Difficulty: Intermediate | Prerequisites: Balancing chemical equations, basic reaction types, Lewis structures
These two chapters introduce two foundational ideas. First, chemical equilibrium: reactions do not always run to completion; they reach a balance point where forward and reverse reactions proceed at equal rates. Second, acids and bases: two frameworks (Arrhenius and Brønsted-Lowry) for classifying substances that produce or transfer protons. The exam coverage here is conceptual and definitional, with no equilibrium calculations required. If you can balance equations and draw Lewis structures, you are prepared.
Equilibrium is the state where the rates of the forward and reverse reactions are equal, so concentrations stop changing. The equilibrium constant (K) tells you whether products or reactants are favoured. Arrhenius acids produce H⁺ in water, Arrhenius bases produce OH⁻. Brønsted-Lowry broadens this: acids donate protons, bases accept them. Every Brønsted acid-base reaction produces a conjugate pair.
Chemical equilibrium
The state in which the rate of the forward reaction equals the rate of the reverse reaction. Concentrations of reactants and products remain constant (but not necessarily equal). Think of it as a tug-of-war where both sides pull with equal force: nobody moves, but the rope is still under tension.
Equilibrium expression
A mathematical ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients. Written from the balanced equation.
Equilibrium constant (K)
The numerical value of the equilibrium expression at a given temperature. A large K (much greater than 1) means products are favoured at equilibrium. A small K (much less than 1) means reactants are favoured. K = 1 means roughly equal amounts of both.
Arrhenius acid
A substance that produces hydrogen ions (H⁺) when dissolved in water. Example: HCl dissolves in water to give H⁺ and Cl⁻.
Arrhenius base
A substance that produces hydroxide ions (OH⁻) when dissolved in water. Example: NaOH dissolves to give Na⁺ and OH⁻.
Brønsted-Lowry acid
A proton (H⁺) donor. This is a broader definition than Arrhenius because it does not require water as the solvent.
Brønsted-Lowry base
A proton (H⁺) acceptor. Again, broader than the Arrhenius definition.
Conjugate acid-base pair
Two species that differ by exactly one proton. When an acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid. In simple terms, they are dance partners: what one gives up, the other takes on.
Conjugate acid
The species formed when a Brønsted-Lowry base accepts a proton. For example, when NH₃ accepts H⁺, it becomes NH₄⁺ (the conjugate acid).
Conjugate base
The species formed when a Brønsted-Lowry acid donates a proton. For example, when HCl donates H⁺, it becomes Cl⁻ (the conjugate base).
Many reactions are reversible. At equilibrium, the forward and reverse reactions occur at the same rate.
Concentrations of reactants and products are constant at equilibrium, but they are not necessarily equal to each other.
Molecular-level images of equilibrium: look for a stable ratio of reactant and product particles that does not change from one snapshot to the next. If the counts are still shifting, the system has not reached equilibrium yet.
Graphs of equilibrium: concentrations change initially, then level off to constant values. The point where all lines flatten is equilibrium.
For a general reaction: aA + bB ⇌ cC + dD
The equilibrium expression is:
K = [C]^c × [D]^d / [A]^a × [B]^b
Products go in the numerator. Reactants go in the denominator.
Each concentration is raised to the power of its coefficient in the balanced equation.
Pure solids and pure liquids do not appear in the expression (their concentrations are effectively constant).
K is a number. Its value tells you where the equilibrium position lies.
K >> 1: products are strongly favoured. The reaction goes nearly to completion.
K << 1: reactants are strongly favoured. Very little product forms.
K ≈ 1: significant amounts of both reactants and products are present.
K depends on temperature. Change the temperature and K changes.
Arrhenius acid: produces H⁺ (or H₃O⁺) in aqueous solution. Examples: HCl, HNO₃, H₂SO₄.
Arrhenius base: produces OH⁻ in aqueous solution. Examples: NaOH, KOH, Ca(OH)₂.
Limitation: the Arrhenius model only works in water. It cannot explain why NH₃ acts as a base (NH₃ does not contain OH⁻).
Brønsted-Lowry acid: a proton donor.
Brønsted-Lowry base: a proton acceptor.
This definition works in any solvent, not just water.
Example: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. Here, H₂O donates a proton to NH₃. Water is the Brønsted-Lowry acid; NH₃ is the Brønsted-Lowry base.
In every Brønsted-Lowry reaction, two conjugate pairs form.
Example: HCl + H₂O → H₃O⁺ + Cl⁻
Pair 1: HCl (acid) and Cl⁻ (conjugate base). They differ by one H⁺.
Pair 2: H₂O (base) and H₃O⁺ (conjugate acid). They differ by one H⁺.
To find a conjugate base: remove one H⁺ from the acid and reduce the charge by 1.
To find a conjugate acid: add one H⁺ to the base and increase the charge by 1.
