thermodynamics

Thermodynamics Overview

• Definition: Thermodynamics is the study of the relationship between heat, work, temperature, and energy. In chemistry, every chemical reaction involves a change in energy, often in the form of heat transfer.

• Related Terminology:
      - Thermochemistry: the study of energy transfers as heat that accompany physical changes and chemical reactions.
      - Temperature Scales: Celsius (°C) and Kelvin (K) are used in thermochemical calculations.

Systems in Thermodynamics

• System: The specific portion singled out for study.

• Surroundings: Everything else outside the system.

Types of Systems

• Open System: Matters and energy are exchanged with surroundings.

• Closed System: Energy (but not matter) is exchanged with surroundings.

• Isolated System: Neither matter nor energy is exchanged with surroundings.

Examples of Systems

• Open System: A pot of boiling water with no lid. Heat enters from the burner, and steam (matter) escapes.

• Closed System: A pot of boiling water with a lid that prevents steam from escaping. Heat enters, but matter remains contained.

• Isolated System: A well-insulated thermos storing hot coffee, preventing heat and matter escape.

Analyzing Systems

• Examples to categorize:
      - The Earth: Closed system due to atmosphere and energy exchange.
      - The Human Body: Open system due to matter and energy intake/exhaust.
      - Electric hair straightener: Often treated as a closed system.

Laws of Thermodynamics

Zeroth Law of Thermodynamics

• Equilibrium: A system whose properties do not change over time without external influences.

• Statement: If two systems are in thermodynamic equilibrium with a third system, they are in equilibrium with each other.

First Law of Thermodynamics

• Law of Conservation of Matter: Matter cannot be created or destroyed, only transformed.

• First Law of Thermodynamics: States that energy must be conserved; the energy in a system can change forms, but not in quantity.

Second Law of Thermodynamics

• Entropy (S): Represents the unavailability of a system's thermal energy for work; indicative of system disorder.

• Irreversible Processes: Any process increases total entropy; heat flow occurs spontaneously from hot to cold bodies.

• Example: In a car engine, some energy as heat cannot be utilized for work, increasing system's disorder.

Third Law of Thermodynamics

• Kinetic Molecular Theory: States all matter is in perpetual motion. As temperature decreases, kinetic energy and entropy decrease.

• Absolute Zero: The state at which all molecular motion stops. Entropy approaches a constant value (zero for perfect crystalline solids).

Entropy Formula

$S = k ext{log} W$

• Where $k$ is the Boltzmann constant, and $W$ is the number of microstates of a system.

Energy Dynamics

Nature of Energy

• Energy Definition: The capacity to do work or transfer heat.

• Types of Energy:
      - Work (W): Energy used to cause an object to move against a force.
      - Heat (q): Energy causing temperature increases in an object.

• Kinetic Energy (Ek):
      - Formula: $E_k = rac{1}{2}mv^2$
      - Depends on mass ($m$) and speed ($v$).

• Potential Energy: Stored energy from attractions or repulsions between objects.

Energy Units

• SI Unit: Joule (J), defined as $1 J = 1 ext{kg} rac{m^2}{s^2}$.

• Conversions:
      - 1 calorie (cal) = 4.184 J
      - 1 kilocalorie (kcal) = 1000 cal.

Internal Energy and Changes

Internal Energy

• Internal Energy (E): Sum of all kinetic and potential energies in a system.

• Change in Internal Energy ($ riangle E$): Defined as
  $riangle E = E_{final} - E_{initial}$

• Implications of $ riangle E$:
      - Positive $ riangle E$: System gains energy from surroundings.
      - Negative $ riangle E$: System loses energy to surroundings.

Heat and Work Relation

• Change in internal energy can be expressed as:
  $riangle E = q + w$

• Meaning:
      - Heat added ($q > 0$) or work done on the system ($w > 0$) results in increased internal energy.

