Chemical Thermodynamics and the 1st Law of Thermodynamics

Lesson 1: CHEMICAL THERMODYNAMICS AND THE 1ST LAW OF THERMODYNAMICS

General Concepts

• Thermodynamics
  • Represents the movement of thermal energy within a system/universe.
  • Related to the forces due to heat.
  • Involves the flow of heat into or out of a system during physical or chemical transformations.
• Chemical Thermodynamics
  • Studies the interrelation of heat and work with chemical reactions or states of matter within the confines of thermodynamic laws.
  • Essentially, it refers to the study of heat and work changes during chemical and physical processes.

Fundamental Definitions

Heat

• Definition: Transfer of thermal energy between two bodies at different temperatures.
• Note: Heat is NOT equal to thermal energy itself.
• Representation: denoted by q.
• Unit: Joules (J).

Work

• Definition: Force applied to transfer energy between a system and its surroundings.
• Essential for creating heat and transferring thermal energy.
• Representation: denoted by w.
• Unit: Joules (J).

Jargons / Terms in Chemical Thermodynamics

• System:
  • Defined as the “focus” of study, any part of the universe we wish to monitor.
  • May be separated from the rest of the universe by a boundary (real or imaginary).
  • Example: A notebook.
• Surroundings:
  • Everything external to the system.
  • Example: Air surrounding the notebook.
• Universe:
  • Comprises the entirety of the system and surroundings.
  • Example: The classroom containing the notebook and air.

Overview of the Laws of Thermodynamics

1. Zeroth Law:
   • If two systems are in equilibrium with a third system, they are in thermal equilibrium with each other.
2. First Law:
   • The energy of the universe is constant.
   • Energy can be transformed but cannot be created or destroyed.
3. Second Law:
   • The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process.
4. Third Law:
   • At absolute zero, the entropy of a perfect crystal is zero.
   • The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.

Thermodynamic Properties

Examples of Exothermic and Endothermic Reactions

Exothermic Reactions

• Definition: Processes that release energy in the form of heat.
• Examples:
  • Combustion of fuels (e.g., wood, petrol, LPG, candle burning).
  • Respiration in living organisms converting glucose to energy.
  • Neutralization between acids and bases (e.g., HCl and NaOH reaction).
  • Rusting of iron in moist air.
  • Explosion of fireworks releasing light, heat, and sound.
  • Thermite reaction producing molten iron.
  • Hand warmers generating heat via chemical reactions.
  • Dissolving certain salts like CaCl2 releases heat.

Endothermic Reactions

• Definition: Processes that absorb energy in the form of heat.
• Examples:
  • Melting ice into water requires heat input.
  • Sublimation of dry ice (solid CO2) turning directly into gas.
  • Photosynthesis in plants absorbing sunlight to produce oxygen and glucose.
  • Cooking or baking processes absorbing heat.
  • Ice packs used for injury cooling absorb heat from surroundings.

First Law of Thermodynamics

• Energy Representation:
  egin{align} ext{Change in energy of a system} & : \, \Delta E{sys} = - \Delta E{sur} \, \end{align}
• Energy Change Equation:
  • The change in energy of a system equals the heat absorbed (q) plus the work done on it (w):
    $\Delta E = q + w$
  • Rearranged:
    egin{align} w & = \Delta E - q \, \text{ and }\, \ q = \Delta E - w \, \end{align}

Numerical Examples

1. Example 1
   • Given:
     • $\Delta E = 49 J$
     • $w = -35 J$
   • Solution:
     • $\Delta E = -q + w$
2. Example 2
   • Given:
     • $q = 47 J$
     • $\Delta E = 12 J$
   • Solution:
     • Compute work: $w = \Delta E - q = 12 J - 47 J = -35 J$
3. Example 3
   • Given:
     • $q = 62 J$
     • $w = 474 J$
   • Solution:
     • $\Delta E = q + w = 62 J + 474 J = 536 J$
4. Example 4
   • Given:
     • $\Delta E = -23 J$
     • $w = 45 J$
   • Solution:
     • Compute heat: $q = \Delta E - w = -23 J - 45 J = -68 J$

Lesson 2: SPONTANEOUS AND NON-SPONTANEOUS PROCESS

Important Points

1. Awareness of whether calculations pertain to the surroundings or system.
   • Example: If surroundings gain 67 J of heat, then $q = -67 J$ for the system.
2. Significant figures must be retained based on the least significant figures provided in the problem.

