Question

Consider the reversible reaction A ( g ) − ⇀ ↽ − B ( g ) Which K values would indicate that there is more B than A at equilibrium? K = 9 × 10 – 5 K = 0.4 K = 5000 K = 8 × 10 7

Answer 1

Answer:

5000 and $8\times 10^{7}$ indicate that there is more B than A at equilibrium

Explanation:

For the given reaction: $K=\frac{[B]}{[A]}$

where [B] and [A] represents equilibrium concentration B and A respectively. K represents equilibrium constant

More B than A at equilibrium means, [B] > [A]

So, $K=\frac{[B]}{[A]}>1$

As, both 5000 and $8\times 10^{7}$ are greater than 1 therefore these two K values indicate that there is more B than A at equilibrium