What is the equilibrium constant K?

K quantifies the position of equilibrium. Kc uses concentrations; Kp uses partial pressures. Large K favors products; small K favors reactants.

What is an ICE table?

ICE tables organize Initial, Change, and Equilibrium concentrations for solving equilibrium problems. Essential for finding K from initial conditions.

Does K change with concentration?

No. K depends only on temperature. Changes in concentration shift the equilibrium position but not the constant.

What is Le Chatelier's principle?

When disturbed, a system at equilibrium shifts to minimize the disturbance. Increase reactant → shifts right; increase product → shifts left.

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PHYSICAL CHEMISTRYChemical EquilibriumChemistry Calculator

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Equilibrium Constant Keq: Kc, Kp, and Gibbs Relation

The equilibrium constant K quantifies the position of equilibrium. Kc uses concentrations; Kp uses partial pressures. They relate via Kp = Kc(RT)^Δn. Thermodynamic K° links to ΔG° = -RT ln K. Large K favors products; small K favors reactants.

Concept Fundamentals

[products]/[reactants]

Kc × (RT)^Δn

-RT ln K

ΔG°

Initial, Change, Eq

ICE

Calculate Equilibrium ConstantsKc, Kp, ICE tables, Q vs K

Why This Chemistry Calculation Matters

Why: Equilibrium constants predict product yields, guide process design, and connect kinetics to thermodynamics. K changes only with temperature.

How: Kc = [C]^c[D]^d / ([A]^a[B]^b) at equilibrium. For gases, Kp = Kc(RT)^Δn. Use ICE tables to find equilibrium concentrations from initial conditions.

●Pure solids and liquids have activity 1 and are omitted from K.
●K depends only on temperature; concentration changes shift position, not K.
●ΔG° = -RT ln K links thermodynamics to equilibrium.

Sources:IUPAC Gold BookNIST Chemistry WebBook

Equilibrium Examples

⚗️ Haber Process

N₂ + 3H₂ ⇌ 2NH₃ at 25°C

🧪 Acetic Acid Dissociation

CH₃COOH ⇌ CH₃COO⁻ + H⁺

🔄 Kc to Kp Conversion

Convert Kc = 3.5e8 to Kp

📊 ICE Table Example

Calculate Kc from ICE table

⚖️ Q vs K Comparison

Determine reaction direction

💨 Water Gas Shift

CO + H₂O ⇌ CO₂ + H₂

🌡️ N₂O₄ Decomposition

N₂O₄ ⇌ 2NO₂

📈 Kp from Pressures

Calculate Kp from partial pressures

Calculate Equilibrium Constant

Calculation Mode

Temperature (°C)

Reactants

Reactant 1 Concentration (M)

Stoichiometry

Products

Product 1 Concentration (M)

Stoichiometry

Δn (moles gas products - moles gas reactants)

For educational and informational purposes only. Verify with a qualified professional.

🔬 Chemistry Facts

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Kc uses concentrations (M); Kp uses partial pressures (atm).

— IUPAC

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Kp = Kc × (RT)^Δn where Δn = moles gas products − moles gas reactants.

— Physical chemistry

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ΔG° = -RT ln K; K > 1 means ΔG° < 0 (spontaneous as written).

— IUPAC

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ICE tables: Initial, Change, Equilibrium for solving equilibrium problems.

— NIST

What is an Equilibrium Constant?

The equilibrium constant (K) quantifies the position of equilibrium for a reversible chemical reaction. It relates the concentrations (Kc) or partial pressures (Kp) of products and reactants at equilibrium. A large K (>1) favors products, while a small K (<1) favors reactants.

Kc = [Products]^stoichiometry / [Reactants]^stoichiometry

For aA + bB ⇌ cC + dD: Kc = [C]^c[D]^d / [A]^a[B]^b

Kc vs Kp: Concentration vs Pressure

Kc (Concentration-Based)

Used for reactions in solution (aqueous or liquid phase). Units depend on stoichiometry.

Kc = [C]^c[D]^d / [A]^a[B]^b

Example: CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)

Kp (Pressure-Based)

Used for gas-phase reactions. Expressed in terms of partial pressures.

