HS Chemistry - Reversible Reactions

Reversible Reactions & Dynamic Equilibrium

Overview of The Page

This page will cover:

• What are reversible reactions?
• What is dynamic equilibrium?
• What is the position of equilibrium?
• What factors affect equilibrium?

So far, we've only studied reactions that go in one direction - that is, the reactants form the products, and once all the reactants have formed the products, the reaction stops. This can be shown as:

Reactant 1 + Reactant 2 + … → Product 1 + Product 2 + …

Or alternatively:

Reactants → Products

In these reactions, all the reactants will combine to form the products, until the supply of one of the reactants is exhausted. For example, in a reaction where there are five moles of H₂ and 3 moles of Cl₂, and HCl is being produced:

The reaction continues until all the moles of Cl₂ have been exhausted, and no more product can be produced. At the end of the reaction, the products $6 moles of HCl$ and the leftovers $2 moles of H~2~$ remain.

But sometimes the reaction will also occur in reverse - that is, the products will react and re-form the original reactants. In the above example, that would mean 2 HCl molecules reacting with each other to form 1 H₂ and 1 Cl₂ molecule.

This is called a reversible reaction, and it can be written as:

Reactants ⇌ Products

This really shows two reactions simultaneously happening:

Reactants → Products

AND

Products → Reactants

The example above, if it were a reversible reaction, would be H₂ + Cl₂ ⇌ 2HCl, and it would consist of both H₂ + Cl₂ → 2HCl and 2HCl → H₂ + Cl₂. The forward reaction is H₂ + Cl₂ → 2HCl, and the reverse reaction is 2HCl → H₂ + Cl₂. In theory, all chemical reactions are reversible.

In a reversible reaction, both the forward and reverse reactions occur at the same time.

But unless they are at equilibrium, they don't occur at the same rate.

For example, at the beginning of the reaction:

There are H₂ and Cl₂ molecules present, so the forward reaction will occur, but since there are no HCl molecules present, the reverse reaction won't occur.

Similarly, at the other end of the reaction:

There are HCl molecules present on the other side, so the reverse reaction will occur. There are also leftover H₂ molecules, but since there are no Cl₂ molecules, the forward reaction won't occur.

At one end of the reaction, the reverse reaction had a rate of 0, and at the other end, the forward reaction had a rate of 0. If the reverse reaction's rate increases from 0 to some value as the reaction progresses, and the forward reaction's rate simultaneously decreases from some value to 0, then at some point the rates of both the forward and reverse reactions will be equal. At that point, the reaction is in dynamic equilibrium. Note that the quantities of molecules on either side don't need to be equal in dynamic equilibrium

This allows us to deduce a few important things:

• At equilibrium, since the rate of the forward and reverse reactions stays the same, the concentration of the reactants and products remains constant.

• Any system that isn't at equilibrium will correct itself so that it is at equilibrium, with more reactants or products being formed until equilibrium is reached. Therefore, in a system that is not in equilibrium, the concentrations of reactants and products will tend towards the concentrations at which equilibrium is achieved, and will stay there.

• The rate of the forward and reverse reactions depends on the concentration of the molecules on either side of the reaction. Therefore, since equilibrium is achieved when these rates are equal, the concentration of molecules on either side of the reaction affects when equilibrium is achieved.

• If matter is added or removed from the system, the concentration of molecules on one side of the reaction is changed, and the point at which equilibrium is achieved changes. Therefore, achieving equilibrium requires a closed system $one which prevents matter inside the system from escaping, or matter outside the system from entering$, as the amount of matter in a closed system remains equal.

The equilibrium is called dynamic equilibrium because reactant molecules are still reacting and forming product molecules, and product molecules are still reacting and forming reactant molecules, but the rate at which these two processes happen is equal, and thus there is no net change.

The point at which equilibrium is achieved in a reaction is called the position of equilibrium. If we try to visualize it on a spectrum ranging from 100% of the molecules at equilibirum being reactant molecules to 100% of them being product molecules, with a black arrow representing the position of equilibrium $not a formal way of showing it$:

The diagram above shows that the position of equilibrium is roughly in the center. When this system is in equilibrium, roughly 50% of the molecules will be reactant molecules and roughly 50% will be product molecules. It favors neither the reactants nor the products.

