Difficulty: Advanced | Prerequisites: Chapters 10-12 (balancing equations, stoichiometry), Chapter 16 (enthalpy, entropy, Gibbs free energy)
Source: Chem 1 Final Study Guide, General Chemistry (Purdue University, Chhabra)
Tags: reaction rate, collision theory, activation energy, catalyst, rate law, reaction order, chemical equilibrium, Keq, Le Chatelier's principle, Ksp, solubility product, common ion effect, acid, base, Arrhenius, Bronsted-Lowry, pH, pOH, Ka, Kb, titration, equivalence point, buffer, neutralisation, oxidation, reduction, redox, half-reaction, oxidation number, electrochemical cell, voltaic cell, electrolytic cell, standard reduction potential, electrolysis, galvanising, electroplating
These final five chapters cover the most conceptually demanding material in the course. Chapter 17 asks how fast reactions happen (kinetics). Chapter 18 asks what happens when forward and reverse reactions reach balance (equilibrium). Chapter 19 tackles acids and bases, their strength, pH calculations, and titrations. Chapter 20 introduces electron-transfer reactions (redox), and Chapter 21 applies redox to real electrochemical devices like batteries and electrolysis cells. These topics are heavily interconnected: equilibrium concepts underpin acid-base chemistry, and redox reactions are the foundation of electrochemistry.
Reaction rates depend on concentration, temperature, surface area, and catalysts. At equilibrium, forward and reverse rates are equal and concentrations are constant (though rarely equal). Le Chatelier's principle predicts how a system responds to stress. Acids donate H⁺ ions, bases accept them, and pH quantifies acidity on a logarithmic scale. Redox reactions involve electron transfer, and electrochemistry harnesses that transfer to produce or consume electrical energy.
Reaction rate
The change in concentration of a reactant or product per unit of time, in mol/(L·s). Always expressed as a positive value.
Collision theory
For a reaction to occur, particles must collide with sufficient energy and correct orientation.
Activated complex (transition state)
A temporary, unstable arrangement of atoms at the peak of the energy profile. It either forms products or reverts to reactants.
Activation energy (Ea)
The minimum energy that colliding particles must have to form the activated complex and proceed to products.
Catalyst
A substance that increases the rate of a chemical reaction without being consumed. It lowers the activation energy.
Enzyme
A biological catalyst, typically a protein.
Inhibitor
A substance that slows down or prevents a reaction. The opposite of a catalyst.
Heterogeneous catalyst
A catalyst in a different physical state from the reactants (e.g. a solid catalyst for a gas-phase reaction).
Homogeneous catalyst
A catalyst in the same physical state as the reactants.
Rate law
An equation expressing the relationship between reaction rate and reactant concentrations: Rate = k[A]^m[B]^n.
Rate constant (k)
A proportionality constant specific to a reaction at a given temperature.
Reaction order
The exponent on a reactant's concentration in the rate law. Determined experimentally, not from the balanced equation (unless the reaction occurs in a single elementary step, which is rare).
Instantaneous rate
The rate at a specific moment in time, found from the slope of the tangent line on a concentration-vs-time curve.
Reaction mechanism
The sequence of elementary steps that make up a complex reaction. The slowest step is the rate-determining step.
Intermediate
A substance produced in one elementary step and consumed in a subsequent step. Does not appear in the overall equation.
Chemical equilibrium
A state in which the forward and reverse reaction rates are equal. Concentrations of reactants and products are constant but usually not equal.
Equilibrium constant (Keq)
The ratio of product concentrations to reactant concentrations, each raised to the power of its coefficient: Keq = [C]^c[D]^d / [A]^a[B]^b. Unitless. Constant at a given temperature.
Le Chatelier's principle
When a stress is applied to a system at equilibrium, the system shifts to counteract the stress and re-establish equilibrium.
Homogeneous equilibrium
All reactants and products are in the same physical state.
Heterogeneous equilibrium
Reactants and products are in more than one physical state. Pure solids and liquids are omitted from the Keq expression.
Solubility product constant (Ksp)
The equilibrium constant for the dissolving of a sparingly soluble ionic compound. Ksp = [cation]^x[anion]^y.
Ion product (Qsp)
The product of ion concentrations at any given moment. Compare to Ksp: Qsp < Ksp (unsaturated, no precipitate), Qsp = Ksp (saturated), Qsp > Ksp (precipitate forms).