You should be able to draw Lewis structures of an acid/base and its conjugate partner, showing where the proton was lost or gained.
General equilibrium expression
For aA + bB ⇌ cC + dD:
K = \frac{[C]^c [D]^d}{[A]^a [B]^b}Remember: products over reactants, each raised to its stoichiometric coefficient. Pure solids and liquids are omitted.
Conjugate pair identification (no formula, but a systematic method)
Write the balanced acid-base reaction.
Identify which species lost a proton (the acid) and which gained one (the base).
The acid minus one H⁺ is its conjugate base.
The base plus one H⁺ is its conjugate acid.
Draw Lewis structures of each pair side by side, showing the lone pair that accepted the proton or the bond that released it.
Students often think equilibrium means the concentrations of reactants and products are equal. They are constant, not equal. A reaction with K = 1000 has far more product than reactant at equilibrium.
A common error is including pure solids or pure liquids in the equilibrium expression. Only aqueous species and gases appear.
Students sometimes confuse Arrhenius and Brønsted-Lowry definitions. Every Arrhenius acid is a Brønsted-Lowry acid, but not every Brønsted-Lowry base is an Arrhenius base (e.g., NH₃ is a Brønsted-Lowry base but not an Arrhenius base because it does not directly produce OH⁻).
Forgetting that water can act as either an acid or a base (it is amphoteric/amphiprotic). In one reaction it donates a proton; in another it accepts one.
⚠️ Be able to write the equilibrium expression from a balanced equation. Know what to include and what to exclude.
⚠️ Interpret the meaning of K: large K vs. small K vs. K near 1.
⚠️ Identify whether a given molecular-level image or graph shows a system at equilibrium or still approaching it.
⚠️ Define Arrhenius and Brønsted-Lowry acids/bases. Know the difference and the overlap.
⚠️ Identify both conjugate pairs in a Brønsted-Lowry reaction. Be able to draw the Lewis structures.
True or False: At equilibrium, the forward reaction has stopped. Answer: False. Both forward and reverse reactions continue, but at equal rates.
Fill in the blank: A large equilibrium constant (K >> 1) means ______ are favoured. Answer: Products.
True or False: NH₃ is an Arrhenius base. Answer: False. NH₃ is a Brønsted-Lowry base (it accepts H⁺) but not an Arrhenius base (it does not directly produce OH⁻ in its formula).
Fill in the blank: The conjugate base of H₂O is ______. Answer: OH⁻.
True or False: Pure water appears in the equilibrium expression for an aqueous reaction. Answer: False. Pure liquids are excluded from K expressions.
Q: Write the equilibrium expression for the reaction: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
A: K = [NH₃]² / ([N₂] × [H₂]³). Products in the numerator, reactants in the denominator, each raised to its coefficient.
Q: For the reaction above, if K = 6.0 × 10⁵ at a given temperature, are products or reactants favoured?
A: Products are strongly favoured. K is much greater than 1, so at equilibrium the concentration of NH₃ is far larger than those of N₂ and H₂.
Q: Identify the two conjugate acid-base pairs in: HF + H₂O ⇌ H₃O⁺ + F⁻
A: Pair 1: HF (acid) and F⁻ (conjugate base). Pair 2: H₂O (base) and H₃O⁺ (conjugate acid).
Q: Is NH₃ an Arrhenius base, a Brønsted-Lowry base, or both? Explain.
A: NH₃ is a Brønsted-Lowry base (it accepts a proton from water to form NH₄⁺ and OH⁻). It is not an Arrhenius base because its formula does not contain OH⁻. The OH⁻ forms as a product of the reaction with water, not from NH₃ itself.
Q: What is the conjugate acid of SO₄²⁻?
A: HSO₄⁻. Add one H⁺ to SO₄²⁻ and increase the charge by 1 (from -2 to -1).
Equilibrium connects back to solutions (Ch. 12): the dissolution of a solid in a solvent reaches equilibrium when the solution is saturated, and vapour pressure above a liquid is an equilibrium between evaporation and condensation.
Acid-base chemistry extends into buffer systems, pH calculations, and titrations in later chapters. The conjugate pair concept is the foundation for understanding how buffers resist pH changes.
The equilibrium constant reappears in solubility product (Ksp) and acid dissociation constant (Ka) calculations later in the course.
Equilibrium, chemical equilibrium, dynamic equilibrium, equilibrium expression, equilibrium constant, K, Keq, products favoured, reactants favoured, reversible reaction, Arrhenius acid, Arrhenius base, Brønsted-Lowry acid, Brønsted-Lowry base, proton donor, proton acceptor, conjugate acid, conjugate base, conjugate pair, amphoteric, amphiprotic, H₃O⁺, hydronium, hydroxide, OH⁻, Lewis structure of acids, CHM 11200, Purdue general chemistry, Chapter 14, Chapter 15