Thermodynamics - CHEMISTRY

Definition

A system is the specific part of the universe under observation, while the surroundings are the rest of the universe that can interact with the system. Together, the system and surroundings make up the universe.

Explanation

• Open System: Can exchange both matter and energy with the surroundings (e.g., an open pan of boiling water).
• Closed System: Can exchange only energy, not matter, with the surroundings (e.g., a sealed metal can of soda in a refrigerator).
• Isolated System: Can exchange neither matter nor energy with the surroundings (e.g., a tightly sealed thermos flask).

Examples

• A living human body is an open system taking in food and oxygen while releasing heat and sweat.
• A sealed, insulated vacuum flask keeping soup hot approximates an isolated system.

Definition

The total energy stored inside a thermodynamic system, represented by the symbol U, comprising microscopic kinetic and potential energies.

Explanation

• Internal energy is a state function depending only on the current state of the system.
• Absolute internal energy cannot be measured directly, but changes ($\Delta U$) can be measured when heat or work is exchanged.

Examples

• Rubbing hands together increases the internal energy of skin molecules due to friction.
• Burning wood converts chemical potential energy into thermal energy.

Definition

Energy can neither be created nor destroyed, it can only be transformed from one form to another.

Explanation

• Mathematical Expression: $\Delta U = q + w$
• $q$ is heat added to the system; $w$ is work done on the system.

Examples

• Inflating a bicycle tire with a hand pump does work on the air and generates heat.

Numerical Illustration

delta U = q + w
If q = +250 J and w = -100 J:
delta U = 250 + (-100) = +150 Joules.

Definition

A thermodynamic property defined as $H = U + pV$, representing the total heat content of a system at constant pressure.

Explanation

• At constant pressure, the change in enthalpy equals heat transferred ($\Delta H = q_p$).
• Exothermic reactions have a negative $\Delta H$; endothermic reactions have a positive $\Delta H$.

Examples

• Melting of ice is endothermic ($\Delta H > 0$).
• Burning natural gas is exothermic ($\Delta H < 0$).

Definition

The amount of heat required to raise the temperature of a given amount of a substance by one degree Celsius (or Kelvin).

Explanation

• Formula: $q = m \cdot c \cdot \Delta T$
• Water has a high specific heat capacity, resisting rapid temperature changes.

Examples

• Water is used in car radiator coolants due to its high specific heat capacity.

Definition

Calorimetry is the measurement of heat transfer during physical changes or chemical reactions using a calorimeter.

Explanation

• Bomb Calorimeter: Constant-volume device used to measure $\Delta U$.
• Coffee-Cup Calorimeter: Constant-pressure device used to measure $\Delta H$.

Definition

If a reaction can be expressed as a series of steps, the total enthalpy change for the overall reaction is equal to the sum of the enthalpy changes for the individual steps.

Explanation

• Direct consequence of enthalpy being a state function.
• Reversing an equation reverses the sign of $\Delta H$; multiplying an equation multiplies $\Delta H$ by the same factor.

Definition

Specific enthalpy terms assigned to distinct chemical or physical processes (e.g., formation, combustion, atomization, sublimation, phase transition, ionization, solution, and dilution).

Explanation

• Standard enthalpy of formation ($\Delta H_f^\circ$) of any pure element in its reference state is zero.
• Bond breaking absorbs energy; bond formation releases energy.

Definition

The total entropy of an isolated system must always increase for any spontaneous process ($\Delta S_{\text{universe}} > 0$).

Explanation

• Explains the natural directionality and irreversibility of processes.

Definition

Entropy ($S$) is a thermodynamic property measuring the degree of randomness or disorder of particles in a system.

Explanation

• Formula: $\Delta S = \frac{q_{\text{rev}}}{T}$
• Solids have low entropy, liquids have moderate entropy, and gases have high entropy.

Definition

Gibbs Free Energy ($G = H - TS$) combines enthalpy and entropy to predict reaction spontaneity at constant temperature and pressure.

