Rates & Equilibrium — Print · Physical Science: Chemistry, Grade 11

Reaction Rates & Kinetics

Learning goals

• I can use collision theory to explain what must happen for particles to react.
• I can predict how concentration, temperature, surface area, and catalysts change the rate of a reaction.
• I can calculate an average reaction rate from the change in an amount over time.

Readings for this unit

• The catalyst can change the journey without changing the destination ratio

Some reactions are over in a flash — a stick of dynamite, a spark in a gas stove — while others crawl along for years, like iron slowly rusting or rock weathering into soil. The branch of chemistry that studies how fast reactions go, and what speeds them up or slows them down, is called kinetics. We measure a reaction rate as a change in amount over a change in time: how much reactant disappears, or how much product appears, each second.

To understand what sets that speed, chemists use collision theory. Its central claim is simple: particles can only react when they actually collide with one another — and not just any collision will do. A collision only leads to a reaction if the particles hit with at least the activation energy (enough energy to break the starting bonds) AND with the correct orientation (the reactive parts lined up). A collision that meets both conditions is an effective collision. Most collisions are gentle glancing blows that bounce apart unchanged, so anything that makes effective collisions more frequent will speed the reaction up.

Changing concentration can change reaction rate by affecting encounters and the measured rate law. For some elementary steps the collision picture gives a direct relationship, but real multistep reactions can have different concentration dependence. A rate law is determined experimentally; the balanced overall equation does not generally supply its exponents.

Higher temperature often increases the fraction of particles capable of crossing the activation barrier. The rate depends on a temperature-sensitive rate constant and the mechanism; doubling per 10°C is at most a rough heuristic for some ranges, not a law. Cooling usually slows many spoilage processes, but actual biological and chemical responses need data.

The third factor is surface area. When a solid reacts with a liquid or gas, only the particles on its exposed surface can be struck. Grind that solid into a powder and you expose an enormous amount of new surface, so vastly more collisions can occur at once. A single log is hard to light, but the same wood as fine sawdust can ignite explosively — the same principle behind the dangerous dust explosions that can level a flour mill or grain silo.

The fourth factor is a catalyst — a substance that speeds up a reaction without being consumed. A catalyst works by offering an alternate reaction pathway with a lower activation energy, so a far greater share of ordinary collisions now have enough energy to succeed. Because it is regenerated at the end, a tiny amount can process mountains of reactant. Your body relies on protein catalysts called enzymes for nearly every reaction of life, and the catalytic converter in a car speeds the reactions that scrub toxic gases from its exhaust. Crucially, a catalyst changes only the speed — it does not change ΔH or how much product a reaction can ultimately form.

Model check: Concentration effects depend on the rate law; catalysts do not change equilibrium yield. The kinetic pathway and rate change; overall reaction ΔH and equilibrium K stay the same at fixed conditions.

Average Reaction Rate

average rate = Δamount / Δtime

the amount can be a concentration (mol/L), a volume of gas, or a number of moles; time is usually measured in seconds.

Worked examples

Avoid these mistakes

• Concentration effects depend on the rate law; catalysts do not change equilibrium yield.

Words to know

reaction rate

How fast reactants are used up or products form, often measured as a change in concentration or amount per unit of time.

collision theory

The model that reactions happen only when particles collide with enough energy and the correct orientation.

effective collision

A collision that actually leads to a reaction because it has sufficient energy and proper orientation.

catalyst

A species that changes the reaction pathway and rate without being consumed overall; it does not change the same reaction’s equilibrium constant at fixed conditions.

surface area

The total exposed area of a solid; breaking a solid into smaller pieces increases it and speeds up reactions.

inhibitor

A substance that slows a reaction down, the opposite of a catalyst.

Rate laws and the Arrhenius relationship

A measured rate law can have form rate=k[A]^m[B]^n. Doubling [A] multiplies rate by 2^m when other conditions stay fixed; m and n are measured, not generally copied from overall coefficients. For a first-order dependence, doubling A doubles rate. The Arrhenius form k=Aexp(−Ea/RT) relates a model rate constant to kelvin temperature; use Ea and R in consistent energy units. A plot of ln k against 1/T has slope −Ea/R when the model applies. A catalyst can change Ea and the pathway without changing reaction ΔH or equilibrium K.

In the real world

Food is refrigerated to slow the reactions of spoilage, while a car's catalytic converter speeds up the reactions that turn toxic exhaust gases into safer ones before they leave the tailpipe.

Exit check

What changes most directly when a catalyst provides a lower barrier?

Chemical Equilibrium

Learning goals

• I can explain how a reversible reaction reaches a dynamic equilibrium where forward and reverse rates are equal.
• I can use Le Châtelier's principle to predict how a change in concentration, temperature, or pressure shifts an equilibrium.
• I can apply equilibrium reasoning to a real industrial system such as the Haber process.

