Overview
This lecture explains the concept of enthalpy, its role as a state function in chemistry, and introduces Hess's Law for calculating the heat of chemical reactions.
Energy and State Functions
- Energy is necessary for work, heat, and is present in all chemical bonds.
- Internal energy can be transferred as heat (q) or work (w), but the split is path dependent.
- State functions depend only on initial and final states, not the path taken (e.g., energy change).
- The total energy in a molecule is hard to measure, but changes in energy are measurable.
Introduction to Enthalpy
- Enthalpy (H) is the internal energy plus the energy needed to displace the environment at constant pressure and volume.
- We focus on the change in enthalpy (ΔH), not the absolute value.
- At constant pressure, and assuming only pressure-volume work occurs, ΔH equals the heat gained or lost by the system.
- Enthalpy is more practical than internal energy for reactions at constant pressure.
Measuring Enthalpy: Calorimetry
- Calorimetry measures temperature changes in an insulated vessel to determine heat changes (ΔH) during reactions.
- Scientists have compiled standard enthalpy changes for many compounds, making predictions possible without repeating experiments.
Hess's Law and Standard Enthalpy of Formation
- Hess’s Law states that the total enthalpy change for a reaction depends only on the initial and final states, not the path taken.
- Standard state: 25°C and one atmosphere; elements in their most stable form at standard state have enthalpy defined as zero.
- Standard enthalpy of formation (ΔHf) is the heat change when one mole of a compound forms from its elements in standard states.
- The enthalpy change for a reaction: ΔH = Σ(np·ΔHf,products) – Σ(nr·ΔHf,reactants).
- Number of moles must be factored in for each substance in the reaction.
Example: Hand Warmer Reaction
- Iron powder reacts with oxygen to form Iron(III) Oxide (Fe₂O₃), releasing heat (exothermic).
- ΔHf for Fe and O₂ are zero; ΔHf for Fe₂O₃ = -826 kJ/mol.
- For 4 moles of Fe, ΔH = -1652 kJ; this heat warms hands in hand warmers.
Key Terms & Definitions
- State function — Property depending only on initial and final states, not the path (e.g., enthalpy, energy change).
- Enthalpy (H) — Internal energy plus pressure-volume energy; measures heat at constant pressure.
- Calorimetry — Technique to measure heat changes during chemical reactions.
- Hess’s Law — The enthalpy change for a reaction is path-independent, depending only on reactants and products.
- Standard state — Defined as 25°C and 1 atm for chemicals; most stable form for elements.
- Standard enthalpy of formation (ΔHf) — Heat change when one mole of a compound forms from its elements at standard state.
Action Items / Next Steps
- Review calorimetry techniques for measuring enthalpy changes.
- Practice using Hess’s Law and standard enthalpy tables to calculate reaction heats.
- Balance chemical equations before performing enthalpy calculations.