Thermochemistry Chapter 06 Quiz

Chapter 06: Thermochemistry

• Energy: The capacity to do work. Kinetic energy is due to motion, potential energy is due to position or composition, thermal energy is associated with temperature, and chemical energy is associated with positions of electrons and nuclei.
• Forms of Energy:
  • Electrical: Kinetic energy from the flow of electrical charge.
  • Heat/Thermal: Kinetic energy from molecular motion.
  • Light/Radiant: Kinetic energy from energy transitions in atoms.
  • Nuclear: Potential energy in the nucleus of atoms.
  • Chemical: Potential energy in the attachment or position of atoms.
• Units of Energy:
  • Joule (J): The amount of energy needed to move a 1 kg mass one meter.
  • Calorie (cal): The amount of energy needed to raise the temperature of 1 gram of water by 1°C.
  • Kilocalorie (kcal): 1000 calories.
  • Kilowatt-hour (kWh): 3.6 x 10⁶ J.

Energy Use

• Units and conversions for energy required to raise water temperature, light a bulb, run a mile, and average US citizen use in a day.
• System and Surroundings:
  • System: The material or process being studied.
  • Surroundings: Everything else that can exchange energy with the system.

Conservation of Energy

• The amount of energy gained/lost by the system must equal the amount gained/lost by the surroundings.
• Energy cannot be created or destroyed, only transferred or converted.
• Total energy present at the beginning must be present at the end.

Three Types of Systems

• Open System: Mass and energy can transfer across the boundaries.
• Closed System: Only energy can transfer across the boundaries.
• Isolated System: No transfer of mass or energy across the boundaries.

The First Law of Thermodynamics

• Thermodynamics: The study of energy and interconversions.
• First Law: Energy cannot be created or destroyed, only transferred between the system and surroundings.

Internal Energy (E)

• Sum of kinetic and potential energies of all particles in a system.
• A state function: Change in internal energy depends only on the initial and final states, not the path taken.

Energy Diagrams

• Graphical depiction of energy flow direction.
• A positive change in internal energy means the final state has more energy than the initial, and vice versa.

Energy Flow

• Energy flows out of a system, ΔE_system is negative.
• Energy flows into a system, ΔE_system is positive.

Heat Exchange

• Heat is the transfer of thermal energy between a system and its surroundings.
• Heat flows from a warmer object to a cooler object until thermal equilibrium is reached.

Quantity of Heat Absorbed

• Increased temperature is directly proportional to the amount of absorbed heat.
• Heat capacity: The proportionality constant.
• Material-dependent: Depends on the mass and type of material.

Specific & Molar Heat Capacity

• Specific heat capacity: The amount of heat energy required to raise the temperature of 1 gram of a substance 1°C.
• Molar heat capacity: The amount of heat energy required to raise the temperature of one mole of a substance 1°C.

Pressure-Volume Work

• If the external pressure is constant, the work done is calculated by the product of the external pressure and the volume change.

Measuring ΔE

• Reaction enthalpy measured by measuring heat (q) and work (w).

Calorimeter use

• Example is bomb calorimeter: used to measure heat under constant volume reaction conditions.

Enthalpy (H)

• Enthalpy is the sum of internal energy and the product of pressure and volume.
• Enthalpy change (ΔH) measures the heat involved in a constant-pressure reaction.
• ΔH and ΔE are similar in value, but the difference is significant for reactions involving gases.

Endothermic and Exothermic Reactions

• Exothermic: releasing heat, ΔH is negative
• Endothermic: absorbing heat, ΔH is positive.

Hess's Law

• Enthalpy change of a reaction is independent of the pathway taken.
• Total enthalpy change equals the sum of enthalpy changes for individual steps in the reaction pathway.

Standard Conditions

• Defined set of conditions for measuring enthalpy changes of reactions:
  • Pure gas at exactly 1 atm pressure.
  • Pure solid/liquid most stable form at exactly 1 atm pressure and a specified temperature.
  • Substance in solution (usually 25°C or 298 K).

Standard Enthalpy of Formation (ΔH_f°)

• Standard enthalpy change for the formation of 1 mole of a compound from its constituent elements in their standard states.
• Standard enthalpy of formation for pure elements in their stable states is zero.
• Calculation of standard enthalpy change for any reaction from the standard enthalpies of formation of the reactants and products.

Test your knowledge on the concepts of thermochemistry as discussed in Chapter 06. Explore different forms of energy, their units, and how they are measured. This quiz will help reinforce your understanding of how energy works in various applications.