Thermochemistry Notes

Thermochemistry

What is Thermochemistry?

• Thermochemistry is the study of energy changes that occur during chemical reactions and physical changes.

• It aims to differentiate between physical changes and chemical changes based on energy transformations.

What is Energy?

• Energy is the ability to do work or produce heat, measured in Joules (J).

• Different Forms of Energy:

  • Kinetic Energy: Energy of matter in motion (atoms).

  • Potential Energy: Ability to do work due to position, composition, or configuration (stored energy).

  • Chemical Energy: Potential energy stored in the chemical bonds of a substance.

• Temperature: The average kinetic energy of a substance, measured in Kelvin, °C, or °F.

Energy Forms

• Breaking Bonds: Requires an input of energy.

• Forming Bonds: Results in an output of energy.

Heat

• Temperature vs. Heat:

  • Temperature: Measures the hotness or coldness of matter.

  • Heat (Thermal Energy): The form of energy transferred between matter due to a difference in their temperatures; moves from a warmer object to a cooler object.

  • Thermal Energy: The total potential and kinetic energy of the molecules of a substance caused by their arrangement and constant motion.

Law of Energy Conservation

• In an insulated system, any heat lost by one quantity of matter must be gained by another.

• Heat transfer continues until both objects reach the same temperature.

• Example: Heat transfers from hand to ice, breaking the bonds in the ice, causing it to melt.

Energy Transfer

• Most common form of energy transfer: heat.

• Symbol for heat: Q.

Calorimeter

• A calorimeter is a device used to measure energy given off or absorbed during a chemical or physical change.

Heat Units

• Units: Joules (J) or Calories (Cal).

• Calorie: Amount of energy the body gets from food.

• $1 \text{ calorie} = 4.184 \text{ Joules}$

• $1000 \text{ calories} = 1 \text{ Calorie}$

• $1 \text{ kJ} = 1000 \text{ J}$

Heat Unit Conversion Sample Problem

• An average size baked potato has an energy value of 686,000 Joules. To convert this value to Calories, divide by 4.184 to get calories, then divide by 1000 to get Calories (food calories).

Specific Heat

• Specific Heat (Cp): Amount of heat needed to raise the temperature of 1 gram of a substance by 1°C.

• Every substance has its own specific heat.

• Symbol: $C_p$

• Unit: J/g·°C

• Specific heat of H2O: 4.184 J/g·°C

• The specific heat is why sand heats up faster than water; sand has a lower specific heat (0.33 J/g·°C).

Heat Formula

• $Q = (m)(\Delta T)(C_p)$

  • Q = heat/energy lost or gained (J)

  • m = mass (grams)

  • $\Delta T$ = change in temp (final - initial) (°C or K)

  • $C_p$ = specific heat (J/g°C or J/g°K)

Heat Sample Problem 1

• How much heat is lost when a solid aluminum ingot with a mass of 4110g cools from 660.0°C to 25°C? (The specific heat for aluminum is 0.903J/g°C)

Heat Sample Problem 2

• Suppose a piece of iron with a mass of 21.5 grams at a temperature of 100.0°C is dropped into an insulated container of water. The mass of the water is 132 grams and its temperature before adding iron is 20.0°C. What will the final temperature of the system be? (The specific heat of iron is 0.449 J/g°C)

Entropy

• Measure of the molecular disorder, or randomness, of a system.

• Dependent on a system’s state, determines the amount of possible energy.

• Related to phases: Solid vs Liquid vs Gas.

Enthalpy

• A thermodynamic quantity equivalent to the total heat content of a system.

• It is equal to the internal energy of the system plus the product of pressure and volume.

Laws of Thermodynamics

• 0th Law: If two systems are in equilibrium with a third system, the two original systems are in equilibrium with each other.

• 1st Law: Energy can be converted, but not created or destroyed.

• 2nd Law: In an insulated system, entropy can only increase.

• 3rd Law: Entropy of a perfect crystal at a temp of 0 K (absolute 0) is equal to zero.

Types of Energy Transfers

• Conduction: Transfer of energy through direct contact.

• Convection: Transfer of energy through liquids or gas.

• Radiation: Transfer of energy through waves.

Insulators vs. Conductors

• Conductors: Allow the transfer of thermal energy.

• Insulators: Prevent the transfer of thermal energy.

Phase Diagrams

• Melting: Solid to liquid.

• Freezing: Liquid to solid.

• Vaporization: Liquid to gas.

• Condensation: Gas to liquid.

• Sublimation: Solid to gas.

• Deposition: Gas to solid.

Heating Curve Points

• Identify the 3 states of matter (solid, liquid, gas).

• Flat line = phase change.

• At each phase change, energy is used for breaking or forming intermolecular forces rather than changing temperature.

Endothermic vs. Exothermic

Endothermic Reactions

• Absorb energy from the surroundings.

• Products have more energy than reactants.

• Heat is a reactant.

• Surroundings get "cold."

• Example: Cold packs.

Exothermic Reactions

• Release energy into the surroundings.

• Reactants have more energy than products.

• Heat is a product.

• Surroundings get "warm."

• Example: Burning match.

Activation Energy

• Activation Energy ($E_A$): Energy required for a reaction to occur.

  • Low $E_A$ = fast reaction rate

• $\Delta H$: Amount of energy lost or gained during a reaction.