Have you heard of the word equilibrium? What about dynamic equilibrium? Sometimes chemical reactions are not unidirectional; they are reversible. The reactants react to create the products, but then the products can also react to recreate the original reactants. This means there is both a forward reaction and a reverse one. In this article, we will learn how reversible reactions can reach an equilibrium state and how different factors affect this equilibrium in the system.
The definition of dynamic equilibrium
When the forward and reverse reaction rates are the same, this system is said to be at dynamic equilibrium. It is called dynamic because the forward and reverse reactions are happening, but since their rates are the same, they cancel each other's effects, so there doesn't appear to be any activity.
Dynamic equilibrium occurs when the forward reaction rate equals the reverse reaction rate, and the amounts of reactants and products stay constant.
It is essential to understand that the number of reactants and products are not necessarily identical in an equilibrium. In an equilibrium state, the ratio of products to reactants can be any number as long as their concentration or amount is not changing.
The equilibrium state can occur in both physical and chemical processes. A physical equilibrium develops between different phases or physical properties and does not involve chemical changes in the system.
And as you can already guess, a chemical equilibrium develops in chemical processes.
The rate of the forward reaction is equal to the rate of the backward reaction in chemical equilibrium. As a result, the reactant and product concentrations do not vary throughout time at the equilibrium state as long as other factors are kept constant.
The system's temperature, pressure, and concentration are all factors that influence equilibrium. When one of these components changes, the system's balance is disrupted, and the system readjusts until it returns to equilibrium. Depending on the nature of the reaction, changes in the system's conditions may favour the forward or the reverse reactions. In that case, the equilibrium position would shift in line with the favoured reaction.
Before heading into how these factors impact the equilibrium state, we need to go over a fundamental principle, Le Chaterlier's.
Le Chaterlier's principle
French chemists Le Chatelier and Braun came up with a generalisation to explain the impact of changes in temperature, concentration, or pressure on the state of an equilibrium system. This generalisation is known as Le Chaterlier's principle. According to Le Chatelier's principle, if a change is made to equilibrium, the equilibrium will act to resist the change. This simple theory is beneficial for predicting the overall effect of a change in concentration, pressure, or temperature on an equilibrium system.
Le Chaterlier's principle applies to both chemical and physical equilibria.
Effect of concentration on equilibrium
Following the Le Chatelier principle, when the concentration of reactants or products in an equilibrium reaction changes, the makeup of the equilibrium mixture changes to lessen the impact of the concentration change.
Take the reaction in the Haber process for example: N2(g) + 3H2(g) ⇌ 2NH3(g)
The concentration of chemicals on either side of the equilibrium can be changed. Suppose we add more reactants (N2 + H2) to the mixture. This will upset the equilibrium, causing the forward reaction to accelerate to consume some new reactants and convert them into products, restoring equilibrium. In this example, the equilibrium position is considered to have moved to the right.
On the other hand, when we add more products (NH3), the equilibrium shifts to the left, converting the extra ammonia to nitrogen and hydrogen and minimising the impact of the ammonia added.
You might wonder what would happen to the equilibrium if we remove only one of the reactants or products. The principle is still the same. The equilibrium shifts to create more of the substance that has been added to restore balance. If we add more nitrogen to the mixture, for example, the equilibrium will shift right to create use up the added nitrogen and produce more ammonia, opposing the change.
Effect of temperature on equilibrium
The impact of temperature on an equilibrium depends on which reaction, forward or reverse, is exothermic and which one is endothermic.
Let's quickly refresh our memories on what exothermic and endothermic mean.
Exothermic chemical reactions are those that release energy, often in the form of heat.
In exothermic reactions, the collective energy released from breaking the chemical bonds in the reactants is greater than the energy needed for forming new chemical bonds in the products. As a result, the surplus in energy is released to the surroundings during the reaction. Due to the energy released, the temperature of the exothermic reaction mixture rises as the reaction takes place.
Endothermic chemical reactions, on the other hand, absorb energy during the reaction.
In endothermic reactions, the bonds in the products hold more energy than the energy released from breaking the bonds in the reactants. Hence, more energy needs to be absorbed by the reaction mixture to form the bonds in the products. In other words, the kinetic energy of the particles in the mix becomes converted to the chemical energy stored in the bonds of the products. Since temperature measures the average kinetic energy of the particles in a system, the temperature of the reaction mixture decreases in endothermic reactions.
The two forward and reverse reactions cannot both be exothermic or endothermic. This is because they are opposite reactions to one another. So, if the forward reaction is endothermic, for example, then the reverse reaction will be exothermic.
Le Chaterlier's principle is still standing when it comes to changes in temperature. The makeup of the equilibrium shifts to oppose any changes in the system's temperature.
The forward reaction is exothermic in the Haber process, making the reverse endothermic.
N2(g) + 3H2(g) ⇌ 2NH3(g) qforward < 0, ΔHforward < 0
When heat is released to the reaction mixture, the reaction is exothermic, and q for the reaction has a negative value. When the reaction mixture absorbs heat, then the reaction is endothermic, and the q for the reaction has a positive value.
At constant pressure and temperature, q is the same as the enthalpy change (ΔH).
'q' denotes the heat absorbed or released during a reaction.
If heat is absorbed, q is positive, and if heat is released to the system and the environment, q is negative.
If the temperature of the equilibrium mixture is increased, the equilibrium will shift towards the endothermic side to oppose and decrease the mixture's temperature. So in the example above, when the mixture's temperature is raised, the equilibrium shifts to the left towards the reactants (N2 + H2) until it reaches a new equilibrium state.
You can probably guess what would happen when the mixture's temperature is decreased. The equilibrium would shift towards the exothermic reaction, in this case towards the product (NH3), since the energy released from the forward exothermic reaction would raise the temperature to oppose the change.
Effect of pressure or volume on equilibrium
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