Collision Theory in Equilibrium – HSC Chemistry

1. Rate of collision between molecules. This is the frequency at which molecules collide. Greater rate of collision leads to greater reaction rate.

 

2. Activation energy of the reaction. For a reaction to occur, the reactants must overcome a certain amount of activation energy. Even if the collision rate is high, if the particles don’t have enough energy reaction will still not occur.

 

Figure: Molecular orientation is an important factor of reaction rate as stated by collision theory. If the orientation of reactants is not correct (effective) during collision, reaction does not occur.

 

3. Molecular orientation of reactants. Rate of reaction can increase if reactants collide in the right orientation. Conversely, rate can decrease if the orientation of molecules is not favourable for the formation of product(s).

 

Review Collision Theory in more detail from Year 11 Chemistry here 

Collision rate increases at a higher concentration

• Concentration of particles/molecules. A higher concentration means there are more particles moving about in a given volume, increasing the frequency at which particles collide.
  
  

Collision rate increases at a higher pressure for gases

• Pressure/volume. Changes in pressure and volume affect the collision rate between gaseous particles as they occupy the most volume. An increase in pressure (or reduction in volume) increases the rate of collision.

• Temperature is the measurement of the average kinetic energy in a system. A higher temperature (greater energy) means particles move about quicker at greater kinetic energy. This in turn increases the collision rate between them.

What Affects Activation Energy?

Catalyst reduces the activation energy

• Catalysts reduce the activation energy of a reaction, resulting in an increased rate.

Molecular Energy Distribution

• In addition to increasing the collision rate, a higher temperature also allows more collisions to result in a chemical reaction (more successful collisions). This is because, at a higher temperature, more molecules have energy greater than the activation energy to react. This causes the Maxwell-Boltzmann energy distribution graph to shift to the right, without changing the activation energy.

 

 

• Catalysts reduce the activation energy of a reaction. This does not affect the energy distribution of molecules as it does not change the energy inside the system. However, by reducing the activation energy, more molecules have enough energy to result in a chemical reaction when collisions occur.

 

 

• In both cases (higher temperature and adding a catalyst), the frequency of a successful collision (one that results in a chemical reaction) increases, and so does the reaction rate.

$$2NO_2(g) \rightleftharpoons N_2O_4(g)$$

 

 

• As the N₂O₄ molecules decompose to produce NO₂, their concentration decreases and so does their collision rate. This in turn reduces the rate of the forward reaction.

• As NO₂ molecules form, they can also collide and result in a reaction to re-form N₂O₄. This is represented by the reverse reaction.

• In the beginning, the forward reaction rate is greater than the reverse reaction rate as there are far more N₂O₄ than NO₂ molecules. This results in a net decrease in [N₂O₄] and a net increase in [NO₂]

 

 

• However, as the reaction proceeds, the rate of the forward reaction gradually decreases, and the rate of the reverse reaction gradually increases.
  
  

• This continues until the rates of the forward and reverse reaction become equal, where equilibrium is established. At this point, forward and reverse reactions still occur due to the collision between molecules but no changes in concentration are observed.

Practice Question

Consider a closed chemical system will only N₂O₄ molecules. 

Using collision theory, explain the changes in [N₂O₄] and [NO₂] shown by the two graphs below. 

 

Figure: change in reaction rate with time. At t₁, rate of forward and backward reactions becomes equal, and equilibrium is reached. 

 

 

Figure: change in concentration of reactants and products of a reversible reaction. [`N_2O_4`] decreases while [`NO_2`] increases. A reaction can reach equilibrium starting with any quantities of reactants and products.

 

Sample answer:

• Initially, N₂O₄ collide and will result in a reaction to produce NO₂ if the activation energy is met. This causes [NO₂] to increase and [N₂O₄] to decrease.
• As [N₂O₄] falls, there are fewer collisions between N₂O₄ molecules, so the rate of the forward reaction decreases.
• As [NO₂] increases, there are more collisions between NO₂ molecules, so the rate of the reverse reaction increases.
• Forward and backward reaction rates will continue to change until they become equal in value. When this occurs, the reaction reaches dynamic equilibrium. Since the two rates are equal, there is no change in [N₂O₄] and [NO₂]

Figure: Reaction between dinitrogen tetroxide and nitrogen dioxide can be monitored by its colour change. Nitrogen tetroxide is colourless whereas nitrogen dioxide has a distinctive brown colour. 

 

• Nitrogen dioxide (`NO_2`) has a distinctive brown colour while dinitrogen tetroxide (`N_2O_4`) is colourless. 

• In the reaction above, the reaction mixture is pale brown as it consists of mostly `N_2O_4`. As the reaction approaches equilibrium, the concentration of `NO_2` increases, making the mixture appear browner.

• When the reaction reaches a dynamic equilibrium, the concentration of `NO_2` and `N_2O_4` remain constant, so the colour of the mixture stays unchanged. This is used as an indication that the reaction has reached equilibrium.

Collision theory is useful in explaining the effect on a reaction at equilibrium when the concentration of a chemical species changes. 

Let's consider the following equilibrium:

$$N_2O_4(g) \rightleftharpoons 2NO_2(g)$$

 

 

If `[N_2O_4]` increases, collision between reactants increases which in turn increases the forward reaction rate. Now that the forward reaction rate is greater than the reverse reaction rate (reaction is no longer at equilibrium), `[NO_2]` will gradually increase, and `[N_2O_4]` will gradually decrease. 

 

 

Effect of Catalyst on Equilibrium

Figure: catalysts reduce the activation energy of both forward and reverse reactions. 

• Addition of catalyst(s) increases the rate of both forward and reverse reactions by reducing their respective activation energies.

 

Catalysts do not disturb chemical equilibria

Figure: addition of catalyst at the 10-minute mark increases the rate of both forward and reverse reaction. The equilibrium is not disturbed.

• Adding a catalyst to a reaction at equilibrium increases the forward and reverse reaction rates to the same extent. As a result, the reaction rates remain equal and equilibrium is not disturbed. 

 

Reactions reach equilibrium faster

Figure: addition of a catalyst allows the equilibrium to be reached faster.

 

In the presence of a catalyst, a reversible reaction reaches equilibrium faster. In the rate versus time graph above, note the following:

The catalysed reaction compared to the uncatalysed reaction

• has a higher initial reaction rate
• reaches equilibrium (equal reaction rates) faster, at an earlier time point
• both forward and reverse reaction rates are higher (but remains equal) at equilibrium

 

Previous section: Non-equilibrium Systems

Next section: Le Chatelier's Principle

 

RETURN TO MODULE 5: EQUILIBRIUM AND ACID REACTIONS