Calculate the equilibrium constant (Kc) from balanced chemical equations and equilibrium concentrations. Also convert between Kc and Kp, compare with reaction quotient Q, and determine the direction of reaction.
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Kc — Equilibrium Constant Calculator
aA + bB ⇌ cC + dD | Kc = [C]c[D]d / [A]a[B]b
Select a common reaction or enter custom coefficients. Then input the equilibrium concentrations (M) for each species.
aA + bB ⇌ cC + dD
🧪 Custom Stoichiometric Coefficients
📊 Equilibrium Concentrations (M)
Enter the equilibrium concentrations in mol/L (M). Leave blank or 0 if a species is not present or coefficient is 0.
🌡️ Temperature (for Kp conversion)
Kc (Equilibrium Constant)
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In terms of concentration
Kp Value
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In terms of partial pressure
Δn (Change in moles of gas)
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Moles of products − reactants
📝 Step-by-Step Calculation
🔄 Reaction Quotient Q & Reaction Direction
Enter initial (non-equilibrium) concentrations to calculate Q and determine reaction direction.
Reaction Quotient (Q)
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Using initial concentrations
Kc
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Equilibrium constant
📝 Q Calculation Steps
Kp ↔ Kc Conversion
Kp = Kc (RT)Δn
Convert between Kp and Kc. Enter the known equilibrium constant, temperature, and the change in moles of gas (Δn).
moles of gaseous products − reactants
Result
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📝 Conversion Steps
Important Notes
R = 0.08206 L·atm / (mol·K) — the ideal gas constant
Temperature must be in Kelvin for the formula. If using °C, add 273.15.
Δn = (moles of gaseous products) − (moles of gaseous reactants)
If Δn = 0, then Kp = Kc (no conversion needed)
Kp and Kc are related but have different units (pressure vs concentration)
🧪 Equilibrium Constant Examples
Example 1: H₂ + I₂ ⇌ 2HI
Problem: At 445°C, the equilibrium concentrations for the reaction H₂(g) + I₂(g) ⇌ 2HI(g) are [H₂] = 0.50 M, [I₂] = 0.50 M, and [HI] = 1.0 M. Calculate Kc.
Equilibrium expression: Kc = [HI]² / ([H₂][I₂])
Substitute values: Kc = (1.0)² / (0.50 × 0.50)
Calculate: Kc = 1.0 / 0.25 = 4.0
Interpretation: Kc = 4.0 means at equilibrium, product concentration is favored. Since Kc > 1, the reaction favors products at this temperature.
Example 2: N₂O₄ ⇌ 2NO₂
Problem: For the decomposition of dinitrogen tetroxide, N₂O₄(g) ⇌ 2NO₂(g), at 100°C the equilibrium concentrations are [N₂O₄] = 0.20 M and [NO₂] = 0.60 M. Find Kc.
Interpretation: Kc ≈ 0.93 indicates that at 500°C, the equilibrium mixture contains significant amounts of both reactants and products.
Example 4: Using Q to Predict Direction
Problem: For H₂(g) + I₂(g) ⇌ 2HI(g) (Kc = 4.0 at 445°C), initial concentrations are [H₂] = 0.10 M, [I₂] = 0.10 M, [HI] = 0.50 M. Which way will the reaction proceed?
Direction: Since Q > Kc, the reaction proceeds in the reverse direction (toward reactants) to reach equilibrium.
Check: At equilibrium, [HI] will decrease and [H₂] and [I₂] will increase until Q = Kc.
Understanding Chemical Equilibrium
Chemical equilibrium is a state in a reversible reaction where the rates of the forward and reverse reactions are equal, and the concentrations of reactants and products remain constant over time. The equilibrium constant (Kc or Kp) quantifies the position of equilibrium.
Equilibrium Constant Formula
Kc = [C]c[D]d / [A]a[B]b
For the general reaction: aA + bB ⇌ cC + dD
Kp = Kc (RT)Δn
Where R = 0.08206 L·atm/(mol·K), T = temperature in Kelvin, Δn = (gaseous products − gaseous reactants)
What the Equilibrium Constant Tells You
1
Kc >> 1 (Large, e.g., 10³): The equilibrium strongly favors products. The forward reaction proceeds almost to completion. At equilibrium, there is much more product than reactant.
2
Kc << 1 (Small, e.g., 10⁻³): The equilibrium strongly favors reactants. Very little product is formed. The forward reaction barely proceeds.