Readings for this unit

• The catalyst can change the journey without changing the destination ratio

A reversible reaction can proceed in both directions. Its rate and starting composition determine how it approaches equilibrium; a sample need not begin with only reactants. Under fixed conditions in an appropriate closed system, forward and reverse rates can become equal while microscopic events continue. The double arrow denotes reversibility rather than proving that a given sample has already reached equilibrium. Distinguish the direction of net change from the simultaneous presence of both processes.

At dynamic equilibrium, forward and reverse reaction rates are equal and macroscopic composition remains stable under fixed conditions. Individual microscopic events continue. Equal reaction rates mean equal stoichiometric rates of progress, not necessarily one molecule turning back for each molecule formed in an equation with unequal coefficients.

Equilibrium does not require equal reactant and product concentrations. In an introductory concentration model, Kc is built from equilibrium concentrations raised to their balanced-equation coefficients, with pure solids and liquids omitted. A large K indicates a product-favored equilibrium expression, but actual concentrations also depend on starting amounts and stoichiometry. The thermodynamic equilibrium constant uses activities. Specify the equation and temperature before comparing K values.

The real power of equilibrium comes from being able to move it on purpose, and the rule for doing so is Le Châtelier's principle: if you disturb a system at equilibrium, it responds by shifting in the direction that partially undoes your disturbance. Change the concentration and the effect is intuitive — add more reactant and the system shifts toward the products to consume the excess; remove a product as it forms and the system shifts toward the products to replace it. Evaluate the new Q relative to K under the actual constraints; pure-phase additions and other special cases need their stated model.

Temperature and pressure work through the same principle but need a little more thought. Treat heat as if it were a reactant or product: in an exothermic reaction heat is a 'product,' so raising the temperature shifts the equilibrium backward (toward reactants), while cooling shifts it forward. For reactions involving gases, raising the pressure (by squeezing the volume smaller) shifts the equilibrium toward whichever side has fewer moles of gas, because that relieves the crowding. Counting the moles of gas on each side of the equation tells you which way it will go. This compression statement assumes the specified ideal reacting gas mixture and fixed temperature. Adding inert gas at fixed volume is different: total pressure rises without changing ideal reacting partial pressures, so that action need not shift composition.

The Haber process illustrates a rate-yield tradeoff: N₂(g) + 3H₂(g) ⇌ 2NH₃(g), with an exothermic forward reaction. Compression favors the side with fewer gas moles in the introductory model, while lower temperature favors ammonia at equilibrium but can slow production. Industrial processes choose conditions that balance rate, conversion, separation, energy cost, and equipment constraints, and use a catalyst and product separation. A catalyst accelerates the approach to equilibrium; it does not change K or the equilibrium composition at the same temperature.

Model check: Equal forward/reverse rates do not imply equal concentrations or that events stop. Reduced volume raises partial pressures; the adjustment favors the side with fewer gas moles, ammonia, under the model.

Equilibrium Constant

Keq = [products] / [reactants]

each concentration is raised to its coefficient from the balanced equation; for H₂ + I₂ ⇌ 2HI, Keq = [HI]² / ([H₂][I₂]). Keq greater than 1 favors products; less than 1 favors reactants.

Worked examples

Avoid these mistakes

• Equal forward/reverse rates do not imply equal concentrations or that events stop.

Words to know

reversible reaction

A reaction that can proceed in both the forward and reverse directions, written with a double arrow (⇌).

dynamic equilibrium

The state in which the forward and reverse reactions occur at equal rates, so the concentrations of reactants and products stay constant.

Le Châtelier's principle

If a system at equilibrium is disturbed, it shifts in the direction that partially counteracts the disturbance.

equilibrium shift

A temporary imbalance between the forward and reverse rates that moves the system to a new equilibrium position.

equilibrium constant (Keq)

A number equal to the product concentrations divided by the reactant concentrations (each raised to its coefficient) that shows whether products or reactants are favored.

Haber process

The industrial reaction N₂ + 3H₂ ⇌ 2NH₃, used to manufacture ammonia for fertilizer.

Haber, Bosch, and the reaction that feeds billions

In 1909 Fritz Haber demonstrated that nitrogen from the air could be forced to react with hydrogen to make ammonia, and engineer Carl Bosch then scaled it into the roaring high-pressure factories that still run today. The Haber process turned the fertilizer supply of the whole planet from a scarce natural resource into something manufactured on demand, and it is often said that the nitrogen in roughly half the protein in your body passed through a Haber-Bosch reactor. Both men won Nobel Prizes. Yet the same mastery of nitrogen chemistry let Haber develop poison gases in World War I — a stark reminder that a single reaction can nourish the world and be turned to terrible ends.

