Chapter 5 - Thermo Chemistry Pre-Inked Postable
Page 1: Introduction
Course: General Chemistry (CHEM 111)
Chapter: Thermochemistry
Instructor: Hans Mentzen
Lecture Date: November 15, 2021
Page 2: Types of Energy
Potential Energy: Energy stored in an object due to its position or arrangement.
Kinetic Energy: Energy of an object in motion.
Energy associated with various forms such as chemical, nuclear, etc.
Page 3: Reaction Energies
Kinetic Energy in Reactions: Energy associated with molecular motion in chemical reactions.
Example of an OH explosion involving H2 and O2:
H2 + O2 → OH
Page 4: Enthalpy Change in Reactions
Enthalpy Change (ΔH): Measure of energy change due to breaking or forming of bonds during a chemical reaction.
Exothermic Reaction:
Reactants have higher enthalpy than products (energy released).
Example:
Reactants → Products (ΔH < 0)
Page 5: Endothermic Reactions
Endothermic Reaction:
Reactants have lower enthalpy than products (energy absorbed).
Example:
Products → Reactants (ΔH > 0)
Page 6: Enthalpy Changes and Stoichiometry
Example Reactions:
N2(g) + 3 H2(g) → 2 NH3(g)
½ N2(g) + ¾ H2(g) → NH3(g)
2 NH3(g) → N2(g) + 3 H2(g)
Examining how stoichiometry affects ΔH depends on the reaction coefficients.
Page 7: Hess’s Law
Hess’s Law: If chemical equations are added, their enthalpy changes can also be added.
Enthalpies of reactions are state functions, meaning the value depends only on initial and final states, not on the path taken.
Page 8: Example of Hess’s Law
Reaction: CH4(g) + ½ O2(g) → CH3OH(g)
ΔHrxn = ?
Related Reactions:a. CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) ΔHrxn = −802 kJb. CH3OH(g) + 3/2 O2(g)→ CO2(g) + 2 H2O(g) ΔHrxn = −676 kJ
Page 9: Another Hess’s Law Example
Reaction: C(s) + 2 H2(g) → CH4(g) ΔHrxn = ?
Related Reactions: a. H2(g) + ½ O2(g) → H2O(l) ΔHrxn = −285.8 kJb. C(s) + O2(g) → CO2(g) ΔHrxn = −393.5 kJc. CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) ΔHrxn = −890.3 kJ
Page 10: Standard Enthalpy of Formation
Definition: Enthalpy change for formation of 1 mole of a compound from its elements under standard conditions (1 bar, 25 °C).
Example:
Na(s) + ½ Cl2(g) → NaCl(s) ΔHf° = −411.12 kJ/mol
Page 11: Standard State of Elements
The ΔHf° value for an element in its standard state is equal to 0 kJ/mol.
Page 12: Example of Enthalpy Formation Equation
Write the equation for the standard molar enthalpy of formation of SO3(g).
Page 13: Using ΔHf° in Hess’s Law
ΔHf° values can be combined with Hess’s law to calculate the standard enthalpy of reaction (ΔHrxn°).
Page 14: Combustion of Dimethyl Ether
Reaction: CH3OCH3(g) + 3 O2(g) → 2 CO2(g) + 3 H2O(l) ΔHrxn = ?
Standard enthalpy values:
Dimethyl ether: -184.1 kJ/mol
CO2: -393.5 kJ/mol
H2O: -285.8 kJ/mol
Page 15: Defining the System
System: The part of the universe being studied.
Surroundings: Everything external to the system.
Universe: System + Surroundings
Page 16: Types of Systems
Open System: Exchanges both matter and energy with surroundings.
Closed System: Exchanges energy but not matter.
Isolated System: Neither matter nor energy is exchanged.
Page 17: First Law of Thermodynamics
Principle: Energy cannot be created or destroyed, only transformed.
Related quantity: Internal energy (U) = Work (W) + Heat (q).
Page 18: First Law of Thermodynamics - Energy Transfer
Endothermic Reaction:
Heat transfers into the system (q > 0).
Exothermic Reaction:
Heat transfers out of the system (q < 0).
Internal energy change (ΔU):
ΔU = q + W
Page 19: Internal Energy Calculation Example
Chemical reaction generates 367 Joules of heat and does 717 Joules of work.
Calculate ΔU = q + W where q = 367 J and W = -717 J.
Page 20: Heat Transfer Overview
Heat transfer (q) is associated with temperature change.
Related parameters include:
Energy gained/lost (q)
Specific heat capacity (c_sp)
Mass (m)
Temperature change (ΔT).
Page 21: Specific Heat Capacity
Definition: Amount of energy required to raise the temperature of 1 g of a substance by 1 °C.
Page 22: Heating Water Example
Example Calculation: Energy added to a sample of water to achieve a final temperature of 36.0 °C from an initial temperature using known specific heat capacity.
Page 23: Thermal Equilibrium
Thermal Equilibrium: The state where hot objects transfer heat to colder objects until they reach the same temperature.
Page 24: Thermal Equilibrium Illustration
Repeated definition of thermal equilibrium emphasizing the heat transfer process.
Page 25: Specific Heat of Silver Experiment
Experiment: Student determines the specific heat of silver by conducting heat transfer experiments with water and measuring temperature changes.
Calculations based on mass and final temperatures lead to the determination of specific heat.
Page 26: Energy and Phase Changes
Heat of Fusion (ΔH_fus)
Heat of Vaporization (ΔH_vap)
Page 27: Phase Change Graph
Graph illustrating energy changes during phase transitions of water with respect to heat added (kJ).
Page 28: Energy Units
SI unit of energy: Joule (1 J = 1 Kg·m²/s²).
Definition of calorie: The amount of energy needed to raise 1 g of water by 1 °C (1 cal = 4.184 J).