6.4 Reaction Rates and Equilibrium

Reaction Rates

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Factors Affecting Reaction Speed

Reaction rate measures how quickly reactants are used up or products are formed over time. For a reaction to actually happen, colliding particles need enough energy to break existing bonds and form new ones. That minimum energy threshold is called activation energy. If particles collide without enough energy, they just bounce off each other and nothing changes.

Five major factors control how fast a reaction proceeds:

• Temperature — Higher temperatures give particles more kinetic energy, so they move faster, collide more often, and hit harder. Both the frequency and the force of collisions increase, which speeds up the reaction.
• Concentration — More reactant particles packed into the same space means more collisions per second. Double the concentration, and you roughly double the chance of particles finding each other.
• Surface area — Grinding a solid into smaller pieces exposes more of it to the other reactant. A sugar cube dissolves slowly; powdered sugar dissolves almost instantly.
• Catalysts — These substances speed up a reaction by lowering the activation energy, giving particles an easier path to react. The catalyst itself isn't consumed, so it can be reused over and over.
• Nature of reactants — Some substances are simply more reactive than others. Sodium reacts violently with water, while gold barely reacts with anything.

Measuring and Controlling Reaction Rates

Scientists track reaction rates by monitoring observable changes over time. Common techniques include tracking color changes (spectrophotometry), measuring electrical conductivity, or collecting gas produced and recording its volume at regular intervals.

The rate law is an equation that links the reaction rate to the concentrations of reactants. For a generic reaction where A turns into products, a rate law might look like:

$\text{Rate} = k[A]^n$

Here, $k$ is the rate constant (specific to each reaction and temperature), $[A]$ is the concentration of reactant A, and $n$ is the order of reaction, which tells you how strongly that reactant's concentration affects the rate. If $n = 2$, doubling $[A]$ quadruples the rate. If $n = 1$, doubling $[A]$ simply doubles the rate.

For reactions that happen in multiple steps, the overall speed is controlled by the slowest step, called the rate-determining step. It acts like a bottleneck: no matter how fast the other steps are, the reaction can't go faster than its slowest stage.

Chemical Equilibrium

Understanding Chemical Equilibrium

Many reactions don't just go in one direction. In a reversible reaction, products can react to re-form the original reactants. Over time, the forward reaction slows down (as reactants are consumed) and the reverse reaction speeds up (as products accumulate). Eventually, both reactions proceed at the same rate.

This is chemical equilibrium: the point where the rate of the forward reaction equals the rate of the reverse reaction. Concentrations of reactants and products stop changing, but both reactions are still happening continuously. That's why it's called a dynamic equilibrium. It looks static from the outside, but there's constant activity at the molecular level.

The equilibrium constant ($K$) puts a number on where the balance sits. For a reaction $aA + bB \rightleftharpoons cC + dD$:

$K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

Products go in the numerator, reactants in the denominator, and each concentration is raised to the power of its coefficient in the balanced equation.

• A large $K$ (much greater than 1) means products are favored at equilibrium.
• A small $K$ (much less than 1) means reactants are favored.
• A $K$ close to 1 means neither side is strongly favored, and you'll find significant amounts of both reactants and products.

Two types of equilibrium worth knowing:

• Homogeneous equilibrium — all reactants and products are in the same phase (e.g., all dissolved in solution).
• Heterogeneous equilibrium — reactants and products exist in different phases (e.g., a solid reacting with a gas). In these cases, pure solids and pure liquids are left out of the $K$ expression because their concentrations effectively don't change.

Manipulating Equilibrium Systems

Le Chatelier's principle states that if you disturb a system at equilibrium, it will shift to partially counteract that disturbance and establish a new equilibrium. Three main types of disturbance matter here:

• Concentration changes — Adding more reactant pushes the equilibrium toward products (the system "uses up" the extra reactant). Removing a product has the same effect, because the forward reaction speeds up to replace what was taken away. The reverse applies too: adding product or removing reactant shifts equilibrium back toward reactants.
• Pressure changes (for gaseous reactions) — Increasing pressure favors the side with fewer moles of gas. For example, in $N_2 + 3H_2 \rightleftharpoons 2NH_3$, the reactant side has 4 total moles of gas and the product side has 2. Increasing pressure shifts equilibrium toward $NH_3$ production. If both sides have the same number of moles of gas, a pressure change has no effect on the equilibrium position.
• Temperature changes — This depends on whether the reaction is exothermic or endothermic. Raising the temperature favors the endothermic direction (the system absorbs the added heat). Lowering the temperature favors the exothermic direction. Temperature is the only factor here that actually changes the value of $K$ itself.