Unit 4 Chemistry- Comprehensive Review

What Unit 4 Chemistry Actually Covers

Unit 4 Chemistry is the section most students struggle with. Not because it's harder than other units, but because it requires you to think in two directions at once—forward with the reaction and backward toward equilibrium. If you're taking AP Chemistry, IB, or a college-level general chemistry course, this is where things get real.

The core topics are chemical kinetics and chemical equilibrium. Some curricula bundle thermochemistry here. Others throw in acid-base equilibrium. Know your specific course outline before you deep-dive into anything.

Chemical Kinetics: How Fast Things Move

Kinetics answers one question: what determines the speed of a reaction?

You're not asking if a reaction will happen. You're asking how quickly it happens and why.

Reaction Rates

The rate of a reaction is the change in concentration of a reactant or product over time. Units are usually M/s (molarity per second).

For the reaction: aA → bB

Rate = -(1/a)(Δ[A]/Δt) = (1/b)(Δ[B]/Δt)

The negative sign for reactants means concentration is decreasing. Don't forget it.

Rate Laws

The rate law shows how rate depends on concentration:

Rate = k[A]m[B]n

Reaction orders tell you what happens to the rate when you double a concentration:

The overall order is m + n. A common mistake is assuming the order equals the stoichiometric coefficient. It doesn't. Ever.

Integrated Rate Laws

For zero-order reactions:

[A]t = -kt + [A]0

A plot of [A] vs. time gives a straight line. Slope = -k.

For first-order reactions:

ln[A]t = -kt + ln[A]0

A plot of ln[A] vs. time gives a straight line. Slope = -k.

For second-order reactions:

1/[A]t = kt + 1/[A]0

A plot of 1/[A] vs. time gives a straight line. Slope = k.

Half-life (t1/2) is the time for concentration to drop to half. For first-order reactions, it's independent of starting concentration:

t1/2 = 0.693/k

Reaction Mechanisms

Reactions don't happen in one step. They go through a series of elementary steps called the reaction mechanism.

Each step has its own rate law. The overall reaction is the sum of all steps.

The rate-determining step is the slowest step. It controls the overall rate law. Intermediates appear in mechanism steps but cancel out in the overall reaction.

If step 1 is slow and step 2 is fast:

Rate = k[reactants in slow step]

Collision Theory and Activation Energy

For a reaction to occur:

  • Particles must collide
  • Collision must have enough energy (≥ activation energy)
  • Collision must have the correct orientation

Activation energy (Ea) is the energy barrier. Higher Ea = slower reaction. Temperature affects this exponentially, which is why heating sometimes makes reactions explode.

The Arrhenius equation:

k = Ae-Ea/RT

A is the frequency factor. It accounts for collision frequency and orientation.

Catalysts

Catalysts speed up reactions without being consumed. They lower the activation energy by providing an alternative pathway.

Homogeneous catalysts are in the same phase as reactants. Heterogeneous catalysts are in a different phase (like solid platinum in a gas reaction).

Enzymes are biological catalysts. They work the same way—lower Ea, faster rate.

Chemical Equilibrium: When Opposites Balance

Equilibrium is when the forward and reverse rates are equal. Concentrations stop changing macroscopically. This doesn't mean the reaction stops—it means both directions are happening at the same rate.

The Equilibrium Constant (K)

For: aA + bB ⇌ cC + dD

Keq = [C]c[D]d / [A]a[B]b

Only gases and aqueous species go in the expression. Solids and liquids are omitted (their concentrations are constant).

What K tells you:

  • K >> 1: Products dominate at equilibrium
  • K << 1: Reactants dominate at equilibrium
  • K ≈ 1: Significant amounts of both

K is temperature-dependent. Changing concentration, pressure, or adding catalysts doesn't change K.

Kp and Kc

Kc uses molarity (mol/L). Kp uses partial pressures.

Kp = Kc(RT)Δn

Δn = (moles of gaseous products) - (moles of gaseous reactants)

Le Chatelier's Principle

When you disturb an equilibrium, it shifts to counteract the disturbance.

