Unit 9 asks two questions about any reaction: how fast does it go, and how far does it go before it settles? By the end you should be able to explain rate changes with collision theory, write and calculate equilibrium constants, and predict how an equilibrium shifts when you change concentration, pressure, or temperature.
Read the overview first, then work through each topic's notes and checkpoint question before testing yourself on the common mistakes and key terms.
This unit links two ideas that are easy to mix up. Kinetics tells you how quickly a reaction happens, while equilibrium tells you where the reaction ends up. A reaction can be fast and barely make product, or slow and almost completely convert to product.
Honors Chemistry Unit 9 covers chemical kinetics and chemical equilibrium. You learn the factors that change reaction rate (concentration, temperature, surface area, pressure for gases, and catalysts), explained through collision theory and activation energy, along with simple rate laws like . You then study reversible reactions that reach a dynamic equilibrium, write and calculate equilibrium constants ( and ), and use Le Chatelier's principle to predict shifts. The Haber process for making ammonia shows both halves of the unit working together.
Collision theory says most collisions do nothing. A collision has to supply at least the activation energy, , and line up the right atoms. Raising temperature, concentration, or surface area increases effective collisions, and a catalyst lowers so more collisions succeed.
At equilibrium the forward and reverse reactions still run at equal rates, so concentrations stay constant. The equilibrium constant compares products to reactants at that point. A large means products are favored and a small means reactants are favored.
Le Chatelier's principle predicts which way the system moves after you add or remove a substance, change the volume of a gas mixture, or heat or cool it. Only a temperature change alters the value of . The other stresses move the position of equilibrium while stays the same.
Kinetics decides how fast you reach equilibrium, and equilibrium decides what the mixture looks like once you get there. A catalyst gets you there faster but doesn't change the final mixture. Changing temperature affects both, which is why chemists often settle on a compromise temperature.
Reaction rate is how fast reactants are used up or products form, usually in molarity per second (M/s).
A reversible reaction can run both ways.
Le Chatelier's principle says that when you stress a system at equilibrium, it shifts in the direction that partially counteracts the stress.
Reaction rate is how fast reactants are used up or products form, usually in molarity per second (M/s). Collision theory explains every factor. Particles must collide with at least the activation energy, , and the correct orientation.
Higher concentration packs more particles into the same space, so collisions happen more often. Higher temperature makes particles move faster, and more importantly a bigger fraction of collisions now have energy above . A common rule of thumb is that a 10°C rise roughly doubles the rate of many reactions, but that is only an approximation. Grinding a solid into powder increases surface area, so more particles are exposed. Squeezing a gas into a smaller volume raises its concentration. A catalyst provides a different pathway with a lower and is not used up.
A rate law shows how concentration affects rate, . The orders and come from experiments, not from the coefficients in the balanced equation.
Worked example. Trial 1 has M, M, and a rate of M/s. Doubling to 0.20 M gives M/s, which is 4 times faster, so A is second order (). Doubling instead gives M/s, so B is first order. The rate law is , and .
Check yourself. In a reaction with , what happens to the rate if you triple and keep the same? (Answer: The rate increases by a factor of .)
| Factor | What changes | Why the rate goes up |
|---|---|---|
| Higher concentration | More particles per liter | Collisions happen more often |
| Higher temperature | Particles move faster | More collisions per second, and a larger fraction have energy above $E_a$ |
| More surface area | Solid broken into smaller pieces | More particles exposed to collisions |
| Higher pressure (gases only) | Same gas squeezed into less volume | Gas concentration rises, so collisions happen more often |
| Catalyst | New reaction pathway | Lower $E_a$, so more collisions succeed |
A reversible reaction can run both ways. If you start with reactants and no products, the reverse rate is initially zero, while the forward reaction makes products. As products build up, the reverse rate increases. The forward and reverse rates eventually become equal if the system reaches equilibrium. At that point reactions are still happening, but concentrations stay constant. Equilibrium needs a closed system, which keeps reactants and products from entering or leaving.
For , the equilibrium constant is . Coefficients become exponents. uses partial pressures of gases instead of molarities. Leave out pure solids and pure liquids, because their concentrations don't change.
