Question

1. On the Section 7 (Part 2) text handout (pp. 3/4), the implications of the magnitude of K is presented, with respect to predictions about product/reactant favoured equilibria. The general examples used were simple: i.e. A↔B and C↔D. When real examples of chemical equilibrium are used, the conclusions about reactant- or product-favoured versus the value of K must be modified (e.g. real equilibria can have K>1 but still be reactant-favoured). Discuss this idea, using examples of real chemical equilibria and assumed values of equilibrium concentrations.

Answer 1

An example of such a real life chemical reaction is as follows:

Consider the following reaction:

H₂ (g) + O₂ (g) <----> H₂O₂ (g)

This reaction attains equilibrium at the following conc. values:

[H₂] = [O₂] = 5*10⁻³ mol.dm⁻³ and [H₂O₂]= 4*10⁻⁵ mol.dm⁻³

At these values, we can see that Kc = 1.6, but still conc. of reactants is about 250 times greater.

To understand this, it's useful to see Kc as the ratio of forward reaction rate constant to the backward reaction rate constant.

That is: K_c = k_forward / k_backward

K_c > 1 implies that k_forward > k_backward. This doesn't mean that because the reactants are in large excess at equilibrium the reaction is reactants favored. NO !! . K_c > 1 strictly means that reaction has a greater tendency to move forward, if allowed somehow.

Adding up the reactants and products concentrations and using them to tell about the favorability of reaction is totally unfair and wrong at many aspects.