To predict the outcome of a Bronsted-Lowry reaction requires information about the equilibrium constants (a Table of the K_as) for each acid. Listed below are two entries showing the acid and its conjugate base. The acid having the larger K_a is a stronger acid.

Here are four scenarios.

1. H₂SO₃ + H₂C₂O₄ --> No reaction since both are acidic.

2. SO₄^2- + CO₃^2- --> No reaction since both are basic.

In a Bronsted-Lowry reaction, there is an acid on the left or reactant side competing with an acid on the right or product side. The stronger acid will cause a shift to the opposite of the equation. Suppose there are two competing acids HCO₃^- and HSO₄^-. HSO₄^- is the stronger acid and sets up the two following possibilities.

3. SO₄^2- + HCO₃^- --> HSO₄^- + CO₃^2- Reaction is not favored, since the HSO₄^- is the stronger acid than HCO₃^-.

4. HSO₄^- + CO₃^2- --> SO₄^2- + HCO₃^- Reaction is favored, since HSO₄^- is a stronger acid than the HCO₃^-, the acid on the product side.

The fourth scenario works owing to the fact that it has a large K_eq. This is the result of the K_a for the dissociation of HSO₄^- being multiplied by the reverse of K_a for the dissociation for HCO₃^-. Recall that the K_reverse is the reciprocal of the K_forward, or K_r =1/K_f. The K_r = 1/(5.6x10^-11). Examine the addition of the forward reaction of HSO₄^- and the reverse reaction of CO₃^2-. Recall that K₂ = 1/K_a for HCO₃^-. The product of K₁ x K₂ = K_eq. In this case, the K_eq is large and products are favored.