Example Problem

• A system loses 1375 J of heat and experiences 235 J of work done on it:
      - $riangle E = (-1375J) + (235J) = -1140J$

Heat Transfer in Reactions

Endothermic & Exothermic Processes

• Endothermic: System absorbs heat (e.g., ice melting).

• Exothermic: System releases heat (e.g., combustion of gasoline).

• Figure comparison of reaction progress illustrates differences in energy absorption/release between the two types of reactions.

Examples of Endothermic & Exothermic Reactions

• Determine whether the following are endothermic or exothermic:
      - Making an ice cube: Exothermic
      - Melting NaI at its melting point: Endothermic
      - Hydrogen gas fire: Exothermic
      - Evaporation of ammonia: Endothermic

Reaction Rates

Kinetic Molecular Theory and Reaction Rate

• Chemical Reactions: Occur when particles collide at correct orientations due to their motion.

• Impact of Temperature: Increasing temperature increases molecular speed and, consequently, collision frequency, enhancing reaction rates.

Collision Requirements

• Not all collisions yield reactions; correct orientation is required.

• Example Reaction: $ext{Cl} + ext{NOCl} <br>ightarrow ext{NO} + ext{Cl}_2$ (must collide correctly to form Cl₂).

Calculating Reaction Rate

• Reaction rates calculated as concentration per time (M/s).

• Example Reaction:
      - 3H₂ + N₂ = 2NH₃
      - Initial concentration of NH₃ (time 0) = 0.0 M, at time 10s = 0.25 M.
      - Rate determined from concentration change over time.

Calorimetry

Heat Measurement Method

• Calorimetry: Measures heat absorbed or released using a calorimeter (e.g., bomb calorimeter).

• Calorimeter function: Measures temperature changes of water to determine energy absorbed or released in chemical or physical changes.

Specific Heat

Definition and Calculation

• Specific Heat (cᵖ): Energy required to raise the temperature of 1 gram of a substance by 1 °C or 1 K, measured in j/g·K.

• Specific heat formula:
  $c^p = rac{q}{m riangle T}$

• Parameters:
      - $q$ = heat added,
      - $m$ = mass,
      - $ riangle T$ = change in temperature.

Example of Specific Heat Calculation

• Example: A 4.0g sample of glass absorbs 32 J of heat, heated from 274 K to 314 K:
      - $c^p = rac{32J}{4.0g imes 40K} = 0.20 J/gK$

Specific Heats of Common Materials

Enthalpy Changes in Reactions

Enthalpy ($ riangle H$)

• Definition: The heat absorbed or released during a chemical reaction under constant pressure.

• Positive $ riangle H$: System is endothermic (absorbing heat).

• Negative $ riangle H$: System is exothermic (releasing heat).

• Enthalpy change formula:
  $riangle H = H_{final} - H_{initial}$

Enthalpy of Reactions

Example of Enthalpy Change

• Reaction:
      $2H_2(g) + O_2(g) <br>ightarrow 2H_2O(g) ext{ with } riangle H = -483.6 kJ$

• Indicates an exothermic reaction as the heat is released.

• Thermochemical equations: Coefficients indicate moles of reactants/products.

Thermochemical Equation Guidelines

1. Enthalpy ($ riangle H$) is extensive; magnitude is proportional to reaction size.

2. Enthalpy change is equal in magnitude but opposing in sign for the reverse reaction.

3. The state of reactants/products influences the enthalpy change.

Example Problem of Heat Release

• Calculate heat released when 4.5g of methane gas burns:
  $rac{(4.5g CH_4)(1 mol CH_4/16.0g CH_4)(-890kJ/1 mol CH_4) = -250 kJ}$

Additional Example

• Decomposing $2H_2O_2(l) <br>ightarrow 2H_2O(l) + O_2(g) ext{ with } riangle H = -196 kJ$.

• Calculate heat released by decomposing 5.00g of $H_2O_2$.