Book cover for Chemistry: The Molecular Nature of Matter

Chemistry: The Molecular Nature of Matter

Neil D. Jespersen, James E. Brady, Alison Hyslop

ISBN #9781118413920

7th Edition

3,064 Questions

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53,557 Students Helped

Homework Questions

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Summary
Learning Objectives
Key Concepts
Example Problems
Explanations
Common Mistakes

Summary

This chapter on Chemical Equilibrium covers the state where forward and reverse reactions occur at equal rates, resulting in constant concentrations. It highlights the role of equilibrium laws formulated via the mass action expression and introduces key constants Kc and KP, along with their interconversion depending on the change in gas moles (?ng) and temperature. The use of ICE tables simplifies the computation of equilibrium concentrations, and Le Châtelier’s principle provides qualitative insights into how a system responds to disturbances. Mastery of these concepts is essential for effective laboratory analysis and optimizing industrial chemical processes.

Learning Objectives

1

Describe the concept of chemical equilibrium and differentiate between dynamic equilibrium and static conditions.

2

Explain and apply equilibrium laws, including the mass action expression and equilibrium constants (Kc and KP).

3

Utilize ICE tables to simplify equilibrium calculations and determine concentration changes.

4

Analyze the impact of system disturbances using Le Châtelier’s principle.

5

Connect equilibrium concepts to laboratory analyses and industrial chemical process optimization.

Key Concepts

CONCEPT

DEFINITION

Chemical Equilibrium

A state in a reversible reaction where the forward and reverse reaction rates are equal, leading to constant concentrations of reactants and products.

Dynamic Equilibrium

A condition in which the processes of the forward and reverse reactions continue to occur at equal rates, resulting in no net change in concentration.

Mass Action Expression

An equation relating the concentrations or partial pressures of reactants and products at equilibrium, typically written in the form Kc = [products]^coefficients / [reactants]^coefficients.

Equilibrium Constant (Kc, KP)

A numerical value that quantifies the ratio of the concentration or pressure of products to that of reactants at equilibrium; Kc is based on concentrations while KP is based on pressures.

∆ng

The change in the number of moles of gas in a reaction, which affects the conversion between Kc and KP.

ICE Table

A tabular method (Initial, Change, Equilibrium) used to organize and simplify equilibrium concentration calculations.

Le Châtelier’s Principle

A principle stating that when a system at equilibrium is disturbed, it will shift in the direction that counteracts the disturbance to re-establish equilibrium.

Example Problems

Example 1

Sketch a graph showing how the concentrations of the reactant $A$ and product $P$ of a typical chemical reaction $(A \rightleftharpoons P)$ vary with time during the course of the reaction. Assume that no products are present at the start of the reaction. Indicate on the graph where the system has reached equilibrium.

Example 2

Using black circles to represent $A$ and open circles to represent $P$, show the relative amounts of $A$ and $P$ in beakers at two points before equilibrium is reached and two points after equilibrium is reached in your answer to Question 14.1. The original beaker looks like this.

Example 3

Repeat the exercise in Question 14.1 but this time start with a mixture that is $25 \%$ product and $75 \%$ reactant.

Example 4

Using black circles to represent $A$ and open circles to represent $P$, show the relative amounts of $A$ and $P$ in beakers at two points before equilibrium is reached and two points after equilibrium is reached in your answer to Question 14.3. The original beaker looks like this.

Example 5

What meanings do the terms reactants and products have when describing a chemical equilibrium?

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Step-by-Step Explanations

QUESTION

How do you determine the equilibrium concentrations for a reversible reaction using an ICE table?

STEP-BY-STEP ANSWER:

Step 1: Write the balanced chemical equation and the corresponding mass action expression for the equilibrium constant.
Step 2: Construct an ICE table by listing the initial concentrations, the changes (using a variable such as x) as the reaction proceeds, and the final equilibrium concentrations.
Step 3: Substitute the equilibrium concentrations expressed in terms of x into the mass action expression.
Step 4: Solve the resulting equation for x to find the changes in concentrations.
Step 5: Calculate the final equilibrium concentrations by substituting the value of x back into the ICE table.
Final Answer:

Using ICE Tables

QUESTION

How do you convert between equilibrium constants based on concentrations (Kc) and pressures (KP) for a reaction?

STEP-BY-STEP ANSWER:

Step 1: Write the mass action expressions for both Kc and KP for the given reaction.
Step 2: Determine the change in moles of gas (∆ng) for the reaction by subtracting the total moles of gaseous reactants from those of gaseous products.
Step 3: Use the relation KP = Kc(RT)^∆ng, where R is the gas constant and T is the temperature in Kelvin.
Step 4: Substitute the known values for R, T, and ∆ng to calculate KP from the given Kc (or vice versa).
Final Answer:

Converting Between Kc and KP

QUESTION

What is the effect of increasing the concentration of a reactant on a system at equilibrium?

STEP-BY-STEP ANSWER:

Step 1: Recognize that an increase in the concentration of a reactant disturbs the equilibrium by increasing the forward reaction rate.
Step 2: According to Le Châtelier’s principle, the system will shift the equilibrium to reduce the added reactant, favoring the production of products.
Step 3: Predict that the concentrations of products will increase as the system counteracts the disturbance.
Step 4: Validate the prediction by confirming that the mass action expression remains satisfied at the new equilibrium state.
Final Answer:

Applying Le Châtelier’s Principle

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Common Mistakes

  • Confusing static equilibrium with dynamic equilibrium, not recognizing that reactions are still occurring at equilibrium.
  • Neglecting the effect of temperature and gas mole changes (?ng) in converting between Kc and KP.
  • Incorrectly setting up ICE tables, such as misidentifying initial amounts or changes in concentration.
  • Overlooking the direction of the shift in equilibrium when applying Le Châtelier’s principle under various disturbances.