Definitions

• Spontaneous Process:
  • A process requiring no energy input to occur; a physical or chemical change that happens by itself.
• Non-Spontaneous Process:
  • A process requiring energy input to occur; one needing outside intervention.

Lesson 3: ENTROPY

Spontaneous & Non-Spontaneous Process Examples

Spontaneous Processes

• Examples:
  • Rusting of metals (e.g., iron in moist air).
  • Decaying of radioisotopes.
  • Drying of leaves.
  • Fireworks.
  • Oxidation of gold.
  • Dissolving salt.

Non-Spontaneous Processes

• Examples:
  • Separation of sugar from coffee grounds.
  • Charging of a battery.
  • Dissolution of sand in water.

Entropy

• Definition: A measure of disorder within a system or energy that is unavailable to perform work.
• Behavior:
  • Higher disorder of molecules indicates increased entropy; lower disorder indicates decreased entropy.
• Notation:
  • Increasing entropy is denoted by +; decreasing entropy is denoted by -.

Law of Disorder

1. Entropy of the gaseous phase is greater than that of the liquid or solid phases.
   • In summary: Solid < Liquid < Gas (in terms of increasing entropy).
2. Chemical reactions that yield more product molecules than reactants lead to increased entropy.
3. Entropy increases when substances divide; for instance, NaCl dissolving in water.
4. Entropy tends to increase with rising temperatures; molecular speed increases, resulting in greater disorder.
   • Relationship: $↑ ext{ Temperature} = ↑ ext{ Entropy} = ↑ ext{ disorder}$

Processes Accompanied by Entropy Change

• Processes:
  • Melting: Solid → Liquid
  • Vaporization: Liquid → Vapor
  • Dissolving: Solute → Solution
  • Heating: System at T₁ → System at T₂ (where T₂ > T₁)

Example Conditions Illustrating Entropy

• Examples:
  • Oxidation of nitrogen.
  • Sublimation of mothballs.
  • Reduction of silicon.
  • Lighting of candles.

Change in Entropy in the Universe

• Entropy change in the universe ($∆S<em>{univ}$) can be defined as the sum of entropy change in the system ($∆S</em>{sys}$) and the surroundings ($∆S<em>{sur}$): $∆S</em>{univ} = ∆S<em>{sys} + ∆S</em>{sur}$
  • If ∆S_{univ} > 0: Process is spontaneous.
  • If $∆S_{univ} = 0$: Equilibrium is achieved.
  • If ∆S_{univ} < 0: Reverse process occurs spontaneously.

Standard Entropy of Reaction

• Formula: $∆S° = ΣS°(products) - ΣS°(reactants)$
  • In terms of coefficients, small letters are used as coefficients, while capital letters represent products and reactants.
  • Example:
    $∆S° = [cS°(C) + dS°(D)] - [aS°(A) + bS°(B)]$

Lesson 4: GIBBS FREE ENERGY

Definitions

• Gibbs Free Energy:
  • Introduced by American physicist Josiah Gibbs, also known as Free Energy.
  • Denoted by the letter G.
• Formula: $G = H - TS$
  • Where:
  • G = free energy
  • H = enthalpy
  • T = temperature
  • S = entropy (quantities pertain to the system only).
• Change in Free Energy (for constant temperature): $∆G = ∆H - T∆S$
  • For chemical reactions:
    $∆G°_{rxn} = Σ∆H° - T∆S°$
• Spontaneity and Equilibrium Conditions:
  • ∆G < 0: Reaction is spontaneous.
  • ∆G > 0: Reaction is non-spontaneous.
  • $∆G = 0$: System is at equilibrium.

Examples in Agriculture

• Importance of the Gibbs free energy reaction between nitrogen gas and hydrogen gas to produce ammonia, essential for plant production.
• Example Calculation:
  • For reaction $N<em>2(g) + 3H</em>2(g) → 2NH_3(g)$, with given standard enthalpy and entropy values.
  • Resulting free-energy change indicates reaction spontaneity:
    • $∆H° = -92.6 kJ/mol$
    • $∆S° = -0.1985 kJ/(K ext{ mol})$
    • Computed $∆G = -33.45 kJ/mol$ denotes the reaction is spontaneous in the forward direction.