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Chemistry Unit 6: Kinetics and equilibrium

Kinetics is how fast a reaction goes; equilibrium is where it settles. Le Chatelier's principle predicts how equilibrium responds to stress.

Covers: Reaction rates and collision theory, catalysts, rate laws, equilibrium constant K, Le Chatelier's principle

Study guide

Chemical kinetics studies reaction rates. Collision theory says particles must collide with enough energy (the activation energy) and the right orientation to react. Rates increase with higher concentration or pressure (more collisions), higher temperature (more energetic collisions), greater surface area, and catalysts. A catalyst speeds a reaction by providing a pathway with lower activation energy and isn't used up; enzymes are biological catalysts. Energy diagrams show reactants, products, the activation energy hill, and how a catalyst lowers it.

A rate law expresses how rate depends on concentration: rate = k[A]ᵐ[B]ⁿ. The orders m and n must be found experimentally, usually by comparing trials where one concentration changes. Doubling a reactant's concentration doubles the rate if it's first order and quadruples it if second order. Reaction mechanisms describe the step-by-step path, and the slowest step limits the overall rate.

Many reactions are reversible. At dynamic equilibrium, the forward and reverse rates are equal, so concentrations stay constant even though both reactions continue. For aA + bB ⇌ cC + dD, the equilibrium constant is K = [C]ᶜ[D]ᵈ/[A]ᵃ[B]ᵇ, leaving out pure solids and liquids. K much greater than 1 means products are favored; K much less than 1 means reactants are favored. Comparing the reaction quotient Q to K predicts which direction a reaction will shift.

Le Chatelier's principle says a system at equilibrium shifts to counteract a stress. Adding a reactant shifts it toward products. Increasing pressure shifts it toward fewer gas moles. Raising temperature shifts it in the endothermic direction, and only temperature changes K. A catalyst doesn't shift equilibrium; it only helps reach it faster.

For a slightly soluble salt such as Mg(OH)₂ ⇌ Mg²⁺ + 2OH⁻, Ksp = [Mg²⁺][OH⁻]², and the solid is left out. If Q > Ksp, a precipitate forms; if Q < Ksp, more solid can dissolve. Adding an ion that's already in the equilibrium (the common-ion effect) lowers solubility. To solve an equilibrium problem, make an ICE table that lists the Initial concentrations, the Change (in amounts set by the coefficients, such as −x and +2x), and the Equilibrium values. Substitute the equilibrium values into the K expression and solve for x. If K is very small, x can usually be ignored where it's added to or subtracted from the initial concentration.

For gas-phase equilibria, K can be written with partial pressures instead of concentrations. This form is Kp, and pure solids and liquids are still left out. For example, for N₂ + 3H₂ ⇌ 2NH₃, Kp = P(NH₃)²/(P(N₂)·P(H₂)³). For a first-order reaction, ln[A] = −kt + ln[A]₀, so a graph of ln[A] against time is a straight line with slope −k. Its half-life, t½ = 0.693/k, is constant no matter the concentration. For a zero-order reaction a graph of [A] against time is linear, and for a second-order reaction a graph of 1/[A] against time is linear.

In a reaction mechanism, an intermediate is made in one step and used up in a later step. A catalyst is used up in an early step and made again later, so it is present at the start and the end. Neither one appears in the overall balanced equation.

Key ideas

• Rates rise with concentration, temperature, surface area and catalysts.
• Catalysts lower activation energy without being used up.
• At equilibrium, forward and reverse rates are equal; concentrations stay constant.
• K = [products]/[reactants], each raised to its coefficient; solids and liquids are left out.
• Le Chatelier: a stressed system shifts to partly undo the stress.

Formulas and rules

Common mistakes

• Thinking a catalyst shifts equilibrium. It only speeds both directions.
• Including pure solids or liquids in K.

Words to know

Catalyst

Speeds a reaction by lowering activation energy; not used up.

Dynamic equilibrium

Forward and reverse reaction rates are equal.

Le Chatelier's principle

A system at equilibrium shifts to counteract a change.

Large K (K ≫ 1)

Products are favored at equilibrium.

Temperature and reaction rate

Higher temperature → faster reaction.

Equilibrium constant (K)

For aA + bB ⇌ cC + dD: K = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ

Collision theory

Particles must collide with enough energy and the right orientation to react.

Activation energy

The minimum energy needed for a reaction.

Rate law

rate = k[A]ᵐ[B]ⁿ

Reaction order

The exponent on a concentration in the rate law, found experimentally.

Rate-determining step

The slowest step of a mechanism.

Reaction quotient (Q)

Same form as K but with current concentrations; compare to K.

Effect of pressure on equilibrium

Higher pressure shifts toward fewer gas moles.

What changes K

Only temperature.

Sample practice questions

5 more questions like these, from level 1 to 5, are in the practice set.