Energy and Chemical Change
Energy and Chemical Change
Energy Basics
Definition: Energy is the ability to do work or transfer heat.
Units of Energy:
SI Unit: Joule (J)
When calculated value > 1000 J, use kilojoules (kJ).
Alternate unit: Calorie (cal), where 1 cal = 4.184 J (exact).
Nutritional unit: Calorie (capital C), equivalent to one kilocalorie: 1 Cal = 1 kcal = 4.184 kJ.
Types of Energy
Kinetic Energy (K): Energy of motion.
Equation: K = 1/2 mv²
Potential Energy (P): Stored energy in matter.
Equation: P = mgh (where m = mass, g = gravity, h = height)
Internal Energy (E): Sum of kinetic and potential energy for particles in a system.
Temperature and Energy
Temperature (T): Measures average kinetic energy of particles.
Units: Celsius (°C), Fahrenheit (°F), Kelvin (K).
Formula: KE average = 1/2 mv².
Heat: Energy transferred due to temperature differences, reaching thermal equilibrium.
System and Surroundings
System: Reaction area under study (e.g., water vapor).
Surroundings: Everything else in the universe.
Types of systems:
Open System: Can exchange mass and energy (e.g., human body).
Closed System: Can exchange energy but not mass (e.g., light bulb).
Isolated System: No exchange of mass or energy (e.g., closed Thermos).
Conservation of Energy
First Law of Thermodynamics: Energy cannot be created or destroyed, only transformed.
For isolated systems: Internal energy (E) remains constant.
Changes in energy: ΔE = E product - E reactant = 0.
E is a state function.
Energy Transfer Methods
Energy can be transferred as heat (q) and work (W).
Equation: Heat = q + W.
Internal energy changes yield similar total energy regardless of the path taken.
"Dead" battery: Internal energy that does no work but releases heat.
State Functions
State Function: Property depending only on the current state of the system, not on the path taken.
Example: Position – different paths (train vs. car) yield the same endpoint.
Heat Transfer
Heat Transfer (q): Transfer of energy from high to low temperature regions.
Units: J, cal, kgm²/s².
Heat transfer is a state function; the path taken does not affect the heat transferred.
Heat Capacity and Specific Heat
Heat Capacity (C): Ability of an object to absorb heat, varies with mass and substance.
Units: J/°C.
Specific Heat (s): Ability of a substance to store heat.
Units: J/g°C or J/mol·K.
Formula: q = m × s × ΔT (where ΔT = change in temperature).
First Law of Thermodynamics Applied
Heat transfers monitored can predict internal energy changes.
Example: Transfer of heat from hot to cold water.
Equation: q = -q_water = +q_sample.
Chemical Potential Energy
Chemical Bonds: Attractive forces binding nuclei and electrons.
Exothermic Reactions: Form stronger bonds, release energy.
Endothermic Reactions: Break stronger bonds, absorb energy.
Calorimetry
Measures the heat of reaction: changes in temperature reflect heat absorbed or released.
Calorimeters: Tools for measuring heat changes during reactions.
Types: Coffee-cup calorimeter for constant pressure, bomb calorimeter for constant volume.
Enthalpy and Heat Transfers
Enthalpy Change (ΔH): Associated with heat transfers at constant pressure.
Enthalpy change of reaction measured through temperature changes in calorimeters.
Hess's Law
Hess's Law: Total enthalpy change equals the sum of enthalpy changes for individual steps regardless of the path.
Example calculations for reactions show how many kJ are released per mole of reactants.
Standard State
Standard State: Most stable form of a pure substance at 1 atm and specified temperature, usually 25°C.
Allotropes: Different forms of the same element in the same physical state (e.g., graphite vs. diamond).
Enthalpy of Formation
Enthalpy of formation (ΔH°f) measures the change for the formation of one mole of a substance from elements in their standard states.
Note: ΔH°f = 0 for elements in their standard state.