chemistry rate of reaction and equilibrium

temperature rate of reaction

Increasing the temperature increases the average kinetic energy of particles, leading to an increased frequency of collisions per unit time. A greater proportion of collisions involve particles with sufficient energy to overcome the activation energy, resulting in more successful collisions per unit time. As a result, the rate of reaction increases.

concentration rate of reaction

When the concentration of a reactant is increased, the number of particles per unit volume increases. This leads to a higher frequency of collisions per unit time between reactant particles. As a result, there are more successful collisions per unit time, which increases the rate of reaction.

surface area rate of reaction

Increasing the surface area of a solid reactant exposes more particles to potential collisions. This increases the frequency of collisions per unit time, as more reactant particles are available to interact. As a result, the number of successful collisions per unit time increases, leading to a faster rate of reaction.

pressure rate of reaction

Increasing the pressure of a gaseous system by decreasing its volume forces gas particles closer together, which increases the concentration of reactant particles per unit volume. This results in a higher frequency of collisions per unit time. As a result, there are more successful collisions per unit time, which increases the rate of reaction.

catalyst rate of reaction

A catalyst provides an alternative reaction pathway with a lower activation energy. This means that a greater proportion of collisions have sufficient energy to overcome the activation energy barrier. Some catalysts, such as enzymes, can also help orient reactant particles correctly, further increasing the likelihood of a successful collision. As a result, there is a greater frequency of successful collisions per unit time, which increases the rate of reaction.

3 factors - collision theory

• must collide with each other

• must collide with sufficient energy

• must collide in correct orientation

measures of testing rate of reaction

• mass loss on balance

• volume of gas produced

• rate of production of precipiate

• colour change end of reaction

transition state

max energy level in an energy profile diagram

activation energy

minimum amount of energy chemical reaction needs to overcome to react

homogenous catalyst

same state as the reactants and products

heterogenous catalyst

different physical state as products and reactants

open system

matter and energy can be lost to surroundings

closed system

no products or reactants escape from reaction mixture

dynamic equilibrium characteristics

• rate of forward reaction is same as the backward reaction

• concentrations of reactants and products is constant

• mixture of products and reactants

rate of reaction

how quick a reaction occurs

extent of reaction

how much of the reactants are converted into products

homogenous reaction

takes place in same physical state

heterogenous reaction

occurs at a boundary/interface between two physical states

meaning k>q

reactants is higher than products

meaning of k<q

reactants is lower than products

reverse equation, effect on k?

1/k

double equation, effect on k?

k²

half equation, effect on k?

square k

k>10⁴

products are higher than reactants

k<10^-4

reactants are higher than products

le chatliers principle

when a change is opposed on a system, it will partially oppose the change to establish a new equilibrium

shift to the left

concentration of reactants increases

shift to the left

concentration of products increases

increasing pressure

favour side with less gas molecules

decreasing pressure

favour side with more gas molecules

effect of inert gas

changes overall pressure but does not affect equilibrium

increase in temperature

favours endothermic direction

decrease in temperature

favours exothermic direction

k value effect of shift to left

is decreased

k value effect of shift to left

is increased