Study Notes for Thermochemistry Chapter 5

Chapter 5 Thermochemistry

5.1 Energy and Energy Changes

Forms of Energy

• Energy is defined as the capacity to do work or transfer heat.

• Common forms of energy: heat, light, electricity.

• All energy is classified as either kinetic or potential.
    - Kinetic Energy (KE): Energy of motion, calculated by the equation:
      $KE = rac{1}{2} m u^2$
      where:
      - $m$ = mass of the object
      - $u$ = velocity of the object
    - Thermal Energy: A specific form of kinetic energy associated with the random motion of atoms and molecules, detectable through temperature changes.
    - Potential Energy (PE): Energy due to the position of an object. Types most relevant to chemistry include:
      - Chemical Energy: Energy stored in chemical bonds, dependent on types and arrangements of atoms in a molecule.
      - Electrostatic Energy: Energy resulting from the interaction of charged particles is influence by distance (d) between charges and the magnitude of charges. The equation is as follows:
      $E = k rac{q_1 q_2}{d}$
      where:
      - $q_1$, $q_2$ = magnitudes of the charges
      - $k$ = constant of proportionality.

Interconversion of Energy

• Kinetic and potential energy can be interconverted. For instance, falling water converts potential energy to kinetic energy.

• Chemical reactions can convert chemical energy into thermal energy, as seen in exothermic reactions.

Conservation of Energy

• The total energy in the universe remains constant; energy can neither be created nor destroyed. This is known as the law of conservation of energy.
  

Energy Changes in Chemical Reactions

Defining the System and Surroundings

• System: The specific part of the universe under study (e.g., a chemical reaction).

• Surroundings: Everything outside the system.

Heat Flow

• Heat: The transfer of thermal energy between two bodies at different temperatures. The concept of heat flow typically refers to heat absorbed or released during a process.

Types of Processes

• Exothermic Process: A process that releases heat to the surroundings.

• Endothermic Process: A process that absorbs heat from the surroundings.

Units of Energy

• SI Unit: The joule (J).
    - Defined as the energy of a 2-kg mass moving at 1 m/s.
    - Also defined: one joule is the energy exerted when a force of 1 newton is applied over a distance of 1 meter.

• Calorie (cal): A non-SI unit defined as the energy required to raise the temperature of 1 g of water by 1°C.
    - Conversion: 1 cal = 4.184 J
    - Nutrition calorie (Cal) = 1 kcal = 1000 cal = 4184 J.

Sample Problems

• Example 5.1(a): Kinetic energy of a helium atom moving at 125 m/s:
    $KE = rac{1}{2} m u^2$,
    where $m = 6.649 imes 10^{-27} ext{ kg}$, yields a result of $5.19 imes 10^{-23} ext{ J}$.

5.2 Introduction to Thermodynamics

Overview of Thermodynamics

• Thermochemistry is a subset of thermodynamics, which is the scientific study of the interconversion of heat and other kinds of energy.

Laws of Thermodynamics

• Important for understanding energetics and directionality of reactions. The first law of thermodynamics is central to thermochemical studies:
    - First Law of Thermodynamics: States energy can be converted but not created or destroyed.

Types of Thermodynamic Systems

1. Open System: Exchanges mass and energy with surroundings.

2. Closed System: Exchanges energy but not mass.

3. Isolated System: Does not exchange mass or energy.

States and State Functions

• State functions are macroscopic properties like energy (U), temperature (T), pressure (P), and volume (V), defined by the state of the system regardless of how the state was achieved.

• Analogy: Altitude (a state function) vs. path traveled to reach it (not a state function).

Work and Heat

• When heat is absorbed or released, or when work is done, the internal energy (U) of a system changes follow these rules:
    - $ext{Change in Internal Energy (ΔU)} = q + w$ where:
      - $q$ = heat (positive if absorbed, negative if released)
      - $w$ = work (positive if done on the system, negative if done by the system).

Heat (q) and Work (w)

• Heat: Energy transfer due to temperature difference (positive if absorbed, negative if released).

• Work: Energy transfer due to mechanical processes.

Sample Problems

• Example 5.2: Calculate the change in internal energy for a system that absorbs 188 J of heat and does 141 J of work on its surroundings:
    - $ext{ΔU} = q - w = 188 ext{ J} - 141 ext{ J} = 47 ext{ J}$.

5.3 Enthalpy

Reactions at Constant Pressure and Volume

• The change in internal energy can often be expressed in terms of enthalpy (H).
    - $ext{ΔU} = q + W$
    - Enthalpy change is given by:
    - For constant volume, $q_v = ΔU$;
    - For constant pressure, $q_p = ΔU + PΔV$.

Thermochemical Equations

• Enthalpy H is defined as:
    $H = U + PV$
    - Change in enthalpy is calculated during reactions at constant pressure.

Significance of Enthalpy Changes

• Exothermic Reactions: Enthalpy decreases (negative ΔH) as heat is released.

• Endothermic Reactions: Enthalpy increases (positive ΔH) as heat is absorbed.

Sample Problems

• Example: Calculate the change in enthalpy (ΔH) for a sodium-water reaction.

5.4 Calorimetry

Specific Heat and Heat Capacity

• Specific Heat (s): Amount of heat required to change 1g of a substance by 1°C.

• Heat Capacity (C): Amount of heat required to change the temperature of an object.

• Example: Calculate the heat required to heat a substance using the equation:
    $q = msΔT$
    where:
    - $m$ = mass,
    - $s$ = specific heat,
    - $ΔT$ = change in temperature.

Constant-Pressure Calorimetry

• Involves using a calorimeter to measure heat changes during reactions.

• At constant pressure, the heat flow out or in is associated with the enthalpy change of the system.

Sample Problems

• Example: Calculate energy content of a cookie after combustion where temperature increases.

5.5 Hess’s Law

Hess's Law Principle

• Hess's Law: The total enthalpy change during a chemical reaction is independent of the pathway taken, thus can be calculated by summing the enthalpy changes for the individual steps.

Practical Applications and Examples

• Sample problem calculating enthalpy change for reactions using Hess's law.

5.6 Standard Enthalpies of Formation

Definition and Importance

• Standard Enthalpy of Formation (ΔHf°): Heat change for the formation of one mole of a substance from its constituent elements in their standard states.

• Standard States:
    - Gases: Pure gas at 1 atm.
    - Liquids/Solids: Pure substance in the most stable form at 1 atm and specified temperature (usually 25°C).
    - Solutions: Concentration of exactly 1M.

Calculation Examples

• Sample problem to calculate standard enthalpy of formation using available enthalpy values.