10: Thermochemistry

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10.1: On Fire, But Not Consumed
10.2: The Nature of Energy - Key Definitions
All chemical changes are accompanied by the absorption or release of heat. In this unit we will review some of the fundamental concepts of energy and heat and the relation between them. We will begin the study of thermodynamics, which treats the energetic aspects of change in general, and we will finally apply this specifically to chemical change. Our purpose will be to provide you with the tools to predict the energy changes associated with chemical processes.
10.3: The First Law of Thermodynamics - There Is No Free Lunch
The first law of thermodynamics states that the energy of the universe is constant. The change in the internal energy of a system is the sum of the heat transferred and the work done. At constant pressure, heat flow (q) and internal energy (U) are related to the system’s enthalpy (H). The heat flow is equal to the change in the internal energy.
10.4: Quantifying Heat and Work
Heat is the amount of energy that is transferred from one system to its surroundings because of a temperature difference. All forms of energy can be interconverted. Three things can change the energy of an object: the transfer of heat, work performed on or by an object, or some combination of heat and work.
10.5: Measuring ΔE for Chemical Reactions- Constant-Volume Calorimetry
A bomb calorimeter operates at constant volume and is particularly useful for measuring energies of combustion.
10.6: Enthalpy- The Heat Evolved in a Chemical Reaction at Constant Pressure
Enthalpy is a state function used to measure the heat transferred from a system to its surroundings or vice versa at constant pressure. Only the change in enthalpy (ΔH) can be measured. A negative ΔH means that heat flows from a system to its surroundings; a positive ΔH means that heat flows into a system from its surroundings. Calorimetry measures enthalpy changes during chemical processes, where the magnitude of the temperature change depends on the amount of heat released or absorbed and on t
10.7: Measuring ΔH for Chemical Reactions- Constant-Pressure Calorimetry
a constant-pressure calorimeter, which gives ΔH values directly
10.8: Relationships Involving ΔHrxn
Hess's law argues that for a chemical reaction, the enthalpy of reaction (ΔHrxn) is the difference in enthalpy between products and reactants; the units of ΔHrxn are kilojoules per mole. Reversing a chemical reaction reverses the sign of ΔHrxn. The magnitude of ΔHrxn also depends on the physical state of the reactants and the products because processes such as melting solids or vaporizing liquids are also accompanied by enthalpy changes: the enthalpy of fusion (ΔHfus) and the enthalpy of vaporiz
10.9: Determining Enthalpies of Reaction from Bond Energies
10.10: Determining Enthalpies of Reaction from Standard Enthalpies of Formation
The standard state for measuring and reporting enthalpies of formation or reaction is 25 oC and 1 atm. The elemental form of each atom is that with the lowest enthalpy in the standard state. The standard state heat of formation for the elemental form of each atom is zero. The enthalpy of formation (ΔHf) is the enthalpy change that accompanies the formation of a compound from its elements. Standard enthalpies of formation (ΔHof) are determined under standard conditions: a pressure of 1 atm for ga
10.11: Lattice Energy
The Lattice energy, U, is the amount of energy requried to separate a mole of the solid (s) into a gas (g) of its ions.