Common ion effect
Adding a solution that shares an ion with the equilibrium system decreases the solubility of the sparingly soluble salt. Explained by Le Chatelier's principle.
Arrhenius acid
A substance that produces H⁺ ions in aqueous solution.
Arrhenius base
A substance that produces OH⁻ ions in aqueous solution.
Bronsted-Lowry acid
A proton (H⁺) donor.
Bronsted-Lowry base
A proton (H⁺) acceptor.
Conjugate acid-base pair
A pair of species that differ by one proton. Every Bronsted-Lowry reaction involves two conjugate pairs.
Amphoteric
A substance that can act as either an acid or a base (e.g. water).
Monoprotic acid
Donates one H⁺ per molecule (e.g. HCl, HNO₃).
Polyprotic acid
Donates more than one H⁺, losing them one at a time in successive ionisation steps (diprotic: H₂SO₄; triprotic: H₃PO₄).
Strong acid
Ionises completely in water. Examples: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.
Weak acid
Ionises only partially. Examples: HF, HCN, HC₂H₃O₂ (acetic acid), H₂CO₃.
Strong base
Dissociates completely into OH⁻ and metal ions. Examples: Group 1A and 2A hydroxides (NaOH, KOH, Ca(OH)₂, Ba(OH)₂).
Weak base
Does not fully ionise. Examples: NH₃, CH₃NH₂.
Ka (acid ionisation constant)
The equilibrium constant for the ionisation of a weak acid. Smaller Ka = weaker acid.
Kb (base ionisation constant)
The equilibrium constant for the ionisation of a weak base. Smaller Kb = weaker base.
Ion product constant for water (Kw)
Kw = [H⁺][OH⁻] = 1.0 x 10⁻¹⁴ at 298 K.
pH
pH = –log[H⁺]. Scale from 0 to 14 at 298 K. Below 7 = acidic; above 7 = basic; 7 = neutral.
pOH
pOH = –log[OH⁻]. pH + pOH = 14.
Self-ionisation of water
Water produces both H⁺ and OH⁻: 2H₂O ⇌ H₃O⁺ + OH⁻.
Neutralisation
An acid-base reaction producing a salt and water.
Titration
A method for determining the concentration of an unknown solution by reacting it with a solution of known concentration (the standard solution).
Equivalence point
The point in a titration where moles of H⁺ from the acid equal moles of OH⁻ from the base. The pH at the equivalence point is not always 7; it depends on the strengths of the acid and base.
End point
The point where the indicator changes colour. Ideally coincides with the equivalence point.
Buffer
A solution that resists changes in pH when small amounts of acid or base are added. Made from a weak acid and its conjugate base, or a weak base and its conjugate acid.
Salt hydrolysis
When a salt dissolves, its ions may react with water, affecting pH. Strong acid + strong base salt: neutral. Strong acid + weak base salt: acidic. Weak acid + strong base salt: basic.
Anhydride
An oxide that becomes an acid or base when water is added. Nonmetal oxides form acids (CO₂ + H₂O → H₂CO₃). Metal oxides form bases (CaO + H₂O → Ca(OH)₂).
Oxidation
Loss of electrons (OIL: Oxidation Is Losing).
Reduction
Gain of electrons (RIG: Reduction Is Gaining).
Oxidising agent
The substance that gets reduced (gains electrons) and thereby causes another substance to be oxidised.
Reducing agent
The substance that gets oxidised (loses electrons) and thereby causes another substance to be reduced.
Oxidation number
The apparent charge an atom would have if all shared electrons were assigned to the more electronegative atom. Rules: uncombined element = 0; monatomic ion = its charge; F always = –1; O usually = –2 (except peroxides = –1, with F = +2); sum of oxidation numbers = charge of the species.
Half-reaction
One half of a redox reaction, showing either the oxidation or the reduction alone.
Electrochemical cell
A device that uses redox reactions to produce electrical energy (voltaic cell) or uses electrical energy to drive a nonspontaneous redox reaction (electrolytic cell).
Voltaic (galvanic) cell
Converts chemical energy into electrical energy via a spontaneous redox reaction.
Anode
The electrode where oxidation occurs.
Cathode
The electrode where reduction occurs.
