Overview
This lecture covers the First Law of Thermodynamics, types of enthalpy changes, energy diagrams, and practical aspects of calorimetry, with calculations and problem-solving methods for enthalpy changes in chemical reactions.
First Law of Thermodynamics
- Energy cannot be created or destroyed, only transformed from one form to another.
- Energy in chemical reactions is transferred as heat between the system (reactants/molecules) and the surroundings (everything else).
Exothermic and Endothermic Reactions
- Exothermic reactions release heat from the system to the surrounding, causing a decrease in energy of the product.
- Endothermic reactions absorb heat from the surrounding into the system, increasing the energy of the product.
Energy Profile Diagrams
- The y-axis represents energy and the x-axis represents reaction progress or time.
- Exothermic: products have lower energy than reactants (ΔH negative).
- Endothermic: products have higher energy than reactants (ΔH positive).
- The difference between reactant and product energy is enthalpy change (ΔH).
- Activation energy is the energy required for reactants to reach the transition state.
Standard Enthalpy Change & Conditions
- Standard conditions: 25°C temperature, 1 atm pressure, 1 mol/dm³ solution concentration, carbon as graphite solid.
- Standard enthalpy change calculations use these conditions.
Calculating Enthalpy Changes
- Formula: ΔH = (mcΔθ) / n, where m = mass of water (kg), c = specific heat capacity, Δθ = temperature change, n = moles.
- Ignore sign in calculation; assign sign based on temperature change (negative for exothermic, positive for endothermic).
- Alternatively, use sum of enthalpy of formation of products minus reactants if all formation values are given.
Types of Enthalpy Changes
- Atomization: ΔHat—energy absorbed to form one mole of gaseous atoms (always positive).
- Formation: ΔHf—enthalpy change when one mole of compound is formed from elements in standard states.
- Neutralization: ΔHn—enthalpy change when one mole of water is formed from acid-base reaction (always negative).
- Combustion: ΔHc—enthalpy change when one mole of substance is completely burned in oxygen (always negative).
Calorimetry and Its Limitations
- Bomb calorimeter measures heat from combustion by tracking temperature change in water surrounding the reaction.
- Sources of error: heat loss to container, stirrer, wires, incomplete combustion due to limited oxygen, inability to observe reaction completion, incorrect heat capacity assumptions.
Problem Solving Strategy
- Always write balanced chemical equations for stoichiometry.
- Use mass/volume and molar mass/density to find moles when needed.
- Choose correct mass (usually mass of water) for calculations involving heat transfer.
- In practical calorimetry, use measured temperature change of water to calculate energy released or absorbed.
Key Terms & Definitions
- System — the part of the universe being studied (e.g., reactants).
- Surrounding — everything outside the system.
- Enthalpy (H) — heat content of a system at constant pressure.
- ΔH — enthalpy change during a reaction.
- Activation Energy — minimum energy required for a reaction to occur.
- Standard Enthalpy Change — enthalpy change under standard conditions (25°C, 1 atm, 1 mol/dm³).
- Bomb Calorimeter — device to measure heat of combustion.
- Atomization — formation of one mole gaseous atoms from elements.
- Formation — creation of one mole compound from elements in standard state.
- Neutralization — formation of one mole of water from acid and base.
- Combustion — complete burning of one mole of substance in oxygen.
Action Items / Next Steps
- Review and understand definitions of atomization, formation, neutralization, and combustion.
- Prepare the provided template for Hess’s law (questions 13–16) for the next class.
- Practice writing and balancing chemical equations for enthalpy change problems.
- Review quiz results and feedback in Microsoft Teams.