Notes on Chapter 10: Energy Changes in Chemical Reactions

Chapter 10: Energy Changes in Chemical Reactions

10.1 Energy and Energy Changes

• Definition of System and Surroundings:
    - The system is the part of the universe that is of specific interest, typically defined as the substances involved in chemical and physical changes.
    - The surroundings refer to everything outside the system.

• Thermochemistry:
    - Defined as the study of heat (transfer of thermal energy) in chemical reactions.
    - Heat is the transfer of thermal energy, which can either be absorbed or released during a process.
    - SI Unit for Energy: The Joule (J).
    - Caloric Conversion:
      - 1 calorie = amount of heat required to raise 1 g of water by 1°C.
      - 1 Calorie (Cal) = 4.184 Joules (J)
      - 1 Cal = 1000 calories (cal).

• Exothermic Process:
    - Occurs when heat is transferred from the system to the surroundings.
    - Example: Hot packs, combustion reactions.
    - Observation: “Feels hot!”

• Endothermic Process:
    - Occurs when heat is absorbed from the surroundings by the system.
    - Example: Photosynthesis, melting ice.
    - Observation: “Feels cold.”

10.2 Introduction to Thermodynamics

• Definition: Thermodynamics is the study of interconversion of heat and other types of energy.

Types of Systems

1. Open System:
      - Can exchange both mass and energy with surroundings.
      

2. Closed System:
      - Allows energy transfer but does not allow mass transfer.
      

3. Isolated System:
      - Does not exchange either mass or energy with the surroundings.

States and State Functions

• State Functions: Properties determined by the state of the system, independent of the process taken to achieve that state. Examples include:
    - Energy
    - Pressure
    - Volume
    - Temperature

The First Law of Thermodynamics

• Statement: Energy cannot be created or destroyed; it can only be transformed from one form to another.

• Mathematical Expression: The change in internal energy (ΔU) is given by:
    - $ext{ΔU = q + w}$,
    - Where:
      - $ext{q}$ = heat added to the system,
      - $ext{w}$ = work done on the system.

Work and Heat

• Work (w): The energy transfer via mechanical means.
    - Sign Conventions:
      - For Heat (q):
        - q > 0 for endothermic processes (heat absorbed) for the system
        - q < 0 for exothermic processes (heat released)by the system.     - For Work (w):       - w > 0 if work is done on the system,
        - w < 0 if work is done by the system.
  

• The overall change in internal energy also considers both heat and work contributions.

10.3 Enthalpy: Reactions Carried Out at Constant Volume or at Constant Pressure

• Definition of Enthalpy (H):
    - A thermodynamic function defined as:
      - $H = U + PV$
    - Where $U$ is internal energy, $P$ is pressure, and $V$ is volume.

Reactions at Constant Volume

• For constant volume, changes in enthalpy relate to internal energy with no work done since $ext{ΔV} = 0$.
   

Reactions at Constant Pressure

• For constant pressure, there can be volume change, hence pressure-volume (PV) work is considered.
    - PV Work Formula:
      - $w = -PΔV$,
    - Where $P$ is external pressure and $ΔV$ is change in volume.

Worked Example 10.1: Change in Internal Energy

• Problem Statement: Calculate $ext{ΔU}$ for a system absorbing 188 J of heat and doing 141 J of work on surroundings.

• Solution:
    - $ΔU = q + w = 188 J + (-141 J) = 47 J$

10.4 Calorimetry

• Calorimetry defined as the measurement of heat changes in a chemical process using a calorimeter.

• Heat capacity defined as the amount of heat required to raise an object's temperature by 1°C.

Specific Heat ($s$)

• Specific heat defined as:
    - $s = rac{q}{mΔT}$
    - Where:
      - $q$ = heat absorbed,
      - $m$ = mass of substance,
      - $ΔT$ = change in temperature.

• Specific Heat Values for Common Substances:
    - Water: $s = 4.184 rac{J}{g imes °C}$
    - Aluminum: $s = 0.900 rac{J}{g imes °C}$
    - Gold: $s = 0.129 rac{J}{g imes °C}$

Worked Example 10.4: Heat Required Calculation

• Problem Statement: Calculate the heat required to heat 255 g of water from 25.2°C to 90.5°C.

• Use the formula:
    - $q = smΔT$ where:
      - $m = 255 g$,
      - $s = 4.184 rac{J}{g°C}$,
      - $ΔT = 90.5°C - 25.2°C$.

Constant-Pressure Calorimetry

• Coffee Cup Calorimeter: Used to measure heat exchange during reactions at constant pressure.

• For an exothermic reaction, the system loses heat while the surroundings gain heat.

10.5 Hess’s Law

• Hess’s Law Statement: The enthalpy change for a stepwise process is the sum of enthalpy changes for individual steps.

• Application Rules:
    1. Specify the physical states of reactants and products.
    2. Multiply by factor $n$, and multiply $ΔH$ by the same factor.
    3. Reverse an equation: Change sign of $ΔH$.
  

Worked Example 10.7: Determining Enthalpy Change

• Arrange equations to sum to a desired equation, manipulating as necessary to find total enthalpy change.

10.6 Standard Enthalpies of Formation

• Defined as the heat change when one mole of a compound forms from its constituent elements in standard states (1 atm, 1 M).
    - Enthalpy of formation for elements in their standard state is zero.

• Standard Enthalpy of Reaction Formula:
    - $ΔH_{rxn} = ext{products} - ext{reactants}$

10.7 Bond Enthalpy and Stability of Covalent Molecules

• Bond Enthalpy: Enthalpy change associated with breaking a mole of bonds in gaseous molecules.

• Bond enthalpy can differ for endothermic and exothermic reactions, with tables available for reference.

Worked Example 10.10: Enthalpy of Reaction for Combustion of Methane

• Use bond enthalpies to estimate the enthalpy change for methane combustion.

10.8 Lattice Energy and Stability of Ionic Compounds

• Born-Haber Cycle: Relates lattice energy to measurable quantities, assisting in understanding ionic compound stability.

• Worked Example 10.11: Calculate lattice energy of cesium chloride (CsCl) using thermodynamic data.

Comparison of Ionic and Covalent Compounds

• Physical properties differ based on bond nature.

• Comparison Table: NaCl vs CCl4 listed with properties such as melting point, boiling point, density, etc.

Chapter Summary: Key Points

• Understanding of systems and surroundings, heat concepts, units of energy, thermochemistry fundamentals, concepts of exothermic and endothermic reactions, thermodynamics basics, states and state functions, laws of thermodynamics, heat and work, the concept of enthalpy, thermochemical equations, calorimetry types, and Hess’s Law.