3
Kc ≈ 1 (0.1 to 10): Significant amounts of both reactants and products are present at equilibrium. The reaction is reversible and neither direction is strongly favored.
Reaction Quotient Q vs Kc
The reaction quotient Q has the same mathematical form as Kc, but uses initial or current concentrations rather than equilibrium concentrations. Comparing Q to Kc tells you the direction the reaction will proceed:
Q
If Q < Kc: The reaction proceeds forward (toward products). More products will form until Q = Kc.
Q
If Q > Kc: The reaction proceeds in reverse (toward reactants). More reactants will form until Q = Kc.
Q
If Q = Kc: The system is already at equilibrium. No net change occurs.
Le Chatelier's Principle
Le Chatelier's principle states that if a dynamic equilibrium is disturbed by changing the conditions, the system will shift its equilibrium position to counteract the change. Here's how different factors affect equilibrium:
📈 Concentration Changes
Adding more reactant shifts equilibrium toward products. Removing a product also shifts toward products. The equilibrium constant Kc remains unchanged — only the position shifts.
🌡️ Temperature Changes
For exothermic reactions (ΔH < 0), increasing T shifts equilibrium toward reactants and changes Kc. For endothermic reactions (ΔH > 0), increasing T shifts toward products and also changes Kc. Temperature is the only factor that changes the value of K.
⚡ Pressure/Volume Changes
Increasing pressure (decreasing volume) shifts equilibrium toward the side with fewer moles of gas. If Δn = 0, pressure has no effect on equilibrium position. Kc remains unchanged.
🧪 Catalysts
A catalyst speeds up both forward and reverse reactions equally. It helps the system reach equilibrium faster but does not change the equilibrium position or the value of Kc.
Common Equilibrium Reactions
Reaction
Kc Expression
Typical Kc (at 25°C)
H₂ + I₂ ⇌ 2HI
[HI]²/([H₂][I₂])
50.5 (at 445°C)
N₂ + 3H₂ ⇌ 2NH₃
[NH₃]²/([N₂][H₂]³)
6.8 × 10⁵ (at 25°C)
N₂O₄ ⇌ 2NO₂
[NO₂]²/[N₂O₄]
0.212 (at 100°C)
2SO₂ + O₂ ⇌ 2SO₃
[SO₃]²/([SO₂]²[O₂])
4.0 × 10²⁴ (at 25°C)
PCl₅ ⇌ PCl₃ + Cl₂
[PCl₃][Cl₂]/[PCl₅]
0.0415 (at 250°C)
CO + 2H₂ ⇌ CH₃OH
[CH₃OH]/([CO][H₂]²)
14.5 (at 100°C)
Key Rules for Equilibrium Constants
Pure solids and liquids are not included in the equilibrium expression (their activity is 1).
Kc is temperature-dependent — changing the temperature changes the value of Kc.
Kc is unitless in thermodynamic contexts, but in practice we use concentrations in mol/L.
If you reverse a reaction, the new Kc = 1 / (original Kc).
If you multiply a reaction by n, the new Kc = (original Kc)ⁿ.
If you add two reactions, the new Kc = Kc₁ × Kc₂.
Equilibrium Calculator Features
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Kc & Kp Calculation
Calculate Kc from equilibrium concentrations, and automatically convert to Kp using the temperature and Δn. Full step-by-step working shown.
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Q vs Kc Comparison
Enter initial concentrations to compute the reaction quotient Q and compare it with Kc to determine which direction the reaction must proceed to reach equilibrium.
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Common Reactions Library
Choose from 10+ pre-loaded common equilibrium reactions including the Haber process, contact process, and more. Or enter custom coefficients.
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Temperature Conversion
Built-in Kp ↔ Kc converter with temperature support. Enter °C or K and get instant conversion using Kp = Kc(RT)^Δn.
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Responsive design works perfectly on phones, tablets, and desktops for studying chemistry on the go.
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Our Equilibrium Constant Calculator simplifies calculating Kc and Kp from experimental concentration data with step-by-step solutions for chemistry students and professionals.
Whether studying for AP Chemistry or working in industry, this calculator helps determine equilibrium constants and predict reaction behavior quickly and accurately.
⚠️ Important Note: This calculator is for educational and reference purposes. Results should be verified with established chemical data for critical applications. Equilibrium constants are highly temperature-dependent — always note the temperature when reporting Kc or Kp values.