What is the difference between a homogeneous and a heterogeneous equilibrium system in the context of $K_c$?

In a homogeneous system, all reactants and products are in the same phase and are all included in the $K_c$ expression. In a heterogeneous system, species are in different phases, and pure solids (or sometimes liquids) are omitted from the expression.

How does the $K_c$ expression change if the stoichiometric coefficients of a balanced equation are doubled?

Since coefficients become exponents, doubling the coefficients squares the entire $K_c$ expression. The new equilibrium constant would be the square of the original $K_c$ value.

When comparing $K_c$ values, what does a value of $K_c = 10^{-8}$ indicate compared to $K_c = 10^{8}$?

A very small $K_c$ ($10^{-8}$) indicates the equilibrium lies far to the left (mostly reactants), while a very large $K_c$ ($10^{8}$) indicates the equilibrium lies far to the right (mostly products).

What is the most common error when writing the notation for concentration in a $K_c$ expression?

Using round brackets $( )$ instead of square brackets $[ ]$. Square brackets are the specific chemical notation required to represent molar concentration ($mol\ dm^{-3}$).

Why is it incorrect to include a solid reactant, such as $Mg(s)$, in a $K_c$ expression?

The concentration of a pure solid is constant because its density does not change. Including it would be redundant as its constant value is already built into the $K_c$ constant itself.

What happens to the $K_c$ expression if you accidentally place reactants in the numerator and products in the denominator?

You would be calculating the reciprocal of the correct $K_c$ ($1/K_c$). This represents the equilibrium constant for the reverse reaction rather than the forward reaction.

Define the equilibrium constant $K_c$.

It is the ratio of the equilibrium concentrations of products to reactants, with each concentration raised to the power of its stoichiometric coefficient at a constant temperature.

What do the square brackets $[ ]$ specifically represent in a $K_c$ expression?

They represent the equilibrium concentration of a substance, typically measured in units of $mol\ dm^{-3}$.

What is a 'homogeneous' reaction mixture?

A reaction mixture where all the reactants and products are in the same physical state or phase, such as all being gases or all being in aqueous solution.

Why does $K_c$ only change with temperature and not with pressure or concentration?

Changes in pressure or concentration shift the position of equilibrium to maintain the ratio defined by $K_c$. Only temperature alters the ratio of the rate constants for the forward and backward reactions, thus changing $K_c$.

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1. Physical Chemistry

Deducing Kc Expressions

Summary

The equilibrium constant ( $K_c$ ) provides a quantitative measure of the position of equilibrium for a reversible reaction at a specific temperature. It is derived from the concentrations of products and reactants at equilibrium, adjusted for their stoichiometric coefficients.

1. Definition & Core Concepts

A diagram showing the mathematical structure of the Kc expression, highlighting that products are in the numerator, reactants are in the denominator, and coefficients become exponents.

2. Underlying Principles

Stoichiometric Exponents: The coefficients from the balanced chemical equation (the 'moles' of each substance) must be used as the powers for the concentrations in the expression.
Temperature Dependence: $K_c$ is a constant for a specific reaction at a constant temperature. If the temperature changes, the value of $K_c$ will change, reflecting a shift in the equilibrium position.
Dynamic Nature: The expression uses equilibrium concentrations, not initial concentrations. At equilibrium, these concentrations remain constant even though the forward and backward reactions continue to occur at equal rates.

3. Methods & Techniques

4. Key Distinctions

5. Interpreting the Kc Value

6. Exam Strategy & Tips