1. What is the Equilibrium Constant (Kc)?

The equilibrium constant (Kc) quantifies the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients. It indicates the extent of a chemical reaction at equilibrium.

2. How Does the Calculator Work?

The calculator uses the equilibrium constant formula:

Where:

• \( [Products] \) — Equilibrium concentrations of products (mol/L)
• \( [Reactants] \) — Equilibrium concentrations of reactants (mol/L)
• \( coefficients \) — Stoichiometric coefficients from balanced equation

Explanation: The equation shows that Kc is the product of product concentrations (each raised to their coefficient) divided by the product of reactant concentrations (each raised to their coefficient).

3. Importance of Equilibrium Constant

Details: Kc helps predict reaction direction, extent of reaction, and how changes in conditions affect equilibrium. Large Kc values favor products, small values favor reactants.

4. Using the Calculator

Tips: Enter concentrations in mol/L and coefficients as whole numbers. Separate multiple values with commas. For example, for the reaction 2A ⇌ B + 3C, enter reactant concentration as "0.5" and coefficient as "2"; products as "0.1,0.3" and coefficients as "1,3".

5. Frequently Asked Questions (FAQ)

Q1: What units does Kc have?
A: Kc is technically unitless, but its numerical value depends on the concentration units used (typically mol/L).

Q2: How does temperature affect Kc?
A: Kc changes with temperature. Exothermic reactions have Kc decrease with temperature increase, endothermic reactions have Kc increase.

Q3: What if solids or pure liquids are in the reaction?
A: Their concentrations are constant and not included in Kc expressions.

Q4: What does Kc = 1 mean?
A: At equilibrium, product and reactant concentrations are balanced (but not necessarily equal).

Q5: How is Kc different from Kp?
A: Kc uses concentrations (mol/L), while Kp uses partial pressures (atm) for gas-phase reactions.