Thermodynamics in Chemical Reactions

Chapter 7.8 Review of Thermodynamics in Chemical Reactions

Overview

• Objective: Review the relationships involving the change in enthalpy (ΔH) of a reaction, noting how these changes depend on the specific reaction conditions.

• Key Theme: The alteration of reactants or products leads to a corresponding change in the enthalpy, with measurable quantitative relationships.

Key Concepts

1. Change in Enthalpy

• Definition of ΔH: Change in enthalpy is specific to each reaction and varies based on reactants and products involved.

• Reagents and Products: If you alter them, the change in enthalpy must be recalculated.

2. Direction of Reaction

• Changing the direction of the reaction modifies ΔH accordingly.

  • Example: Multiplying the reaction coefficients alters ΔH in a proportional manner.

  • If a reaction is multiplied by 2, then ΔH becomes 2 * ΔH.

3. Quantitative Relationships in Reactions

• Three major relationships regarding ΔH:

  1. If a chemical equation is multiplied by a factor, ΔH is also multiplied by that same factor.

  2. If a reaction is reversed, ΔH changes sign (positive to negative and vice versa).

  3. Hess's Law can be applied where multiple reactions are summed to derive ΔH for a target reaction.

Example of Multiplying Reaction

• Hypothetical Reaction:

  • 
    $A + 2B <br>ightarrow C$

• If coefficients are multiplied by 2:

  • Resulting Reaction:
    $2A + 4B <br>ightarrow 2C$

  • New ΔH = 2 * ΔH(1)

4. Reversing Reactions and Sign Change

• If:
  $A + 2B <br>ightarrow C$ gives ΔH(1)

• Then reversing:
  $C <br>ightarrow A + 2B$ results in ΔH(2) = -ΔH(1)

5. Hess's Law

• Definition: If a chemical reaction can be constructed from a series of steps,
  then the total ΔH for the overall process is the sum of ΔH values for the steps.

• Steps to Apply Hess's Law:

  1. Identify all steps needed to form the desired equation.

  2. Ensure terms that appear on both sides cancel out when summing ΔH's.

  3. Formulate the equation such that reactants and products align appropriately.

  • Example:

    • Step 1:
      $A + 2B <br>ightarrow C$

    • ΔH(1)

    • Step 2:
      $C <br>ightarrow 2D$

    • ΔH(2)

    • Net Reaction:
      $A + 2B <br>ightarrow 2D$

  • Resulting ΔH = ΔH(1) + ΔH(2)

6. Energy Diagrams in Thermodynamics

• Energy diagrams illustrate the total change in enthalpy during reactions.

  • They visually depict the relationship between reactants and products, including exothermic and endothermic processes.

Example Problems and Applications

1. Hypothetical Example: Reaction and ΔH Calculation

• Consider the reaction:
  $2A + B <br>ightarrow C$

• Given a ΔH = 122 joules.

• For the reaction: $2C ightarrow 4A + 2B$

  • Calculating ΔH gives -244 joules after applying Hess's Law by reversing and adjusting coefficients.

2. Industrial Applications

• Production of hydrogen gas from carbon and water:
  $C + H<em>2O ightarrow CO + H</em>2$

• Employed in industrial methods of hydrogen production, highlighting practical applications of the concepts.

3. Determining Unknown ΔH Values

• Strategies include:

  1. Using known reactions to derive ΔH for unknown reactions.

  2. Analyzing coefficients and ensuring they match the desired reaction's coefficient arrangements.

Standard State and Calculating ΔH

1. Standard State for Reagents

• Standard state for gases: pure gas at 1 atmosphere pressure.

• Standard state for liquids: pure liquid at 1 atmosphere and designated temperature (typically 25°C).

• Standard states simplify the process for calculating ΔH values further.

2. Standard Enthalpy of Formation (ΔHf°)

• Definition: Change in enthalpy when one mole of a compound is formed from its constituent elements in their standard states.

• For pure elements, ΔHf° = 0 (e.g., C, H2).

3. Calculation Example for ΔHf°

• Formation of methane: $C (graphite) + 2H<em>2 (gas) ightarrow CH</em>4 (gas)$

  • ΔHf° = -74.6 kJ/mol

4. Practical Implications of Understanding ΔHf°

• Understanding these principles allows predictions of energy changes in chemical reactions.

• Important for industrial applications where energy efficiency and reaction feasibility are critical.

Conclusion

• Mastering the relationships among reaction coefficients, ΔH, and Hess's Law enriches understanding of thermodynamics in chemical applications.

• Students are encouraged to practice problems involving Hess's Law from the textbook for additional understanding and retention.