Stress the system → it adjusts → new equilibrium forms

Changes and their effects:

  • Add reactant: equilibrium shifts right (toward products)
  • Add product: equilibrium shifts left (toward reactants)
  • Increase pressure: equilibrium shifts toward fewer moles of gas
  • Decrease pressure: equilibrium shifts toward more moles of gas
  • Increase temperature: equilibrium shifts toward endothermic direction
  • Decrease temperature: equilibrium shifts toward exothermic direction
  • Add catalyst: no shift (just reaches equilibrium faster)

For pressure changes, only gaseous species matter. A reaction with equal moles of gas on both sides doesn't shift with pressure changes.

The Reaction Quotient (Q)

Q tells you where you are relative to equilibrium:

Q < K: Reaction shifts right (toward products)

Q > K: Reaction shifts left (toward reactants)

Q = K: Already at equilibrium

Calculate Q the same way you calculate K, using current concentrations.

ICE Tables

ICE = Initial, Change, Equilibrium. Use these to find equilibrium concentrations.

Example: Start with 2.0 M A and 0 M B. K = 4.0 for A ⇌ 2B

A B
Initial 2.0 0
Change -x +2x
Equilibrium 2.0 - x 2x

K = [B]2/[A] = (2x)2/(2.0 - x) = 4.0

Solve: 4x2 = 4.0(2.0 - x) → x = 1.0

Equilibrium: [A] = 1.0 M, [B] = 2.0 M

Check: Q = (2)2/1 = 4 = K ✓

Common Mistakes That Cost You Points

  • Confusing rate with equilibrium. Fast doesn't mean favorable. Slow reactions can still go nearly to completion.
  • Using stoichiometric coefficients as reaction orders. Experimental data only. Not negotiable.
  • Forgetting to convert to consistent units. Partial pressures and molarity don't mix in the same expression.
  • Omitting solids and liquids from equilibrium expressions. Only gases and aqueous species count.
  • Thinking catalysts affect K. They don't. They only speed up the approach to equilibrium.
  • Not writing units on rate constant calculations. k units depend on overall reaction order.

How to Actually Solve Unit 4 Problems

For Kinetics Problems

  1. Identify what information is given. Rate data? Time? Concentration?
  2. Determine the rate law by comparing experiments where one concentration changes while others stay constant.
  3. Find the order for each reactant from the data.
  4. Calculate k using the rate law and one set of data.
  5. For integrated rate laws, plot the appropriate variable vs. time. The straight line confirms the order.
  6. Calculate half-life or find concentration at a given time using the integrated equation.

For Equilibrium Problems

  1. Write the balanced equation. Get coefficients right.
  2. Write the K expression. Only gases and aqueous species.
  3. Set up the ICE table. Use x for the unknown change.
  4. Substitute equilibrium values into K expression.
  5. Solve for x. If K is small, you can often approximate (2.0 - x ≈ 2.0). If K is large, you might need the quadratic formula.
  6. Check your approximation by calculating Q with your answer. Should equal K within reason.

Quick Reference: Key Equations

Topic Equation
Rate Law Rate = k[A]m[B]n
First-order half-life t1/2 = 0.693/k
First-order integrated ln[A]t = -kt + ln[A]0
Equilibrium constant K = [products]/[reactants]
Kp to Kc Kp = Kc(RT)Δn
Arrhenius k = Ae-Ea/RT

What to Study Before the Exam

Work through problems until you can do them without checking the book. The formulas are useless if you don't know when to use them.

Focus on:

  • Determining rate laws from experimental data
  • Setting up and solving ICE tables
  • Interpreting graphs for integrated rate laws
  • Le Chatelier shifts with multiple changes at once
  • Connecting kinetics to equilibrium (they're related through the reaction pathway)

Unit 4 isn't about memorizing. It's about understanding how reactions behave and being able to predict outcomes. Get the concepts right and the math follows.