Ch06_Thermochemistry

Chapter 6: Thermochemistry

Energy Definitions

• Energy: The capacity to do work; it exists in various forms and can be converted from one form to another.

• Radiant Energy: This energy originates from the sun and is the primary energy source for Earth, essential for processes such as photosynthesis in plants.

• Thermal Energy: Linked to the random motion of atoms and molecules, it plays a crucial role in determining the temperature of a substance.

• Chemical Energy: Stored within the bonds of chemical substances; this energy is released or absorbed during chemical reactions, making it a key factor in thermochemistry.

• Nuclear Energy: Found in the collection of neutrons and protons in an atom, nuclear energy is released during nuclear reactions such as fission or fusion.

• Potential Energy: This type of energy is available based on an object's position within a gravitational field or its configuration within a system.

Heat and Temperature

• Heat: Defined as the transfer of thermal energy between two bodies at different temperatures, heat flow is critical for understanding thermal processes in both chemical reactions and physical changes.

• Temperature: A measure of the average thermal energy of the particles in a substance, temperature is a key determinant of the direction of heat transfer.

• Thermochemistry: The branch of chemistry that deals with the study of heat changes that occur during chemical reactions and phase changes.

Thermodynamics Basics

• System: Refers to the specific area of study within the universe, allowing us to focus on particular interactions and energy changes.

• Types of Systems:

  • Open System: Allows for the exchange of both mass and energy with its surroundings.

  • Closed System: Energy can be exchanged, but mass remains constant within the system.

  • Isolated System: Neither mass nor energy can be exchanged with the surroundings, idealized to simplify calculations.

• Surroundings: Consists of everything outside the defined system that can interact with it in terms of heat and work.

Energy Changes in Reactions

• Exothermic Process: A reaction that releases heat to the surroundings, effectively transferring thermal energy from the system. Example Reaction: 2H₂ (g) + O₂ (g) → 2H₂O (l) + energy. This process results in an increase in temperature of the surroundings.

• Endothermic Process: A reaction that absorbs heat from the surroundings into the system. Example Reaction: H₂O (g) → H₂O (l) + energy. This process typically results in a decrease in the temperature of the surroundings.

Thermodynamics Principles

• First Law of Thermodynamics: States that energy can change forms but is conserved; it cannot be created or destroyed. Mathematically represented as DE_system + DE_surroundings = 0, indicating the total energy change in a system is equal and opposite to the energy change in its surroundings.

• Change in Internal Energy (DE): DE is calculated from the contributions of heat exchange (q) and work (w) done on or by the system. Formula: DE = q + w.

Work and Energy Calculations

• Work (w): Calculated from pressure (P) and volume change (DV). Common formulas include w = -PDV, showing work done on/by the system depending on whether the volume increases or decreases.

• Example Calculation of Work Done:

  • (a) Expansion against vacuum:

    • DV = 3.8 L, P = 0 atm → w = 0 joules

  • (b) Expansion against constant pressure:

    • DV = 3.8 L, P = 3.7 atm

    • w = -14.1 L·atm; converting to joules:

      • w = -14.1 L·atm x 101.3 J/L·atm = -1430 J

Enthalpy Changes

• Enthalpy (H): A measure of heat flow at constant pressure, with significant implications in the study of thermochemistry, as it helps quantify the energy changes during a chemical reaction.

• Relation: The change in enthalpy (ΔH) can be calculated from the equation ΔH = ΔE + PΔV, where P represents pressure, and ΔV represents the change in volume.

Thermochemical Equations

• Endothermic Reaction Example: H₂O(s) → H₂O(l), ΔH = +6.01 kJ/mol, indicating the energy needed to convert ice to water.

• Exothermic Reaction Example: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), ΔH = -890.4 kJ/mol, showcasing energy released during combustion.

Properties of Thermochemical Equations

• Stoichiometric coefficients represent moles of reactants and products and must be considered when calculating enthalpy changes.

• Reversing a reaction changes the sign of ΔH, reflecting the opposite nature of energy changes when the reaction is run in reverse.

• Multiplying the entire equation by a factor (n) results in ΔH changing by the same factor, facilitating scalability in calculations.

Specific Heat and Heat Capacity

• Specific Heat (s): Defines the amount of heat required to raise the temperature of 1 g of a substance by 1°C, crucial for understanding thermal properties of materials.

• Heat Capacity (C): The total heat needed to raise the temperature of a given quantity (m) of material by 1°C, calculated as C = m x s.

• Heat absorbed or released can be calculated using formulas: q = m x s x Δt (for the change in temperature) or q = C x Δt.

• For temperature changes, Δt is defined as tfinal - tinitial, helping to ascertain changes in thermal energy of the system.

Calorimetry

• Constant-Volume Calorimetry: A method that assumes no heat is exchanged with the surroundings, allowing for precise measurements of reaction heat.

• Constant-Pressure Calorimetry: Verifies that the heat absorbed by the solution equals the heat change occurring in the reaction, crucial for practical applications of thermodynamics in reactions.

Examples: Heat Calculations

• Detailed calculations of heat change through specific example scenarios in calorimetry are essential for demonstrating principles of thermodynamics, heat capacities, and enthalpy changes effectively.

• Molar heat of combustion calculations are performed based on specific examples to illustrate practical applications of thermochemistry.

Hess's Law

• States that the total enthalpy change for a reaction is the sum of enthalpy changes for individual steps, allowing for calculations to be simplified regardless of the reaction pathway taken.

Summary of Enthalpies of Formation

• Tables depicting standard enthalpies of formation for various compounds and elements at specified temperatures allow for reference in calculating ΔH for reactions.

• Enthalpy changes are measured under standard conditions, often at 1 atm pressure and 25°C, providing context for comparing different substances and reactions.