Kp = (P_C)^c(P_D)^d / (P_A)^a(P_B)^b

Example: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Conversion Formula

Kp = Kc × (RT)^Δn

Where Δn = (moles of gas products) - (moles of gas reactants), R = 0.0821 L·atm/(mol·K), T = temperature in Kelvin

ICE Tables: Initial, Change, Equilibrium

ICE tables help organize information about initial concentrations, changes during reaction, and equilibrium concentrations. They're essential for solving equilibrium problems.

📊 ICE Table Structure

Species	Initial (I)	Change (C)	Equilibrium (E)
CH₃COOH	0.10 M	-x	0.10 - x
CH₃COO⁻	0	+x	x
H⁺	0	+x	x

For CH₃COOH ⇌ CH₃COO⁻ + H⁺, Kc = [CH₃COO⁻][H⁺] / [CH₃COOH] = x² / (0.10 - x) = 1.8 × 10⁻⁵

Reaction Quotient (Q) vs Equilibrium Constant (K)

The reaction quotient Q has the same form as K but uses current (non-equilibrium) concentrations. Comparing Q to K predicts the direction the reaction will shift.

Q < K

Reaction proceeds forward (toward products). System shifts right.

→

Q = K

System is at equilibrium. No net change occurs.

⇌

Q > K

Reaction proceeds in reverse (toward reactants). System shifts left.

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How to Calculate Equilibrium Constants

Equilibrium constants are calculated from experimental data or derived from thermodynamic properties. The process depends on whether you're working with concentrations or pressures.

🔬 Step-by-Step Process

1. Write the Balanced Equation

Ensure the reaction is balanced: aA + bB ⇌ cC + dD

2. Set Up the Expression

Kc = [C]^c[D]^d / [A]^a[B]^b (products over reactants, raised to stoichiometric coefficients)

3. Use Equilibrium Concentrations

Only equilibrium concentrations are used. Pure solids and liquids are excluded (activity = 1).

4. Calculate K

Substitute values and calculate. K is dimensionless but has units implied by concentration units.

When to Use Equilibrium Constants

Equilibrium constants are fundamental in chemistry for predicting reaction behavior, designing processes, and understanding chemical systems.

🏭

Industrial Chemistry

Optimize reaction conditions for maximum yield in chemical manufacturing.

Haber process (ammonia)
Contact process (sulfuric acid)
Ostwald process (nitric acid)

🧪

Analytical Chemistry

Determine concentrations, predict solubility, and analyze complex formation.

Acid-base equilibria
Complex ion formation
Solubility products

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Biochemistry

Understand enzyme kinetics, protein folding, and metabolic pathways.

Enzyme-substrate binding
Protein-ligand interactions
Oxygen binding to hemoglobin

Common Equilibrium Reactions

Reaction	Equation	Kc (25°C)	Type
Haber Process	N₂(g) + 3H₂(g) ⇌ 2NH₃(g)	3.50e+8	gas
Water Gas Shift	CO(g) + H₂O(g) ⇌ CO₂(g) + H₂(g)	5.000	gas
Dissociation of Acetic Acid	CH₃COOH(aq) ⇌ CH₃COO⁻(aq) + H⁺(aq)	1.80e-5	aqueous
Formation of HI	H₂(g) + I₂(g) ⇌ 2HI(g)	54.30	gas
Dissociation of Water	H₂O(l) ⇌ H⁺(aq) + OH⁻(aq)	1.00e-14	aqueous
Decomposition of N₂O₄	N₂O₄(g) ⇌ 2NO₂(g)	0.2110	gas
Formation of NO	N₂(g) + O₂(g) ⇌ 2NO(g)	4.10e-31	gas
Dissociation of H₂S	H₂S(aq) ⇌ H⁺(aq) + HS⁻(aq)	9.10e-8	aqueous

Le Chatelier's Principle

When a system at equilibrium is disturbed, it responds to minimize the disturbance. This principle helps predict how changes in concentration, pressure, or temperature affect equilibrium position.

📈 Effect of Changes

• Increase reactant: Shifts right (toward products)
• Increase product: Shifts left (toward reactants)
• Increase pressure: Shifts toward fewer gas moles
• Increase temperature: Shifts toward endothermic direction
• Catalyst: No effect on equilibrium position

⚖️ Important Notes

• K changes only with temperature (for a given reaction)
• Changes in concentration/pressure shift position but not K
• Pure solids/liquids don't affect equilibrium position
• Dilution shifts toward side with more moles
• Equilibrium is dynamic (forward = reverse rate)

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