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5: Thermochemistry

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    516431

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    • 5.0: Introduction
      When Calories Become Clues: What Thermochemistry Teaches Us About Healing and Nutrition
    • 5.1: Energy and Its Conversion
      Energy is the capacity to do work or transfer heat. Energy exists in different forms, including thermal, chemical, electrical, nuclear, and radiant energy. Energy is classified as either kinetic energy, the energy of motion, or potential energy, stored energy due to position or arrangement. Energy is conserved: it cannot be created or destroyed, only transformed from one form to another. The SI unit of energy is the joule (J).
    • 5.2: The First Law of Thermodynamics
      Internal energy (E) is a state function representing the total energy of all particles in a system. Internal energy is transferred as heat (q) and work (w).  Heat absorbed by the system is positive, while heat released by the system is negative.  Work done on the system is positive, while work done by the system is negative.  Pressure-volume work occurs when gases expand or compress.  Energy is always conserved; it cannot be created or destroyed, only transformed from one form to another.
    • 5.3: Enthalpy
      Most reactions occur at constant pressure, where the change in enthalpy (ΔH) is the heat absorbed or released. Enthalpy is a state function that accounts for both internal energy and pressure–volume work. When there is little or no volume change, ΔH is approximately equal to ΔE. Reactions with negative ΔH are exothermic; those with positive ΔH are endothermic. A thermochemical equation includes a balanced chemical equation and its corresponding ΔH value, reported per mole of reaction as written.
    • 5.4: Calorimetry
      Calorimetry measures heat changes in chemical reactions by observing temperature changes in a calorimeter. The amount of heat transferred depends on the heat capacity of the substances involved and the temperature change. According to the law of conservation of energy, the heat released or absorbed by the system must be gained or lost by the surroundings. Heat measurements can be made using a constant-pressure calorimeter (to obtain ΔH) or a bomb calorimeter at constant volume (to obtain ΔE).
    • 5.5: Hess's Law
      Hess’s Law states that the enthalpy change of an overall reaction is equal to the sum of the enthalpy changes of its individual steps. This principle follows from the fact that enthalpy is a state function—it depends only on the initial and final states, not the path taken.  Because enthalpy is an extensive property, changing the number of moles involved in a reaction will change the magnitude of ΔH. Reversing a chemical equation changes the sign of ΔH, but not its magnitude.
    • 5.6: Enthalpies of Formation
      Standard enthalpy of formation is the heat change when one mole of a compound forms from its elements in their most stable forms under standard conditions (1 atm, 25°C, and 1 M concentration for solutions). The standard enthalpy of reaction can be calculated using these formation values and Hess’s Law, which relies on the fact that enthalpy is a state function. This approach allows us to determine enthalpy changes for reactions without needing to measure them directly.
    • 5.7: Bond Energies
      Bond energy is the energy needed to break a specific bond in a gaseous molecule. Because bond energy values are often averaged across many compounds, they provide useful estimates but not exact values for individual molecules. Bond energies can be used to estimate the enthalpy change of a reaction by subtracting the energy released from forming product bonds from the energy used to break reactant bonds.
    • 5.E: Thermochemistry (Exercises)
      Thermochemistry problems are more than numerical exercises. They strengthen your ability to connect the concepts of energy, bonding, and chemical change. With regular practice, these ideas become second nature, and you develop the skills needed to predict and explain chemical behaviour in real-world contexts.

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