For a certain reaction 2A+BC+3D, Keq=4.2×103, which of the following is true?

a) The reaction is product-favored.
b) The reaction rate is fast.
c) There are more reactants than products.
d) The reaction is neither product-favored nor reactant-favored.

2 Answers
May 13, 2016

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Explanation:

The reaction lies to the right side because of the following:

The expression of the equilibrium constant K is:

K=[Products][Reactants]

Since K>1 this means that the [Products] is greater than the [Reactants] and thus there is more products than reactants, and therefore, the equilibrium lies to the right.

Here is a video that I recently made about this topic:

Chemical Equilibrium | Reaction Quotient & Application of a Large K.

May 13, 2016

The equilibrium constant for this reaction is:

Keq=[products][reactants]=[C][D]3[A]2[B]=4.2×103

Keq>1, thus [products]>[reactants], and the equilibrium lies to the "right" (a), or the "products" side.

(b) We cannot say what the rate of reaction is, because that's a kinetic description of a reaction whose kinetic quantities are unstated. We only know the equilibrium constant, which is a thermodynamic quantity. ΔG=RTlnKeq.

(c) This is backwards. If you had more reactants than products, then 0<Keq<1, since [reactants]<[products].

(d) This is only true if Keq=1, i.e. [reactants]=[products].