⏱️ Rates & Equilibrium

Rates of reaction

There are three common ways to measure the rate of a reaction in the lab: recording the mass lost on a balance as a gas escapes, collecting the volume of gas produced in a syringe, or timing how long a solution takes to turn cloudy.

Collision theory explains why reactions speed up or slow down: a reaction only happens when particles collide with enough energy, at least the activation energy. Anything that makes collisions more frequent or more energetic will speed up the reaction, for the following reasons:

• Increasing the concentration of a solution or the pressure of a gas packs more particles into the same volume, so collisions happen more often.
• Increasing the surface area, for example using a powder instead of a lump, exposes more particles and gives more opportunities for collisions.
• Increasing the temperature helps for two reasons: particles collide more often, and more of those collisions carry enough energy to react.
• A catalyst provides an alternative pathway with lower activation energy and is not used up in the reaction; enzymes are biology's catalysts.

Reversible reactions & equilibrium

Some reactions can run in both directions, shown as A + B ⇌ C + D. If the forward reaction is exothermic, the reverse reaction is endothermic by exactly the same amount. In a closed system, a reversible reaction settles into dynamic equilibrium: both the forward and reverse reactions are still happening, at the same rate as each other, so the amounts of reactants and products stop changing, even though they are not necessarily equal.