Thermochemistry: Energy Changes in Chemical Reactions

Chapter 10: Thermochemistry: Energy Changes in Chemical Reactions

10.1 Energy as a Reactant or Product

• Combustion Reactions of Fossil Fuels: Produce the majority of energy consumed by humanity:

  • Combustion of coal

  • Combustion of natural gas (CH₄)

  • Combustion of gasoline

• Origins of Fossil Fuel Energy: Energy originates in food chains that provide nutrition to plants and animals, tracing back to photosynthesis.

10.2 Forms of Energy

• Internal Energy (E): Total of all kinetic and potential energy within an object.

• Chemical Energy: A type of potential energy that relates to the atomic bonding in a substance.

• Energy Change (ΔE): Occurs when the internal energy of reactants deviates from that of products.

• Thermochemistry: The study of energy changes accompanying chemical reactions.

10.3 Thermodynamics

• First Law of Thermodynamics: Energy cannot be created or destroyed but can only change forms. The universe's total energy remains constant.

• Components: The universe consists of the system (focus of study) and surroundings (everything outside the system).

• Energy Equation:
  $\Delta E<em>{univ} = \Delta E</em>{sys} + \Delta E<em>{surr} = 0$ $\Delta E</em>{sys} = -\Delta E_{surr}$

10.4 State Functions

• Example: A skier's potential energy increase is the same whether the skier takes a lift or hitches a hike.

• State Function Defined: A property determined solely by the current state of the system, not by how that state was achieved.

10.5 Heat and Work

• Heat (q) and Work (w): The two main avenues through which energy transfers occur.

• Energy Change Equation:
  $\Delta E = q + w$

• Work Definition: A force exerted over a distance.
  $w = F \times d$

• Internal Energy: A state function independent of acquisition methods.
  $\Delta E = E<em>{final} - E</em>{initial}$

Table 10.1: Heat and Work Flows

• Rule of thumbs for heat transfer and work:

  • Heat flow to system: q_sys > 0

  • Heat flow from system: q_sys < 0

  • Work done on system: w_sys > 0

  • Work done by system: w_sys < 0

10.6 Thermal Energy

• Thermal Energy: Portion of internal energy proportional to absolute temperature; includes the sum of particle kinetic energies.

10.7 Transferring Heat and Doing Work

• Importance of identifying the system; can range from galaxies to lab apparatus to tiny particles.

• Energy change equation restated:
  $\Delta E = q + w$

10.8 Types of Thermodynamic Systems

• Isolated System: No exchange of energy or matter with surroundings.

• Closed System: Energy can flow, but matter cannot.

• Open System: Both energy and matter can flow freely to and from surroundings.

10.9 Exothermic and Endothermic Processes

• Exothermic: Processes where heat flows from system to surroundings (ΔH < 0).

• Endothermic: Processes where heat flows into the system from the surroundings (ΔH > 0).

10.10 Pressure-Volume Work

• Pressure-volume work occurs when gases expand or compress against external pressure (P).

• Units of Work:
  $1 \text{ L atm} = 101.325 \text{ J}$
  $w = -P \Delta V$
  $\Delta E = q - P \Delta V$

10.11 Sample Exercise: Calculating P–V Work

• Example Problem: Calculation involving a tank of helium used for inflating balloons interrupts management of forces with given parameters.

• Resulting work is calculated using the atmospheric pressure to indicate pressures during inflation.

Key Steps in P–V Work Example

1. Inputs gathered on balloon volume and atmospheric pressure.

2. Estimated work done based on gas behavior at volumes inflated against external pressure (here: P = 1.02 bar).

Final results highlight the change in internal energy due to volume expansion against pressure, indicative of real-world applications involving gaseous substances.

10.12 Enthalpy and Enthalpy Changes

• Enthalpy (H): A function used to evaluate energy and heat flow in thermodynamic systems under constant pressure.
  $H = E + PV$

• Enthalpy Change (ΔH): Heat exchanged during reactions at constant pressure. Given as: $\Delta H = \Delta E + P \Delta V$

  • Exothermic reaction characterized by ΔH < 0

  • Endothermic reaction characterized by ΔH > 0

Phase Change Enthalpies

• Enthalpy of Fusion (ΔH_fus): Energy needed to convert 1 mole of solid to liquid at freezing point.

• Enthalpy of Vaporization (ΔH_vap): Energy needed to convert 1 mole of liquid to vapor at boiling point.

• Enthalpy values reported per mole converted.

Table 10.3: Enthalpies of Fusion and Vaporization

• Examples: Presented various compounds alongside their respective ΔHfus and ΔHvap values.

10.13 Heating Curves and Heat Capacity

• Heating Curve: Illustrates temperature and state changes (like melting/freezing) for substances like water.

10.14 Relating Temperature Change and Heat Flow

• Heat Capacity (CP): Energy required to raise an object's temperature by 1°C at constant pressure. Relates change in temperature (ΔT) to heat (q). $C</em>P = \frac{q}{\Delta T}$

• Specific Heat (s): Energy required to raise the temperature of 1 g of a substance by 1°C.
  $s = \frac{q}{m \Delta T}$

• Molar Heat Capacity (Cp): Energy required to raise the temperature of 1 mol of a substance by 1°C.
  $C</em>p = \frac{q}{n \Delta T}$

Table 10.4: Values of Specific Heat and Molar Heat Capacity

• Detailed values across elements and compounds.

10.15 Sample Exercise: Calculating Energy for Heating Processes

• Series of steps outlined for typical problems involving temperature increases through phase changes, highlighting calculation processes.

10.16 Sample Exercise: Calculating Final Temperature

• A case study involving mixing hot brewed tea and ice, focusing on energy balance to determine final temperature post-reaction.

Key Steps in the Thermal Balancing Exercise

1. Warm ice to melting point using specific heat equations.

2. Calculate energy required for phase change (melting) then warming of resulting water to final equilibrium temperature.

3. Apply the same for the heat loss from the hot tea as it cools to the final temperature.

Conclusion of Thermal Balancing Examples

• Significant effort is involved in setting up systems for changes through heat transfer, but consistency in applying formulas yields results.

10.17 Calorimetry

• Calorimetry: Determines energy quantities transferred during reactions using calorimeters.

  • Calorimeter Defined: A device measuring energy absorption or release.

• Enthalpy of Reaction (ΔH_rxn): Change in enthalpy observed during chemical processes.

Table: Standard Heat Protections

• Explore typical calorimetry scenarios alongside relationships modeled by thermochemical equations.

10.18 Sample Exercise: Calculating ΔH_rxn

• Calculations of ΔH_rxn based on calorimetry data help solidify concepts in reaction energy transfer.

10.19 Bomb Calorimetry

• Use of bomb calorimeter for combustion energy measurements.

10.20 Standard Enthalpy of Reaction

• Hess’s Law: Entropy values of reactions sum to determine unknown values indirectly using known conditions.

  • Defines approaches to solving complex chemical reaction calculations dynamically with alternative processes.

10.21 Lattice Energy and Solutions

• Definitions around the lattice energy equations guide normative evaluations of solid compounds and their behavior in solution.

Conclusion

• Thermochemistry: Deeply rooted in understanding how energy transfers during chemical reactions influence physical changes in systems, aiding in predictions for contextual applications in everyday scenarios like heating methods and fuel value assessments.