General Chemistry: Thermochemistry (Chapter 7) Study Notes

Thermochemistry

Introduction to Thermochemistry

Thermochemistry is the study of energy changes, particularly heat, that occur during chemical reactions and changes of state. It is a branch of thermodynamics focused on the transfer of energy as heat and work in chemical processes.

• Energy is the capacity to do work or produce heat.

• Work is the result of a force acting through a distance.

• Energy can be exchanged between matter in contact.

Law of Conservation of Energy

The law of conservation of energy states that energy cannot be created or destroyed, only transformed from one form to another. This law governs all natural phenomena and has no known exceptions.

• Conservation of energy: The total energy of the universe remains constant.

• Energy changes in a system must be balanced by changes in the surroundings.

• When energy flows out of a system, it enters the surroundings, and vice versa.

Types of Systems

In thermochemistry, the system is the part of the universe being studied, while the surroundings are everything else.

• Open system: Both energy and matter can be exchanged with the surroundings.

• Closed system: Only energy can be exchanged; matter cannot.

• Isolated system: Neither energy nor matter is exchanged.

Forms of Energy

Energy exists in various forms, including kinetic and potential energy.

• Kinetic energy (KE): Energy due to motion. $KE = \frac{1}{2}mv^2$

• Potential energy (PE): Energy due to position or composition.

• Internal energy (E): The sum of all kinetic and potential energies of the particles in a system. $E = KE + PE$

State Functions

A state function is a property whose value depends only on the state of the system, not on the path taken to reach that state.

• Examples: Pressure (P), Volume (V), Temperature (T), Internal Energy (E), Enthalpy (H)

• Non-state functions: Heat (q) and work (w) depend on the process, not just the state.

Energy Transfer: Heat and Work

Energy can be transferred between a system and its surroundings as heat or work.

• Heat (q): Thermal energy transferred due to temperature difference.

• Work (w): Energy transferred when an object is moved by a force.

• Relationship: $\Delta E = q + w$

• Sign conventions: Heat absorbed by the system ($q > 0$), heat released ($q < 0$); work done on the system ($w > 0$), work done by the system ($w < 0$).

Heat Capacity and Specific Heat

Heat capacity is the amount of heat required to change the temperature of a substance by 1°C.

• Heat capacity (C): $q = C \Delta T$

• Specific heat capacity (c): Heat required to raise 1 gram of a substance by 1°C. $q = m c \Delta T$

• Molar heat capacity: Heat required to raise 1 mole of a substance by 1°C.

Thermal Equilibrium and Energy Transfer

When two objects at different temperatures come into contact, heat flows from the hotter to the cooler object until thermal equilibrium is reached.

• Heat lost by hot object = Heat gained by cold object

• Equation: $m_{hot} c_{hot} (T_{final} - T_{initial,hot}) = - m_{cold} c_{cold} (T_{final} - T_{initial,cold})$

Work Done by Expansion/Compression

Work can be done by a system when it expands or contracts against an external pressure.

• Equation: $w = -P_{ext} \Delta V$

• 1 L·atm = 101.3 J

Calorimetry

Calorimetry is the measurement of heat flow in a chemical reaction or physical process.

• Bomb calorimeter: Measures heat at constant volume.

• Coffee-cup calorimeter: Measures heat at constant pressure.

• Heat absorbed by calorimeter: $q_{cal} = C_{cal} \Delta T$

Enthalpy (H)

Enthalpy is a state function that represents the heat content of a system at constant pressure.

• Definition: $H = E + PV$

• Change in enthalpy: $\Delta H = \Delta E + P \Delta V$

• At constant pressure: $\Delta H = q_p$ (heat at constant pressure)

• Exothermic reaction: Releases heat ($\Delta H < 0$)

• Endothermic reaction: Absorbs heat ($\Delta H > 0$)

Thermochemical Equations and Stoichiometry

Thermochemical equations show the enthalpy change associated with chemical reactions.

• Example: $C(s) + O_2(g) \rightarrow CO_2(g) \quad \Delta H_{rxn} = -393.5 \text{ kJ}$

• Enthalpy change is proportional to the amount of substance reacted.

• If a reaction is reversed, the sign of $\Delta H$ is reversed.

• If a reaction is multiplied by a factor, $\Delta H$ is multiplied by the same factor.

• For reactions in steps: $\Delta H_{overall} = \sum \Delta H_{steps}$

Standard States and Standard Enthalpy Changes

Standard state refers to the most stable form of a substance at 1 atm pressure and a specified temperature (usually 25°C).

• Standard enthalpy change ($\Delta H^\circ$): Enthalpy change when reactants and products are in their standard states.

• Standard enthalpy of formation ($\Delta H_f^\circ$): Enthalpy change for the formation of 1 mole of a compound from its elements in their standard states.

• $\Delta H_f^\circ$ for a pure element in its standard state is 0 kJ/mol.

• Calculation: $\Delta H^\circ_{rxn} = \sum n \Delta H_f^\circ (\text{products}) - \sum n \Delta H_f^\circ (\text{reactants})$

Energy Consumption and Fuels

Chemical reactions involving fuels are important sources of energy for society. The combustion of fossil fuels (coal, petroleum, natural gas) releases energy, but also has environmental impacts.

• Average daily energy use per person is very high (e.g., 867,169 kJ/day).

• Combustion reactions are typically exothermic, releasing energy as heat.

• Environmental concerns include air pollution, acid rain, and global warming due to CO2 emissions.

Table: Energy Consumption by Source (Described)

The table compares energy consumption from different sources over time, including petroleum, natural gas, coal, nuclear electric power, and renewable energy. Petroleum and natural gas are the largest sources, with renewable energy increasing slowly.

Source	Trend (1950-2010)
Petroleum	Rapid increase, peak around 2005, slight decline
Natural Gas	Steady increase, plateau after 1970s
Coal	Steady, slight increase
Nuclear Electric Power	Begins in 1970s, slow increase
Renewable Energy	Low, gradual increase

Table: Example Combustion Reactions and Enthalpy Changes

Fuel	Combustion Reaction	$\Delta H^\circ_{rxn}$ (kJ)
Carbon	$C(s) + O_2(g) \rightarrow CO_2(g)$	-393.5
Methane	$CH_4(g) + 2 O_2(g) \rightarrow CO_2(g) + 2 H_2O(g)$	-802.3
Octane	$C_8H_{18}(l) + 25/2 O_2(g) \rightarrow 8 CO_2(g) + 9 H_2O(g)$	-5074.1

Fossil Fuels and Sustainability

Fossil fuels such as coal, petroleum, and natural gas are formed from ancient organic matter over millions of years. Their extraction and use have significant environmental impacts.

• Coal is formed from plant material buried and compressed over geological time.

• Petroleum and natural gas originate from ancient marine organisms.

• Combustion of fossil fuels releases pollutants and greenhouse gases.

Environmental Impact of Fuels

Impurities in fossil fuels and incomplete combustion can release harmful materials into the atmosphere, contributing to air pollution, acid rain, and global warming.

• Major pollutant: Carbon dioxide (CO2), a greenhouse gas.

• Other pollutants: Sulfur oxides, nitrogen oxides, particulates.

Summary

• Thermochemistry connects chemical reactions to energy changes, especially heat and work.

• Understanding energy transfer, enthalpy, and calorimetry is essential for analyzing chemical processes and their environmental impact.