How to Solve Law of Chemical Equilibrium Numericals for Exams
The law of chemical equilibrium numericals topic is fundamental for mastering JEE Main Chemistry computations. It connects conceptual understanding with formula-based applications, allowing students to predict reaction directions, calculate equilibrium constants, and solve practical exam problems. By consistently applying the law of mass action and using relevant equations, you can confidently approach questions on reversible reactions and dynamic equilibrium.
Law of Chemical Equilibrium Numericals: Core Ideas and Applications
The law of chemical equilibrium states that for a reversible reaction at equilibrium, the rate of the forward reaction equals the rate of the backward reaction. At this point, the ratio of the concentrations of products to reactants, each raised to their respective coefficients, is constant (Kc or Kp), at a fixed temperature. This law helps students analyze exam-level questions and understand how chemical systems reach stability.
Let us consider the classic example:
H2(g) + I2(g) ⇌ 2HI(g)
At equilibrium, the equilibrium constant is: Kc = [HI]2 / ([H2][I2])
For calculations involving gases, you may use Kp = Kc(RT)Δng, where Δng is the change in moles of gas.
Key Formulae for Law of Chemical Equilibrium Numericals
| Formula | Meaning & Use |
|---|---|
| Kc = [C]c[D]d / [A]a[B]b | Equilibrium constant in terms of molarity (C, D: products; A, B: reactants) |
| Kp = (PC)c(PD)d / (PA)a(PB)b | Equilibrium constant in terms of partial pressures |
| Kp = Kc(RT)Δng | Relates Kp and Kc for gaseous reactions; R = 0.0821 L·atm·mol-1·K-1 |
| Q = [Products]/[Reactants] | Reaction quotient – compare with Kc to predict shift in direction |
To solve law of chemical equilibrium numericals, always write the balanced equation, set up the equilibrium expression, substitute values, and check for consistency with dimensional analysis.
Step-by-Step Numerical Problems on Law of Chemical Equilibrium
Let’s practice classic types of chemical equilibrium numericals using core concepts and precise calculation methods. Follow the outlined method for full marks.
- Write the balanced chemical equation for the reversible reaction.
- Express the equilibrium law using correct concentrations or pressures.
- Set up an ICE (Initial-Change-Equilibrium) table if required for unknowns.
- Substitute all known values at equilibrium.
- Calculate Kc or Kp as per the question.
- Compare the value of Q to Kc to predict direction, if asked.
Example 1: Given the reaction H2(g) + I2(g) ⇌ 2HI(g) at 717 K, with [H2] = 0.2 mol L-1, [I2] = 0.2 mol L-1 and [HI] = 0.6 mol L-1. Kc = 48. Find Q and predict the direction.
- Q = (0.6)2 / (0.2 × 0.2) = 0.36 / 0.04 = 9
- Since Q < Kc, reaction proceeds in the forward direction.
Example 2: For PCl5(g) ⇌ PCl3(g) + Cl2(g), 1 mole PCl5 is heated in 1L. Equilibrium [Cl2] = 0.6 mol L-1. Calculate Kc.
- [PCl3]eq = 0.6 mol L-1
- [PCl5]eq = 1 - 0.6 = 0.4 mol L-1
- Kc = [PCl3][Cl2] / [PCl5] = (0.6 × 0.6) / 0.4 = 0.9
Practice Worksheet: Law of Chemical Equilibrium Numericals
| Question | Answer | Concept Tested |
|---|---|---|
| For N2O4(g) ⇌ 2NO2(g), Kc = 0.21 at 373 K. If [N2O4] = 0.125 and [NO2] = 0.5 mol/L, find Q and the direction of reaction. | Q = (0.5)2 / 0.125 = 2; Q > Kc, shifts left. | Using Q vs Kc |
| At 298 K for N2(g) + 3H2(g) ⇌ 2NH3(g), Kp = 8.19 × 102. Is Kp > Kc, or vice versa? | Δng = -2; Kp < Kc | Relation between Kp and Kc |
| 1 mol CH4, 2 mol H2S, 1 mol CS2 and 2 mol H2 in 0.5 L for CH4 + 2H2S ⇌ CS2 + 4H2, Kc = 4 × 10-2. Find Q and reaction direction. | [CH4]=2, [H2S]=4, [CS2]=2, [H2]=4 Q = 16; Q > Kc, reverse. |
Initial Q value vs Kc |
Practice with a timer and check your dimensional consistency. For more solved samples, Vedantu offers additional chemical equilibrium worksheets and mock tests designed for the exam environment.
Common Mistakes and Pitfalls in Chemical Equilibrium Numericals
- Forgetting to use equilibrium concentrations (not initial ones) in the Kc and Kp expressions.
- Ignoring the correct exponent or coefficient for each species in the formula.
- Missing out the units or using wrong volume for calculating molarity, especially for gaseous systems.
- Confusing Kc (concentration) and Kp (pressure), especially when Δng ≠ 0.
- Assuming solids and pure liquids appear in the equilibrium law (they do not).
- Mixing up the concepts of reaction quotient Q and equilibrium constant (K).
A common trap is to plug in initial instead of equilibrium values, causing wrong answers. Always build an ICE table and double-check each step.
Shortcuts and Exam Tips for Law of Chemical Equilibrium Numericals
- For Δng=0 (equal moles gas each side), Kp = Kc.
