converts to the requested quantity.

e.g. From 2 H₂ + O₂ → 2 H₂O, 5.0 mol H₂ reacts with 2.5 mol O₂ to give 5.0 mol H₂O.

Strong electrolyte

Solute producing ions extensively in solution under the stated conditions.

Subscript

A small number written after an element symbol in a chemical formula, giving the count of that atom in one formula unit or molecule — not to be confused with the coefficient in front of a formula, which multiplies the whole thing.

Substitution reaction

An organic reaction in which one atom or group of atoms in a molecule is replaced by a different atom or group, with the rest of the molecule's structure unchanged.

Surroundings

Everything outside the defined system that can exchange energy or matter with it — in a coffee-cup calorimeter, the water and cup are commonly treated as the system, and the room is the surroundings.

System

Portion of the universe selected for study.

Systematic error

A consistent bias that shifts every measurement the same way; it lowers accuracy but not precision.

T

Temperature

Property related to the distribution of particle kinetic energy and thermal state.

Theoretical yield

The maximum mass or amount of product predicted by stoichiometry from the limiting reactant, assuming the reaction goes to completion with no losses.

e.g. If the limiting reactant allows 0.250 mol of a product with molar mass 44.0 g/mol, the theoretical yield is 11.0 g.

Titrant

The solution of known concentration added from a burette during a titration to react with a measured volume of the analyte.

Triple point

The one specific temperature and pressure at which a substance's solid, liquid, and gas phases all coexist in equilibrium at once. Every pure substance has exactly one.

e.g. Water's triple point, 0.01 °C and 611.657 Pa, is precise enough that it once defined the kelvin itself.

Related: Phase diagram Critical point Vapor pressureDeveloped in: Module 14

U

Uncertainty

Quantified or described limitation on knowledge of a measurement or result.

V

Valence bond model

A bonding theory that treats a covalent bond as the overlap of two atomic orbitals — often hybridized — each contributing one electron to a shared pair localized between the two nuclei. It is the model behind sigma/pi bond descriptions and hybridization, and is taught alongside (not in place of) Lewis structures and VSEPR.

e.g. Methane's four equivalent C–H bonds are explained by carbon's four sp³ hybrid orbitals, each overlapping head-on with a hydrogen 1s orbital.

Related: Sigma bond Pi bond Hybridization Covalent bondDeveloped in: Module 13

Valence electron

Electron in the occupied outer region most involved in main-group bonding and chemical behavior.

Vapor pressure

The pressure exerted by a substance's own vapor once it reaches equilibrium with its liquid (or solid) at a given temperature — the point where molecules leave and rejoin the surface at equal rates. Stronger intermolecular forces mean fewer molecules have enough energy to escape, so vapor pressure is lower.

e.g. Water's vapor pressure at 25 °C is about 23.8 torr; diethyl ether, with much weaker intermolecular attractions, is over 400 torr at the same temperature and evaporates visibly faster.

VSEPR

Valence-shell electron-pair repulsion: the model that predicts molecular shape by arranging the electron domains (bonds and lone pairs) around a central atom as far apart as possible.

e.g. Four electron domains around carbon in CH₄ point to the corners of a tetrahedron, giving 109.5° bond angles.

W

Wavelength

Distance between corresponding points on successive wave cycles.

Weak acid

An acid that only partially ionizes in water, establishing an equilibrium rather than dissociating completely like a strong acid; its strength is described by Ka, not just its identity.

Weak electrolyte

Solute that forms ions only partly under the stated solution conditions.

Work

Energy transfer through force acting through distance or another generalized force–displacement pair.