If , products are favored at equilibrium. If , reactants are favored. For a given reaction, depends only on temperature, not on how much you started with.
Worked example. For , a mixture at equilibrium has M, M, and M. Write . Then . Since , products are somewhat favored at this temperature.
Check yourself. Write the expression for . (Answer: , because the two solids are left out.)
Le Chatelier's principle says that when you stress a system at equilibrium, it shifts in the direction that partially counteracts the stress. It never fully undoes the change, but it moves to a new equilibrium position.
For concentration, the system shifts away from what you add and toward what you remove. Adding to the ammonia equilibrium pushes it toward .
For temperature, treat heat like a chemical. In an exothermic reaction, write heat as a product, so heating shifts toward reactants and cooling shifts toward products. In an endothermic reaction, heat is a reactant, so the pattern flips. Temperature is the only stress that changes the value of .
For pressure, only gases count. Decreasing volume (raising pressure) shifts toward the side with fewer moles of gas. If both sides have the same number of gas moles, nothing shifts. Adding an inert gas like argon at constant volume doesn't shift anything, because the partial pressures of the reacting gases don't change.
A catalyst never shifts equilibrium. It speeds up the forward and reverse reactions equally.
The Haber process, with kJ, has 4 moles of gas on the left and 2 on the right. Factories use high pressure (around 200 atm) to favor ammonia. Low temperature would also favor ammonia but would be too slow, so they compromise near 450°C and add an iron catalyst.
Check yourself. For , which way does the equilibrium shift if you decrease the volume? (Answer: Toward , because the right side has 2 moles of gas and the left side has 3.)
| Stress | Direction of shift | Does K change? |
|---|---|---|
| Add a reactant | Toward products | No |
| Remove a product | Toward products | No |
| Decrease volume of a gas mixture | Toward fewer moles of gas | No |
| Raise temperature, exothermic reaction | Toward reactants | Yes, K decreases |
| Raise temperature, endothermic reaction | Toward products | Yes, K increases |
| Add a catalyst | No shift | No |
| Term | Definition |
|---|---|
| Reaction rate | The change in concentration of a reactant or product per unit time, often in M/s. |
| Collision theory | The model that says reactions happen when particles collide with enough energy and the correct orientation. |
| Activation energy | The minimum collision energy needed to start a reaction. |
| Catalyst | A substance that lowers activation energy and speeds up a reaction without being used up. |
| Rate law | An equation like $Rate = k[A]^m[B]^n$ that links reaction rate to reactant concentrations. |
| Reversible reaction | A reaction where products can react to re-form the original reactants. |
| Chemical equilibrium | The state where forward and reverse reaction rates are equal and concentrations stay constant. |
| Equilibrium constant (K) | The ratio of product to reactant concentrations or pressures at equilibrium, each raised to its coefficient. |
| Le Chatelier's principle | The rule that a system at equilibrium shifts to partially counteract any stress placed on it. |
| Haber process | The industrial reaction $N_2 + 3H_2 \rightleftharpoons 2NH_3$, run at high pressure with an iron catalyst to make ammonia. |
Equilibrium means equal rates, not equal concentrations. A reaction with strongly favors reactants, but you need the balanced reaction and the starting amounts to calculate the actual equilibrium mixture. Check whether the rates are equal and the concentrations are constant, not whether the amounts match.
A catalyst lowers the activation energy for both the forward and reverse reactions, so it speeds both up by the same factor. You reach equilibrium sooner, but the final mixture and the value of stay the same.
In , the exponents come from experimental data, such as comparing initial rates when one concentration doubles. Coefficients do become exponents in the equilibrium expression , which is why the two get mixed up.
Pure solids and pure liquids have constant concentrations, so they stay out of expressions. Adding more of a pure solid, like extra , to an existing equilibrium doesn't shift it as long as some of that solid is already present. Adding a liquid solvent such as water is a different case, because it dilutes the dissolved substances and can shift an aqueous equilibrium. Changing the concentrations or partial pressures of reacting gases and dissolved substances can shift equilibrium, and a temperature change shifts it and changes the value of .