Standard reduction potential (E°)
The tendency of a substance to gain electrons, measured against the standard hydrogen electrode (SHE).
Cell potential (E°cell)
E°cell = E°(reduction) – E°(oxidation). Positive = spontaneous; negative = nonspontaneous.
Electrolytic cell
Uses electrical energy to drive a nonspontaneous redox reaction (e.g. recharging a battery, electroplating, separating molten NaCl).
Electrolysis
The use of electrical energy to bring about a chemical reaction.
Galvanising
Coating a metal (typically iron) with zinc to prevent corrosion. Zinc acts as a sacrificial anode.
Electroplating
Depositing a thin layer of metal onto an object by making the object the cathode in an electrolytic cell.
Factors affecting rate
Reactivity of reactants: some substances are inherently more reactive
Concentration: higher concentration = more frequent collisions = faster rate
Surface area: grinding a solid into powder exposes more surface, increasing rate (does not change concentration)
Temperature: higher temperature = higher average KE = more collisions above Ea
Catalysts: lower the activation energy; enzymes are biological catalysts
Inhibitors do the opposite
Heterogeneous catalyst: different phase from reactants
Homogeneous catalyst: same phase
Important distinction
ΔG determines whether a reaction is spontaneous; it says nothing about speed
A reaction can be spontaneous but extremely slow (e.g. diamond → graphite)
Rate law and reaction order
Rate = k[A]^m[B]^n
Orders (m, n) are found experimentally
If doubling [A] doubles the rate: m = 1 (first order in A)
If doubling [A] quadruples the rate: m = 2 (second order in A)
If doubling [A] has no effect: m = 0 (zero order in A)
Overall order = m + n
Only for a single-step (elementary) reaction do the orders equal the stoichiometric coefficients
Reaction mechanisms
Complex reactions consist of multiple elementary steps
The slowest step is the rate-determining step
Intermediates are produced in one step and consumed in the next
Dynamic equilibrium
Forward and reverse reactions proceed at equal rates
Concentrations of reactants and products are constant, but usually not equal
Keq expression
For aA + bB ⇌ cC + dD: Keq = [C]^c[D]^d / [A]^a[B]^b
Keq > 1: products favoured
Keq < 1: reactants favoured
Keq is unitless and temperature-dependent
Pure solids and liquids are omitted from heterogeneous equilibrium expressions
Le Chatelier's principle responses to stress
Concentration: removing product shifts right; removing reactant shifts left (and vice versa)
Pressure/volume: increasing pressure (decreasing volume) shifts toward the side with fewer moles of gas
Temperature: increasing temperature favours the endothermic direction; decreasing favours the exothermic direction
Temperature changes also alter the value of Keq; concentration and pressure changes do not
Ksp and precipitate prediction
Ksp = [cation]^x[anion]^y for a sparingly soluble salt
Qsp < Ksp: unsaturated, no precipitate
Qsp = Ksp: saturated
Qsp > Ksp: precipitate forms, ions crash out until Qsp = Ksp
Common ion effect
Adding an ion already present in the equilibrium shifts the equilibrium toward the solid, reducing solubility
Models
Arrhenius: acid produces H⁺(aq); base produces OH⁻(aq). Limited because it cannot explain NH₃ as a base.
Bronsted-Lowry: acid = H⁺ donor; base = H⁺ acceptor. Every reaction involves two conjugate pairs.
