JlDkDi-Thermodynamics - part 2 - enthalpy
Page 1: Enthalpy
Overview
Introduction to Enthalpy concepts relevant for thermodynamics.
Page 2: Key Concepts
Contents Covered
Standard Conditions and Standard State.
Bond Energies to Calculate Enthalpy Changes.
Standard Enthalpies of Formation.
Enthalpies from ΔfH⦵ directly.
Calorimetry techniques.
Hess’s Law in thermodynamics.
Page 3: Reaction State Dependency
The Role of States in Reactions
Example Reaction: C(s) + O2(g) → CO2(g)
Importance of the State of Carbon: diamond, graphite, C60.
For a reaction to occur, reactants must be in their standard state and at a specified temperature.
All products should also be in their standard states during calculations.
Page 4: Definition of Standard State
Standard Enthalpy
Denoted as ΔHo or ΔH⦵.
Defined for 1 mole at given temperature:
Pure substances in their most stable form at 1 bar.
Gases at a pressure of 1 bar (105 Pa).
Solutions at 1 M concentration.
Typically measured at room temperature (298 K).
Definition of a mole, approximately 6 * 10^23 atoms (Avogadro's number).
Page 5: Calculating ∆H
Bond Enthalpies Explanation
Calculate ΔHo using average bond enthalpies.
Energy needed to break 1 mole of bonds.
Bond Breaking:
Requires energy (ΔH⦵ > 0) - Endothermic reaction.
Bond Making:
Releases energy (ΔH⦵ < 0) - Exothermic reaction.
Units are expressed in kJ/mol.
Page 6: Example Calculation of Bond Energies
Breakdown of Energy in Reactions
Example: Breaking a mole of C-H bonds in methane (CH4):
Enthalpy of dissociation for C(s) → C(atoms) = +610 kJ/mol
Break 2 moles of H-H: 2 * 436 kJ/mol = +872 kJ/mol
Making 4 moles of C-H bonds: –4 * X kJ/mol
Calculation:
ΔH⦵ = -76 = +610 + 872 - 4X
Solving collectively leads to X = 389.5 kJ/mol.
Page 7: Properties of Graphite
Energy Characteristics
In graphite, approximately 1 C=C bond per carbon.
Sublimation ΔH⦵ = +704 kJ/mol;
C=C bond ΔH⦵ = +610 kJ/mol.
Comparison with ΔH⦵sublimation pointing to larger values.
Page 8: Further Breakdown of Bond Energies
Calculation Framework
Detailed energy changes for the reaction of C(graphite) + 2H2(g) → CH4(g).
Use of enthalpy values:
C(graphite) + 2 H2 (g):
Reactions and products considering ∆H and summing energies for both reactants and products.
Σ H(making products) - Σ H(breaking reactants).
Page 9: Energy Landscapes
Visual Representation of Reactions
Description of the 'energy mountains' with a focus on the reaction coordinate and energy landscape throughout the reaction path.
Page 10: Reaction Using Bond Energies
Calculation of ΔH⦵ for CH3Cl
Reaction: C(graphite) + 1.5 H2(g) + 0.5 Cl2(g) → CH3Cl(g).
Applying known bond enthalpies for calculations:
ΔH⦵ = Σ Hmaking-products - Σ Hmaking-reactants = -99 kJ/mol.
Page 11: Summary of Bond Energies
Single and Multiple Bonds Data
Display of bond energies for various elements (e.g., H, C, N) in single and multiple bonds.
Important to reference standard bond energies in thermodynamics.
Page 12: Measuring Enthalpy of Combustion
Bomb Calorimetry
Measurement defined as the enthalpy change when a compound is burned fully in oxygen.
Calculation based on heat transfer (ΔT, mass, specific heat capacity) and the moles participating in the reaction.
Page 13: Understanding Reaction Enthalpy
Comprehensive Approach
Overall equation: REACTANTS → PRODUCTS, ∆H⦵ calculation as:
∆H⦵ = Hproducts - Hreactants.
Conceptual analogy to age calculation as it deals with differences in value.
Page 14: Exothermic vs Endothermic Reactions
Example of Water Formation
Reaction: H2(g) + 0.5 O2(g) → H2O(l) demonstrating exothermic nature (∆H⦵ = -286 kJ/mol).
Clarification on manipulating coefficients for clarity in thermodynamic equations.
Page 15: Standard Enthalpy of Formation
Core Definition
ΔH⦵ denotes standard enthalpy when 1 mole of a substance is formed from its elements under standard conditions.
Page 16: Example of Standard Enthalpy of Formation
Ethanol Formation
Reaction: 2 C(s) + ½ O2(g) + 3 H2(g) → C2H5OH(l) with ∆fH⦵ = -277 kJ/mol.
Emphasizes accuracy over bond energies by taking into account the physical state of substances.
Page 17: Calculate ΔH⦵ from Standard Enthalpies of Formation
Methodology for Calculation
Using standard enthalpies to determine ΔHo through the equation:
ΔH⦵ = Σ (∆fH⦵)products - Σ (∆fH⦵)reactants.
Page 18: Table of Standard Enthalpies of Formation
Detailed Reference Table
Listing of various substances and their formation enthalpies in kJ/mol.
Useful for quick reference during calculations.
Page 19: Allotropes
Brief Discussion
Overview of the concept of allotropes and their relevance in chemistry.
Page 20: Combustion of Benzene
Benzene Reaction Setup
Reaction: C6H6(l) + x O2(g) → y CO2(g) + z H2O(l) with focus on determining ∆H⦵.
Page 21: Calculating Enthalpy Change for Benzene
Detailed Calculation
Balanced combustion reaction for benzene:
C6H6(l) + 7.5 O2(g) → 6 CO2(g) + 3 H2O(l).
ΔH⦵ for total enthalpy change calculation.
Page 22: Understanding Hess’s Law
Fundamental Principle
Discussion on how reactions may proceed non-linearly.
Introduction to Hess's Law, emphasizing energy change independence from the path taken.
Page 23: Application of Hess's Law in Reactions
Example Calculation
Reactants and products application illustrating Hess’s law via described reactions.
Page 24: Standard Enthalpy for Acetylene
Detailed Enthalpy Calculation
Consideration of reactions and combustion heat requirements in acetylene formation.
Page 25: Summation of Reactions (Hess’s Law)
Example with Acetylene
Detailed working to find standard enthalpy using Hess’s Law based on various reactions.
Page 26: Hard Example of Hydrogenation Reaction
Reactions and Enthalpy Changes
Calculation involving hydrogenation of benzene, including combustion reactions with their enthalpy values.
Page 27: Continued Example Analysis
Simplifying and Summing Reactions
Further breakdown and use of Hess's Law for clear enthalpy calculations.
Page 28: Final Steps in Calculation
Closing the Example
Final summation and simplification of reaction enthalpy calculations for clarity.