Unit 4: Chemical Reactions

Chemical equation

A symbolic representation of a chemical reaction showing reactants, products, and their relative amounts using formulas and coefficients.

Balanced chemical equation

A chemical equation in which the number of atoms of each element is the same on both sides, reflecting conservation of atoms.

Coefficient

A number placed in front of a chemical formula in an equation that indicates relative amounts (mole ratios) of reactants and products.

Mole ratio

A ratio of moles of two substances in a reaction, taken from the coefficients of a balanced chemical equation.

Conservation of atoms

Principle that atoms are rearranged in chemical reactions but not created or destroyed in ordinary chemical processes.

State symbols

Notation in equations—(s), (l), (g), (aq)—that indicates physical state and often hints at driving forces (e.g., precipitate or gas formation).

Aqueous (aq)

State symbol meaning dissolved in water as ions and/or molecules; important for determining which species dissociate in ionic equations.

Precipitate

An insoluble solid that forms from mixing solutions; often a driving force for an aqueous reaction and written with (s).

Diatomic elements

Elements that exist naturally as two-atom molecules in their standard form: H2, N2, O2, F2, Cl2, Br2, I2.

Balancing by coefficients (not subscripts)

Balancing equations by changing coefficients (counts of particles), not subscripts (which would change the identity of the substance).

Reaction type

A classification (e.g., synthesis, decomposition, combustion) used to help predict products and explain reactions, not just patterns to memorize.

Synthesis (combination) reaction

Reaction in which two or more reactants form a single product; general pattern A + B → AB.

Decomposition reaction

Reaction in which one compound breaks into two or more products; general pattern AB → A + B (often requires heat).

Combustion reaction

Reaction with O2 that typically forms oxides; used frequently for predicting products and practicing balancing.

Hydrocarbon combustion

Combustion of a compound containing C and H (and sometimes O) that produces CO2 and H2O (for complete combustion).

Precipitation reaction

An aqueous reaction where ions combine to form an insoluble ionic compound (solid precipitate) when two solutions are mixed.

Solubility guidelines

Rules that help predict whether an ionic compound is soluble (aq) or insoluble (s), determining whether a precipitate forms.

Acid–base neutralization

Reaction where an acid (H+ source) reacts with a base (OH− source) to form water and a salt; water formation is a common driving force.

Net ionic equation

An equation showing only the species that undergo chemical change in aqueous solution, with spectator ions removed.

Molecular equation

Equation written with compounds as intact formulas (not split into ions), including state symbols.

Complete ionic equation

Equation in which strong electrolytes are written as separate ions; used to identify and cancel spectator ions.

Spectator ion

An ion that appears unchanged on both sides of the complete ionic equation and does not participate in the chemical change.

Strong electrolyte

A substance that dissociates or ionizes essentially completely in water (many soluble ionic compounds and strong acids), so it is written as ions in ionic equations.

Weak electrolyte

A substance that only partially ionizes in water (weak acids/bases); often written mainly as molecules in net ionic equations.

Nonelectrolyte

A substance that dissolves as neutral molecules and produces essentially no ions in solution (e.g., sugar).

Driving forces in aqueous reactions

Factors that favor reaction in water by removing species from solution: formation of a precipitate, a gas, or a weak electrolyte (especially water).

Stoichiometry

Quantitative relationship between reactants and products using a balanced equation’s coefficients as mole ratios.

Limiting reactant

The reactant consumed first, which stops the reaction and determines the maximum amount of product that can form.

Theoretical yield

The maximum possible amount of product predicted by stoichiometry from the limiting reactant.

Excess reactant

A reactant present in more than the stoichiometric amount; it remains after the limiting reactant is used up.

Percent yield

A measure of reaction efficiency: (actual yield ÷ theoretical yield) × 100%.

Percent error

Comparison of experimental to expected value: |experimental − expected| ÷ expected × 100%.

Molarity (M)

Solution concentration defined as moles of solute per liter of solution: M = n/V.

Solution stoichiometry (n = MV)

Method to find moles from solution volume and molarity: n = M·V (with V in liters).

Titration

A technique where a solution of known concentration reacts with an unknown until stoichiometric completion, allowing determination of the unknown concentration.

Titrant

The solution of known concentration added during a titration (often delivered from a buret).

Analyte

The substance/solution of unknown concentration being measured in a titration.

Equivalence point

Point in a titration where stoichiometric amounts have reacted according to the balanced equation (defined by moles, not equal volumes).

Brønsted–Lowry acid–base theory

Model in which acids donate H+ (protons) and bases accept H+; acid–base reactions form conjugate pairs differing by one H+.

Conjugate acid–base pair

Two species related by gain/loss of one H+ (e.g., HC2H3O2 / C2H3O2−).

Amphoteric

Able to act as either an acid or a base depending on the reaction partner; water is a common amphoteric substance.

Dilution (M1V1 = M2V2)

Relationship for dilution (adding solvent) where moles of solute are conserved; not used for chemical reaction stoichiometry.

Gravimetric analysis

Determining amount of an analyte by converting it to a precipitate of known formula, measuring precipitate mass, and using stoichiometry to find the analyte amount.

Combustion analysis

Determining composition (often of C/H/O compounds) by combusting a sample and using measured CO2 and H2O to calculate moles of C and H (and O by mass difference if needed).

Empirical formula

The simplest whole-number ratio of atoms in a compound, often found by converting element amounts to moles and reducing ratios.

Oxidation state

A bookkeeping value assigned to atoms to track electron distribution; used to identify redox processes by changes in oxidation numbers.

Oxidation

Process in which oxidation state increases (loss of electrons) in a redox reaction.

Reduction

Process in which oxidation state decreases (gain of electrons) in a redox reaction.

Half-reaction

An equation showing either oxidation or reduction alone, explicitly including electrons to track electron transfer.

Redox titration (permanganate endpoint)

A titration based on oxidation–reduction; MnO4− is purple and a persistent faint pink color indicates slight excess permanganate at the endpoint.