CH 5 (10/28) (PG 1-6)
Chemistry and Chemical Reactivity: Principles of Chemical Reactivity: Energy and Chemical Reactions
Key Questions to Address
How do we measure and calculate the energy changes that are associated with physical changes and chemical reactions?
What is the relationship between energy changes, heat, and work?
How can we determine whether a chemical reaction is product-favored or reactant-favored at equilibrium? (Refer to Chapter 18)
How can we determine whether a chemical reaction or physical change will occur spontaneously, without outside intervention? (Refer to Chapter 18)
Types of Energy
Energy: The capacity to do work.
Kinetic Energy (KE): Energy associated with motion.
Formula: KE = (1/2)mv^2
Types of kinetic energy include:
Thermal (disorderly motion of submicroscopic particles)
Mechanical (energy of moving macroscopic objects)
Electrical (movement of electrons through a conductor)
Potential Energy: Energy associated with position.
Types of potential energy include:
Gravitational potential energy
Electrostatic potential energy
Chemical potential energy associated with molecular structure and bonds.
Law of Conservation of Energy
Energy can neither be created nor destroyed.
Indicates that the total energy of the universe remains constant.
Mathematically expressed as one of the statements of the first law of thermodynamics:
△U = 0 for an isolated system
△U = q + w
Where:
U = internal energy (the sum of kinetic energy and potential energy)
△U = change in internal energy
q = heat (energy exchanged as heat)
w = work (energy exchanged as work; work is the movement against an opposing force)
Energy at the Molecular Level
Thermal Energy: Related to the kinetic energy of particles (atoms or molecules).
Transfer of thermal energy occurs as heat between a hotter object and a cooler one.
Chemical Energy: Potential energy associated with the chemical composition, specifically the bonds between atoms and possible reactions.
Heat
Heat: Transfer of thermal energy.
Occurs from a hotter object to a cooler object.
Energy flows from an object with higher average kinetic energy to one with lower average kinetic energy.
Symbol for heat (energy transferred as heat):
System and Surroundings
System: The part of the universe selected for study (e.g., contents of a beaker).
Types of systems:
Open System: Can exchange matter and energy with surroundings.
Closed System: Can exchange energy but not matter.
Isolated System: Exchanges neither matter nor energy.
Surroundings: The rest of the universe outside the system that exchanges energy and/or matter with the system.
Energy Changes in a System
General relationship for energy changes: △U = q + w
△U = U{final} - U{initial}
Endothermic Process: The system absorbs heat, with heat flowing from surroundings into the system.
Exothermic Process: The system releases heat, with heat flowing from the system to the surroundings.
Energy Units
Joule (J): The official SI unit of energy.
Kilojoule (kJ): Commonly used; 1 kJ = 1000 J.
Calories:
1 Calorie (Cal) = 1000 calories (cal)
1 cal = 4.184 J (exact)
Heat Capacity
Heat Capacity: The quantity of heat required to raise the temperature of an object by one Kelvin (or one degree Celsius).
Formula: q = C{object} x △T
Where:
= energy gained or lost as heat
= heat capacity of the object (J/K)
△T = change in temperature
△T = T{final} - T{initial}
Extensive and Intensive Properties
Extensive Property: Depends on the quantity of material considered (e.g., mass, volume, heat capacity).
Intensive Property: Independent of the quantity of material (e.g., density, specific heat capacity).
Specific Heat Capacity
Definition: The quantity of heat required to raise the temperature of 1 gram of a substance by one Kelvin (or one degree Celsius).
Depends on:
Composition of the substance
Physical state (solid, liquid, gas) of the substance.
Specific heat capacity formula: q = C x m x △T
Where:
= energy gained or lost as heat
= mass of a substance
= specific heat capacity (J K⁻¹ g⁻¹)
△T = change in temperature
△T = T{final} - T{initial}
Quantitative Aspects of Heat Transfer
Energy transfer as heat occurs spontaneously:
From an object at a higher temperature to an object at a lower temperature.
Continues until both objects reach the same temperature (thermal equilibrium).
Thermal Equilibrium: Point at which the gained and lost thermal energy is balanced.
Energy Changes from Chemical Reactions
Energy Changes: △T = q + w
Discusses the transition of energy from system internal changes.
Endothermic versus exothermic reactions define whether energy is gained or lost by the system.
Energy and Changes of State
Melting: Energy transferred as heat to convert a solid to a liquid is called heat of fusion (denoted as △{fus}H).
Vaporization: Energy transferred to convert a liquid to a gas at its boiling point is called heat of vaporization (denoted as △{vap}H).
Contrast between a change of state and a mere temperature increase through energy addition needs careful consideration of q values across distinct phases.
For example, determining the heat absorbed from a temperature range of -50 °C to 200 °C for a quantity of water requires accounting for heat capacities at various phases (solid, liquid, gas) along with the latent heat transfer in phase changes.