Strong acid → weak conjugate base (ionisation lies to the right)
Weak acid → strong conjugate base (ionisation lies to the left)
pH scale
pH = –log[H⁺]; pOH = –log[OH⁻]; pH + pOH = 14
[H⁺] = 10^(–pH); [OH⁻] = 10^(–pOH)
For strong acids: [H⁺] = [acid] (complete ionisation)
For strong bases: [OH⁻] = [base] (polyprotic case: Ca(OH)₂ gives [OH⁻] = 2 x [Ca(OH)₂])
Using pH to calculate Ka
Given pH, find [H⁺] = 10^(–pH)
For a monoprotic weak acid HA: [H⁺] = [A⁻] at equilibrium
[HA] at equilibrium = initial concentration – [H⁺]
Ka = [H⁺][A⁻] / [HA]
Acid naming rules (review)
Anion ends in "-ide" → hydro____ic acid
Anion ends in "-ite" → ____ous acid
Anion ends in "-ate" → ____ic acid
Neutralisation and titration
Acid + base → salt + water
Titration uses a standard solution (known concentration) to find the unknown concentration
At the equivalence point: moles H⁺ = moles OH⁻
The equivalence point pH depends on acid-base strength:
Strong acid + strong base: pH = 7
Weak acid + strong base: pH > 7
Strong acid + weak base: pH < 7
Can use M(A)V(A) = M(B)V(B) (adjusted for mole ratios)
Salt hydrolysis
Dissolving a salt can shift pH depending on the ions' parent acid/base strength
Strong acid + weak base → acidic solution (strong conjugate acid, weak conjugate base)
Weak acid + strong base → basic solution
Buffers
Weak acid + conjugate base, or weak base + conjugate acid
Resist pH changes when small amounts of acid or base are added
Blood pH is maintained at 7.4 by buffer systems; acidosis (pH drop > 0.3) and alkalosis (pH rise > 0.3) are dangerous
Definitions
Oxidation: loss of electrons; oxidation number increases
Reduction: gain of electrons; oxidation number decreases
OIL RIG mnemonic
Oxidation number rules
Uncombined element: 0
Monatomic ion: equal to its charge
Fluorine: always –1
Oxygen: usually –2 (–1 in peroxides, +2 with fluorine)
Hydrogen: usually +1 (–1 in metal hydrides)
Sum of oxidation numbers = overall charge of the species
Balancing redox equations (oxidation-number method)
Assign oxidation numbers to all atoms
Identify which are oxidised and which are reduced
Write half-reactions showing electron loss and gain
Multiply half-reactions so electrons lost = electrons gained
Add the half-reactions together; cancel common species
Agents
Oxidising agent: gets reduced (accepts electrons)
Reducing agent: gets oxidised (donates electrons)
Voltaic cells
Two half-cells separated by a salt bridge
Anode: oxidation occurs; cathode: reduction occurs
Cell notation: anode | anode product || cathode reactant | cathode
E°(cell) = E°(reduction) – E°(oxidation)
Positive E°(cell) = spontaneous reaction
Batteries
Primary: not rechargeable (e.g. alkaline dry cell)
Secondary: rechargeable (e.g. lead-acid car battery, lithium-ion)
Lithium: lightest metal, lowest standard reduction potential (–3.04 V), excellent for long-lasting batteries
Fuel cells: voltaic cells that continuously convert fuel (H₂) into electrical energy
Corrosion
Spontaneous oxidation of metals (e.g. rusting of iron requires both O₂ and water)
Prevention: paint/coatings, sacrificial anodes (Zn, Mg, Al), galvanising (zinc coating)
Zinc is self-protecting because its oxide layer clings tightly and seals the surface
Electrolysis
Uses electrical energy to force a nonspontaneous reaction
Electrolytic cell: external power source drives the reaction
Applications: electrolysis of molten NaCl (Down's cell), electrolysis of brine (produces H₂, Cl₂, NaOH), aluminium manufacture (Hall-Heroult process), copper purification, electroplating
Ionic compounds conduct electricity only when their ions are free to move (molten or dissolved)
Electroplating
Object to be plated = cathode
Plating metal = anode
Solution contains ions of the plating metal
Smooth, steady current gives an even coating
Rate law: Rate = k[A]^m[B]^n
Keq: [products]^coefficients / [reactants]^coefficients
Ksp: [cation]^x[anion]^y
pH: –log[H⁺]
pOH: –log[OH⁻]
Kw: [H⁺][OH⁻] = 1.0 x 10⁻¹⁴
Titration: M(A)V(A) = M(B)V(B) (account for mole ratios)
Cell potential: E°(cell) = E°(reduction) – E°(oxidation)
Catalytic converters in cars use heterogeneous catalysts to speed up reactions that convert toxic exhaust gases (CO, NOₓ) into less harmful substances (CO₂, N₂).
Le Chatelier's principle governs the industrial Haber process for ammonia: high pressure and moderate temperature maximise yield.
Buffer systems keep blood pH at 7.4. The carbonic acid-bicarbonate buffer is one of the most important systems in human physiology.
Galvanising (zinc-coating) steel is the standard method for protecting exposed structural steel in construction and automotive manufacturing.
