UNIT A_ Thermochemical Changes

UNIT A: Thermochemical Changes

INDEX

  • Topic 1: Enthalpy Change

  • Topic 2: Explaining Chemical Changes

Topic 1: Enthalpy Change

Learning Objectives

  • Explain energy in hydrocarbons originated from the sun.

  • Classify reactions as endothermic or exothermic.

  • Identify water and carbon dioxide's roles in photosynthesis and cellular respiration.

  • Analyze heat transfer using Q=mcΔt.

  • Define enthalpy and molar enthalpy.

  • Use calorimetry data to determine enthalpy changes.

  • Interpret ΔH notation for energy changes in reactions.

  • Write balanced equations for reactions involving energy changes.

  • Explain and apply Hess’ law to calculate energy changes.

  • Predict enthalpy changes using standard enthalpies of formation.

Thermochemistry Overview

  • Based on the First Law of Thermodynamics: energy is converted, not created or destroyed.

  • Photosynthesis: Radiant energy transforms into chemical energy (stored in bonds of carbohydrates).

  • Chemical Energy: Converted into usable forms through various processes.

Key Principles

  • Second Law of Thermodynamics: Energy conversion is never 100% efficient, leading to heat loss.

  • Heat Transfer in Reactions:

    • Endothermic Reactions: Absorb thermal energy (system gains energy, surroundings lose).

    • Exothermic Reactions: Produce thermal energy (system loses energy, surroundings gain).

Endothermic vs. Exothermic Reactions

Endothermic

  • Example: Photosynthesis

    • Plants capture energy and store it in glucose.

    • System temperature decreases as surrounding temperature increases.

Exothermic

  • Examples: Cellular Respiration, Combustion

    • Release energy into surroundings (noticeably warming the surroundings).

Calculating Thermal Energy Change

Factors Influencing Thermal Energy (Q)

  • Mass (m): of the substance undergoing reaction (grams).

  • Specific Heat Capacity (c): Amount of energy to raise temperature (J/(g•°C)).

  • Temperature Change (Δt): Overall change in temperature (°C).

Specific Heat Capacity (Example)

  • Water: 4.19 J/(g•°C) (Energy to raise 1g of water by 1°C).

Applying Q=mcΔt

Sample Calculation

  • Calculate thermal change when 115g of water is heated from 19.6°C to 98.8°C.

  • Determine whether this change is endothermic or exothermic.

Enthalpy

Definition

  • Total potential (stored in bonds, intermolecular forces) and kinetic energy (electron movements) of a chemical system.

Enthalpy Change

  • Indirectly calculated from thermal energy change (Q).

  • Example: Burning candle shows indirect measurement of combustion enthalpy change.

Measuring Enthalpy Change (Calorimetry)

  • Method for indirectly measuring energy change via isolated systems.

Types of Calorimeters

  • Simple Calorimeter

  • Flame Calorimeter

  • Bomb Calorimeter

Understanding ΔH

  • Negative ΔH indicates exothermic (energy lost), positive indicates endothermic (energy gained).

  • Example: N2(g) + 3 H2(g) --> 2 NH3(g), ∆H = -90 kJ implies energy relationship scales with coefficients.

Molar Enthalpy

Definition

  • Change in enthalpy per unit of chemical amount denoted as ΔrH (in J/mol).

Key Calculations involving Molar Enthalpy

  1. Convert between enthalpy and molar enthalpy.

  2. Calculate enthalpy change per specific mass based on molar values.

  3. Use calorimetry data for molar enthalpy calculations.

Standard Molar Enthalpies of Formation

Formula

  • ΔH°r = ∑nΔfH° (products) - ∑nΔfH° (reactants).

  • Theoretical value upon element combination for compound formation.

Energy Change in Chemical Reactions

Communicating Energy Changes

  • Ex: CO(g) + NO2(g) --> CO2(g) + NO(g) ∆H = -224.9 kJ illustrates exothermic interpretation.

Thermochemical Equations

  • Incorporate ΔH into balanced equations to express energy changes.

Example of Endothermic vs. Exothermic

  • H2(g) + I2(g) + 53.0 kJ ---> 2 HI(g) with energy uptake.

Potential Energy Diagrams

Characteristics

  • Exothermic: Lower potential energy in products.

  • Endothermic: Higher potential energy in products.

  • Example: Combustion of magnesium (ΔH = -601.6 kJ).

Hess’s Law

Overview

  • Measures ΔH for complex reactions by summing individual reaction changes.

Steps

  1. Add overall products and reactants.

  2. Simplify by canceling identical terms.

  3. Total ΔH from contributing reactions.

Activation Energy

Definition

  • Minimum energy needed to initiate a chemical reaction.

  • Factor for all reactions, indicating the energy needed to break bonds.

Role of Catalysts

  • Lower activation energy, increasing reaction rates without altering net enthalpy changes.

  • Enzymes in nature (like chlorophyll in photosynthesis) catalyze biological reactions.

Conclusion

  • Understanding thermochemical changes aids in predicting reaction behavior, energy calculations, and application of catalysts.