This system $a different one$ has a position of equilibrium that is closer to the reactants than the products. When this system is in equilibrium, there will be slightly more reactant molecules than product molecules. The position of equilibrium is shifted left $towards the reactants$, and the reaction slightly favors the reactants.

This system has a position of equilibrium that is much closer to the reactants than the products. When this system is in equilibrium, the vast majority of molecules will be reactant molecules, and very few will be product molecules. The position of equilibrium is shifted left, and the reaction strongly favors the reactants.

The position of equilibrium of this system is shifted right, and the reaction slightly favors the products.

The position of equilibrium of this system is shifted right, and the reaction strongly favors the products.

Note: These aren't formal diagrams. They're just shown to help visualize what is meant by position of equilibrium.

Factors that affect equilibrium

There are three factors that affect the equilibrium of a reaction: the concentration of the molecules, their pressure $for gases only$, and their temperature. They affect equilibrium in accordance with Le Chatelier's principle.

Le Chatelier's principle: If one or more of the factors that affect equilibrium changes, the position of equilibrium shifts in the direction that opposes $reduces$ the change.

• Concentration:

  • If the concentration of any of the reactant molecules is increased, the position of equilibrium will shift to the right $towards the products$. This is because if there are more reactant molecules, the forward reaction's rate will increase, and more product molecules will be formed. As this happens, the concentration of product molecules increases $increasing the reverse reaction's rate$ and the concentration of reactant molecules decreases $decreasing the forward reaction's rate$, until the rates of both reactions once again equal each other, at which point equilibrium is once again established. Overall, the position of equilibrium shifted to the right when the concentration of reactant molecules was increased. Conversely, when the concentration of any of the product molecules is increased, the position of equilibrium will shift to the left $towards the reactants$.

• Pressure $only affects equilibrium if gas molecules are a part of the reaction$:

  • The reason increasing the pressure only affects equilibrium if gas molecules are present is because only gases can be compressed to smaller volumes $making pressure significant$. Increasing the pressure of a system will increase the pressure on both the reactant and product sides equally. Therefore, when pressure is increased, whichever side has a larger number of molecules in the equation ends up with more of its molecules closer together. That side of the equation then reacts more, producing more molecules of the other side and shifting the position of equilibrium towards that side $the side with fewer molecules$.

  • For example, in the reaction:

    O₂ + 2F₂ ⇌ 2F₂O

    The reactant side has 3 molecules $1 O~2~ molecule and 2 F~2~ molecules$, while the product side has 2 molecules $2 F~2~O molecules$. Therefore, when the pressure of the system is increased, the reactant molecules will see a greater increase in how much they react with each other, since there are more of them. This will cause more product molecules to be formed, shifting the position of equilibrium to the right $towards the products$. It follows that if the pressure of the system were to be decreased, the position of equilibrium will be shifted to the left $towards the reactants$.

• Temperature:

  • The reason increasing the temperature of the system affects equilibrium is because every reaction has an enthalpy change. That means that in a reversible reaction, one of the reactions $either the forward or reverse reaction$ will be exothermic, and the other must be endothermic $as it is reversing the first one$. If the temperature of the system is increased, there is more heat energy available in the surroundings. The endothermic reaction, which absorbs heat, can do so more easily, and therefore the rate of the endothermic reaction increases.

  • When the temperature of a system increases, the rate of the endothermic reaction in a reversible reaction $it can be either the forward or reverse reaction$ also increases. Conversely, when the temperature of a system decreases, the exothermic reaction starts reacting faster.

    • In the reaction:

      H₂ + Cl₂ ⇌ 2HCl

      The forward reaction $H~2~ \+ Cl~2~ → 2HCl$ is exothermic, while the reverse reaction $2HCl → H~2~ \+ Cl~2~$ is endothermic. When the temperature of the system increases, the rate of the reverse reaction will increase. The position of equilibrium will shift to the left $towards the reactants$. Conversely, when the temperature of the system decreases, the rate of the forward reaction $the exothermic reaction$ will increase, shifting the position of equilibrium to the right $towards the products$.

    • In a reversible reaction where the forward reaction is endothermic and the reverse reaction is exothermic, an increase in the temperature of the system will still cause the rate of the endothermic reaction to increase - in that case, the endothermic reaction is the forward reaction rather than the reverse reaction.