Reaction Rates and Equilibrium

Chapter 18: Pages 541-585

Terms to know•Rate•Collision theory•Activation energy•Inhibitor•Catalyst•Pressure

•Reversible reaction

•Equilibrium•Dynamic

equilibrium•Le Chatelier’s

principle

Collision Theory• In order for chemical reactions to occur,

particles must collide with a certain minimum amount of kinetic energy (activation energy).

• Particles must also collide “head on.” Glancing collisions do not result in reactions.

• As the chance of collisions increases, the chance of a chemical change (reaction) occurring increases correspondingly.

Factors affecting collision Rate

• Temperature: When particles are heated, the particles gain kinetic energy and move faster, increasing the chance of collision between reactant molecules.

• Concentration and/or pressure: When particles are crowded together (Increased concentration or pressure), they are more likely to bump into each other so there is an increased chance of collisions.

• Particle size: Smaller particles have more surface area (compared to volume) in order to achieve head on collisions that result in chemical reactions.

• Let’s take a look!

A Couple of Other factors affecting collision rates

• Catalysts lower activation energy making it easier for particles to collide effectively.

• Inhibitors (like food preservatives) make it so that catalysts don’t work; thus, slowing reducing effective collisions.

Reversible Reactions• Some reactions occur in both directions at

once.• In other words, reactants are making products

at the same time the products are making reactants.• Forward reaction:

• 2SO2 + O2 2SO3 • Reverse reaction:

• 2SO3 2SO2 + O2

• Combined reaction:• 2SO2 + O2 2SO3

Chemical Equilibrium• Remember reversible reactions…

• 2SO2 + O2 2SO3

• At first, the reaction will occur in one direction more than the other

• After some time, the rates of the forward and reverse reactions will equalize.

• Chemical equilibrium is reached when the rates of the forward and reverse reactions equalize.

Dynamic equilibrium• At equilibrium, the amounts of reactant and

product reach a steady state, since the forward and reverse reaction rates are equal.

• It is important to remember that the reactions haven’t stopped. They are happening at the same rate, so the amount of products and reactants stays the same.

• Dynamic equilibrium means that the reactions are still happening (dynamic) but the amount of products and reactants are staying the same.

Equilibrium Point• Chemical equilibrium occurs at

the point when the forward and reverse reactions occur at the same rate.

• This does not mean that there are equal amounts of reactants and products.

• Equilibrium point is the amount of product and reactant present when a reaction is at chemical equilibrium.

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Equilibrium means balance

• Chemical equilibrium is a delicate balance. Just like a person responds to a push in order to keep balance, a chemical reaction does the same.

• Le Chatelier’s Principle says that if a stress is applied to a reaction at equilibrium, the reaction responds in a way that relieves the stress.

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Stresses on Equilibrium: Concentration

• Concentration or amount: As you add more of a reactant or product to a reaction, the reaction will produce more of the other.

2SO2 + O2 2SO3

• If I add more SO2 and/or O2, the reaction will produce more SO3 to relieve the stress.

• If I add more SO3, the reaction will produce more SO2 and O2 to relieve the stress.

Stresses on Equilibrium: Temperature

• Temperature: As you add more heat, the reaction will respond in a way to reduce or give off heat. If you take away heat, the reaction will respond in a way to produce more heat.

2SO2 + O2 2SO3 + heat

• If I add heat, the reaction will produce more SO2 and O2.

• If I take heat away, the reaction will produce more heat and SO3

Stresses on equilibrium: Pressure

• Pressure: Changes in pressure only affect gaseous reactants and products in chemical reactions. • Increasing the pressure causes the reaction to respond

in a way that will reduce the number of moles of gas.• Decreasing the pressure causes the reaction to respond

in a way that increases the number of moles of gas.

• N2(g) + 3H2(g) 2NH3(g)• If I increase the pressure, more NH3 will be produced

since there are only 2 moles of it.• If I decrease the pressure, more N2 and H2 will be

produced since there is more of it (1 + 3 = 4 moles).