- Reaction proceeds forward if Q < K; backward if Q > K.
- Only include gases and solutions in equilibrium expressions; omit solids, pure liquids.
- Practice calculation speed with varied equilibrium problems.
- Know R = 0.0821 L·atm·mol-1·K-1 for Kp–Kc conversions.
- Re-check direction by comparing Q and Kc in every scenario.
By mastering these tactics, you can save precious time in the JEE Main exam and maximize marks on chemical equilibrium numericals. For instant practice, explore the dedicated numerical problem sets and high-yield revision material on Vedantu.
Deeper Connections: More Resources for JEE Main Success
- Read about Le Chatelier’s Principle for equilibrium shifts.
- Master the Law of Mass Action for more advanced theoretical explanations.
- See the difference between ionic equilibrium and chemical equilibrium numericals.
- Use the revision notes for equilibrium to reinforce theory before attempting tougher problems.
- Access all JEE Chemistry mock tests for practice under timed conditions.
Mastering law of chemical equilibrium numericals connects calculation and chemical logic, a core JEE Main skill. Use Vedantu resources and build a systematic approach for consistent top scores.
FAQs on Law of Chemical Equilibrium Numericals: Step-by-Step Solutions
1. What is the law of chemical equilibrium?
The law of chemical equilibrium states that for a reversible reaction at equilibrium, the ratio of the concentrations (or partial pressures) of products to reactants, raised to the power of their stoichiometric coefficients, remains constant at a given temperature.
This constant is called the equilibrium constant (Kc or Kp).
Key points:
- Applies to reversible reactions
- Ratio is determined by the balanced chemical equation
- Only changed by temperature, not by catalyst or pressure (unless gases are involved)
2. How do I solve numericals based on the law of chemical equilibrium?
To solve law of chemical equilibrium numericals, follow a clear step-wise method:
Steps:
- Write the balanced chemical equation.
- Set up the equilibrium expression based on the equation (Kc and/or Kp).
- List initial concentrations or moles given.
- Apply changes to reach equilibrium (use ICE or RICE table if needed).
- Substitute values into the equilibrium law formula.
- Solve for required unknown (K value, concentration, pressure, etc).
3. Can you provide a worksheet or PDF of solved equilibrium problems?
Yes, many reputable educational sites provide law of chemical equilibrium numerical worksheets and PDFs with detailed solutions. These worksheets include:
- Step-by-step solved examples for Kc and Kp
- Practice questions with answer keys
- MCQs, short and long answer types
- JEE Main, NEET, and class 11 exam pattern coverage
4. What are some examples of chemical equilibrium in real life?
Chemical equilibrium occurs in many real-life scenarios, demonstrating how reactions balance reactants and products over time.
Examples:
- Carbon dioxide in soft drinks: CO2 dissolved in water forms equilibrium with gaseous CO2.
- Haber Process: Industrial synthesis of ammonia (N2 + 3H2 ⇌ 2NH3).
- Blood Oxygenation: O2 binding and release with hemoglobin is a reversible equilibrium.
- Buffer systems in biology: Maintain pH through weak acid/base equilibria.
5. How do I calculate Kc and Kp values in equilibria?
To calculate Kc (equilibrium constant for concentration) and Kp (for partial pressure):
For Kc:
- Use the equilibrium concentrations (mol/L) in the expression based on stoichiometry.
- Use partial pressures of gases (in atm or bar) for values at equilibrium.
6. What is the difference between Kc and Kp in chemical equilibrium?
The main difference between Kc and Kp is the type of values used:
- Kc uses molar concentrations (mol/L) for both reactants and products.
- Kp uses partial pressures (atm, bar) for gaseous reactants and products.
- For reactions involving gases, Kc and Kp are related by Kp = Kc(RT)Δn.
7. Why do students often confuse Kc and Kp, and how can I avoid this?
Many students confuse Kc and Kp because both are equilibrium constants but use different units and are applied to different systems.
Avoid confusion by:
- Checking the phase: Use Kc for solutions (aq), Kp for gases.
- Remembering: Kc – concentration (mol/L), Kp – pressure (atm/bar).
- Using the relation Kp = Kc(RT)Δn for conversions.
8. What are some common calculation errors in chemical equilibrium numericals?
Some frequent calculation mistakes in chemical equilibrium numericals include:
- Incorrectly writing the equilibrium expression (forgetting exponents or omitting species).
- Mixing up initial and equilibrium values.
- Using incorrect units for Kc or Kp.
- Ignoring Δn when converting between Kc and Kp.
- Simple arithmetic errors or misusing formulas.
9. Can equilibrium constants be negative or have no units?
Equilibrium constants (Kc and Kp) are always positive; they can never be negative.
- The units of equilibrium constants depend on the reaction and are determined by the change in moles (Δn).
- For some balanced reactions (Δn = 0), Kc and Kp may become dimensionless (no units).
10. Which equilibrium law shortcuts are best for competitive exams?
For JEE Main, NEET, and other competitive exams, use these quick tips:
- Use the ICE/RICE table for organizing changes quickly.
- Memorize the relationship: Kp = Kc(RT)Δn
- For strong acids/bases, assume complete dissociation unless specified.
- Use significant value filters–ignore very small terms when allowed by the question.
- Avoid calculation traps by checking stoichiometry before plugging values.