Electrolysis of brine is a cornerstone industrial process, producing chlorine (for water treatment and PVC), hydrogen (fuel), and sodium hydroxide (used in soap, paper, and aluminium production).
Students often think that a large Keq means the reaction happened quickly. Keq says nothing about speed; it describes the position of equilibrium. A reaction can be strongly product-favoured but very slow without a catalyst.
The equivalence point of a titration is not always at pH 7. That only holds for a strong acid + strong base. A weak acid titrated with a strong base has an equivalence point above 7.
Students frequently assume that a catalyst changes the equilibrium position. It does not. A catalyst speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but at the same position.
In redox, students sometimes confuse the oxidising agent with the species being oxidised. The oxidising agent is the one being reduced (it causes oxidation in the other species).
For cell potential, students sometimes add E° values incorrectly. E°(cell) = E°(cathode) – E°(anode), or equivalently E°(reduction) – E°(oxidation). Do not multiply E° by coefficients.
⚠️ Determining reaction order from experimental data (doubling concentration and observing rate change) is a classic exam question.
⚠️ Le Chatelier problems are tested in multiple formats: predicting the direction of shift, predicting whether Keq changes, and predicting concentration changes.
⚠️ pH calculations for strong acids, strong bases, and weak acids (using Ka) are core exam material. Know all three paths.
⚠️ Titration calculations and interpreting titration curves (finding the equivalence point, reading pH) are heavily tested.
⚠️ Balancing redox equations using the oxidation-number method is a multi-step process that appears frequently. Practise until it is systematic.
⚠️ Cell potential calculations (E°cell = E°reduction – E°oxidation) and knowing which half-reaction is the anode vs cathode are standard final exam questions.
⚠️ Know the key industrial applications: electrolysis of molten NaCl, electrolysis of brine, electroplating setup (what is the cathode, what is the anode, what is in solution).
True or False: A catalyst changes the equilibrium constant of a reaction.
Fill in the blank: If Keq > 1, the __________ are favoured at equilibrium.
True or False: The pH at the equivalence point of a weak acid-strong base titration is exactly 7.
Fill in the blank: In electrochemistry, oxidation occurs at the __________.
True or False: Increasing the temperature of an exothermic reaction at equilibrium shifts it toward the products.
Q: If doubling the concentration of reactant A causes the rate to quadruple, what is the reaction order with respect to A?
A: 2^m = 4, so m = 2. The reaction is second order in A.
Q: For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = –92 kJ, what happens to the equilibrium if the temperature is increased?
A: The reaction is exothermic. Increasing temperature favours the endothermic (reverse) direction. Equilibrium shifts to the left, producing more N₂ and H₂.
Q: Calculate the pH of a 0.010 M HCl solution.
A: HCl is a strong acid and ionises completely. [H⁺] = 0.010 M. pH = –log(0.010) = 2.0.
Q: What is E°cell for a voltaic cell with Zn²⁺/Zn (E° = –0.76 V) and Cu²⁺/Cu (E° = +0.34 V)?
A: Zn is oxidised (anode), Cu²⁺ is reduced (cathode). E°cell = 0.34 – (–0.76) = +1.10 V.
Q: In electroplating a fork with silver, which electrode is the fork?
A: The cathode. Silver ions in solution are reduced to silver metal on the fork's surface.
Q: Identify the oxidising agent in the reaction: Zn + CuSO₄ → ZnSO₄ + Cu.
A: Cu²⁺ (from CuSO₄) is the oxidising agent. It gains electrons (is reduced), causing Zn to be oxidised.
Kinetics (Ch. 17) and equilibrium (Ch. 18) are complementary: kinetics describes how fast a system reaches equilibrium, while equilibrium describes where it ends up.
Acid-base equilibrium (Ch. 19) is a specific application of the general equilibrium concepts from Ch. 18, using Ka, Kb, and Kw.
Gibbs free energy (Ch. 16) connects to electrochemistry: ΔG° = –nFE°cell, where n = moles of electrons transferred and F = Faraday's constant.
Redox balancing (Ch. 20) is a prerequisite for understanding electrochemical cell setups (Ch. 21).
Thermochemistry (Ch. 16) and kinetics (Ch. 17) together answer two distinct questions about any reaction: will it happen (thermodynamics) and how fast will it happen (kinetics).
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