The two great myths about equilibrium

Equilibrium attracts two stubborn misconceptions, and untangling them is the key to the whole topic. The first is that equilibrium means the reaction has stopped — it has not; the forward and reverse reactions both race along at full speed, merely at equal rates, so nothing appears to change on the surface. The second is that equilibrium means reactants and products are present in equal amounts — they usually are not. Equilibrium only guarantees the amounts hold constant, and the equilibrium constant Keq tells you how lopsided that balance is. Keeping 'constant, not stopped' and 'constant, not equal' straight cures most equilibrium confusion.

Le Châtelier's principle inside your own bloodstream

Le Châtelier's principle is not just an industrial trick — it keeps you alive with every breath. In your lungs, where oxygen is plentiful, the equilibrium between hemoglobin and oxygen shifts toward binding; in your busy tissues, where oxygen is scarce, it shifts back toward releasing, delivering oxygen exactly where it is needed. Your blood also runs a carbonic-acid equilibrium (CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻) that buffers its pH, and when you exhale carbon dioxide you deliberately pull that equilibrium to keep your blood from turning acidic. Holding your breath drives it the other way, which is part of why you feel the urge to breathe.

Squeezing more ammonia out: the industrial balancing act

The Haber process is a masterclass in fighting a trade-off with Le Châtelier's principle. High pressure favors ammonia, because the reactant side has 4 moles of gas and the product side only 2, so squeezing the mixture pushes it toward the smaller volume. Low temperature also favors the exothermic forward reaction — but run too cold, and the reaction becomes uselessly slow. Industry therefore settles on a compromise near 450°C and 200 atmospheres, adds an iron catalyst to regain the lost speed, and continuously drains off the ammonia so the equilibrium keeps chasing the product it can never quite reach. Every lever in the plant is a deliberate application of this one principle.

What the size of Keq really tells you (and what it doesn't)

The equilibrium constant is a single number that captures how far a reaction proceeds. A Keq far greater than 1 means the products dominate at equilibrium — the reaction 'goes nearly to completion' — while a Keq far less than 1 means barely any product forms and reactants dominate. A Keq near 1 means reactants and products coexist in comparable amounts. But here is the crucial limit: Keq says nothing about how fast equilibrium is reached. A reaction can have an enormous Keq and still be agonizingly slow (diamond turning to graphite, for instance), which is exactly why kinetics and equilibrium are two separate questions this unit keeps side by side.

In the real world

The Haber process runs the equilibrium N₂ + 3H₂ ⇌ 2NH₃ under high pressure with an iron catalyst to mass-produce the ammonia fertilizer that supports roughly half the world's food supply.

Exit check

How does compressing N₂+3H₂⇌2NH₃ at fixed T affect equilibrium?

Equilibrium Constants & Reaction Quotients

Learning goals

• I can form K and Q expressions with phases and coefficients considered.
• I can predict net adjustment and solve a simple equilibrium composition.

Readings for this unit

• The catalyst can change the journey without changing the destination ratio

For aA+bB⇌cC+dD, an introductory equilibrium expression is K=[C]^c[D]^d/[A]^a[B]^b using equilibrium concentrations. Pure solids and liquids are omitted because their activities are approximately fixed. Rigorous equilibrium constants use dimensionless activities; school calculations often use concentration ratios with an implied standard state.

The reaction quotient Q uses the same form at any moment. Compare Q with K at the same temperature: Q<K means net forward change is favored; Q>K means net reverse change; Q=K means equilibrium. This comparison predicts direction, not the speed of approach.

Large K favors products in the represented expression and small K favors reactants, but neither means every reactant is absent or every product is absent. The equation’s orientation matters: reversing a reaction gives 1/K; multiplying all coefficients by n gives K raised to n.

An ICE table records initial, change, and equilibrium amounts. For A⇌B initially .80M A and .20M B, a forward change x makes A=.80−x and B=.20+x. Substitute these into K and solve, then check positive concentrations and the conservation relationship.

Adding species, changing volume, or removing product can change Q and cause adjustment. Changing temperature can change K. A catalyst speeds both directions and does not change K. Pressure changes influence gas equilibria through partial pressures and gas mole ratios, while adding an inert gas has condition-dependent effects.

Equilibrium is dynamic: forward and reverse events continue at equal rates, so macroscopic composition stays stable. Equal rates do not imply equal concentrations. Always define the reaction, temperature, phase conditions, and approximation before interpreting a numerical constant.

Worked examples

Avoid these mistakes

• Use equilibrium concentrations for K and current concentrations for Q.

Words to know

reaction quotient Q

The equilibrium-expression ratio evaluated for current conditions.

ICE table

A table tracking initial values, stoichiometric change, and equilibrium values.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

Why need not equilibrium concentrations be equal?