Question

For a certain reaction `2A + B rightleftharpoons C + 3D`, `K_(eq) = 4.2 xx 10^3`, 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.

Answer 1

See below>

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.

Answer 2

The equilibrium constant for this reaction is:

`"K"_(eq) = \frac(["products"])(["reactants"]) = \frac([C][D]^3)([A]^2[B]) = 4.2xx10^3`

`"K"_(eq) > 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. `DeltaG^@ = -RTlnK_(eq)`.

(c) This is backwards. If you had more reactants than products, then `0 < "K"_(eq) < 1`, since `["reactants"] < ["products"]`.

(d) This is only true if `"K"_(eq) = 1`, i.e. `["reactants"] = ["products"]`.