Question

Brainliest if answered in the next 5 minutes

Answer 1

The equilibrium constant "k" is used to determine the concentration of every product in a reversible reaction until equilibrium is reached. At equilibrium, the value of k is equal to 1.

When the value of k is less than 1, this means that the concentration of the reactants is more than the concentration of the products. To reach equilibrium again, the reaction will shift to the left forming more products to compensate for the concentration.

Based on this, there are only two choices which are: 
The concentration of one or more of the products is small.
The reaction will not proceed very far to the right.

The other two are incorrect as:
The concentration of the reactants is more and not less
The reaction will generally form more products than reactants.

Answer 2

Answer: The correct statements are the concentration of one or more of the products is small, the reaction will not proceed very far to the right and the reaction will generally form more reactants than products.

Explanation:

$K_{eq}$ is defined as the equilibrium constant of the reaction. It is basically the ratio of concentration of products to the concentration of reactants, each raised to the power their stoichiometric coefficients.

For a reaction:

$aA+bB\rightarrow cC+dD$

The expression for $K_{eq}$ is:

$K_{eq}=\frac{[C]^c[D]^d}{[A]^a[B]^b}$

When $K>1$, forward reaction is favored and when $K$, backward reaction is favored.

When K < 1, the expected possibilities are:

• The reaction will proceed in the left direction
• The reaction will lead to the formation of reactants more than the products.
• The concentration of reactants is more than the concentration of products.

Hence, the correct statements are the concentration of one or more of the products is small, the reaction will not proceed very far to the right and the reaction will generally form more reactants than products.