Explanation

• $\Delta G < 0$: Spontaneous
• $\Delta G > 0$: Non-spontaneous
• $\Delta G = 0$: Equilibrium

Problem 1

Calculate the standard enthalpy change ($\Delta H$) for the reaction: $4\text{Fe}(s) + 3\text{O}_2(g) \rightarrow 2\text{Fe}_2\text{O}_3(s)$ given that $\Delta H_f^\circ(\text{Fe}_2\text{O}_3) = -824.2\text{ kJ/mol}$.

Step 1: delta H_rxn = sum(delta H_f products) - sum(delta H_f reactants)
Step 2: delta H_rxn = [ 2 * (-824.2) ] - [ 0 ] = -1648.4 kJ.

Problem 2

For a reaction, $\Delta H = -35.5\text{ kJ/mol}$ and $\Delta S = -85.0\text{ J/K}\cdot\text{mol}$. At what temperature range will this reaction become spontaneous?

Step 1: delta G = delta H - T * delta S < 0 => T * delta S > delta H
Step 2: delta H = -35500 J/mol, delta S = -85.0 J/K*mol
Step 3: T < (-35500) / (-85.0) => T < 417.65 Kelvin.

• First Law: $\Delta U = q + w$
• Enthalpy definition: $H = U + pV$
• Heat transfer relation: $q = m \cdot c \cdot \Delta T$
• Enthalpy change at constant pressure: $\Delta H = q_p$
• Entropy change definition: $\Delta S = \frac{q_{\text{rev}}}{T}$
• Gibbs Free Energy equation: $\Delta G = \Delta H - T \cdot \Delta S$

Modern chemists utilize Gaussian/ORCA for computational chemistry, Differential Scanning Calorimetry (DSC) for polymer analysis, and Aspen Plus/HYSYS for industrial process simulation.

The Haber-Bosch process ($\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$) is exothermic ($\Delta H \approx -92.2\text{ kJ/mol}$) with negative entropy change. Engineers balance moderate temperatures (400-500°C), high pressures (150-250 atm), iron catalysts, and heat exchangers to optimize yield and energy efficiency.

1. Which of the following is an extensive property?
   Answer: C) Enthalpy
   Explanation: Extensive properties depend on the amount of matter present.
2. For an isolated system, the change in internal energy ($\Delta U$) is always:
   Answer: C) Zero
   Explanation: An isolated system cannot exchange energy or matter.
3. In an exothermic reaction, the value of $\Delta H$ is:
   Answer: B) Negative
   Explanation: Exothermic reactions release heat, lowering system enthalpy.
4. Which state function represents molecular disorder?
   Answer: C) Entropy
   Explanation: Entropy ($S$) measures particle randomness.
5. A reaction is guaranteed spontaneous at all temperatures if:
   Answer: C) $\Delta H$ is negative and $\Delta S$ is positive
   Explanation: Makes $\Delta G$ negative unconditionally.
6. What device measures enthalpy changes at constant pressure?
   Answer: B) Coffee-cup calorimeter
   Explanation: Operates under open atmospheric conditions.
7. Hess's Law is a consequence of:
   Answer: B) Enthalpy being a state function
   Explanation: Path independence allows step-by-step summation.
8. The third law states entropy is zero at:
   Answer: C) Absolute zero (0 K)
   Explanation: Thermal motion ceases entirely.
9. When ice melts at room temperature, $\Delta S$ for the water is:
   Answer: A) Positive
   Explanation: Solid turns into a more disordered liquid.
10. If $\Delta G$ for a reaction is positive, the reaction is:
    Answer: C) Non-spontaneous
    Explanation: Indicates non-spontaneity in the forward direction.

1. State the First Law of Thermodynamics and write its mathematical expression.
   Answer: Energy can neither be created nor destroyed. Expression: $\Delta U = q + w$.
2. Why is internal energy a state function while work is a path function?
   Answer: Internal energy depends only on the current state, whereas work depends on the pathway taken between states.
3. Differentiate between open, closed, and isolated systems with examples.
   Answer: Open exchanges matter and energy (boiling water in open pan); closed exchanges only energy (sealed flask on heater); isolated exchanges neither (thermos flask).
4. Why is standard enthalpy of formation for pure elements in reference state taken as zero?
   Answer: Absolute enthalpy cannot be measured, so pure stable elements are set as a reference baseline of zero.
5. State Hess's Law and give its significance.
   Answer: Total enthalpy change is independent of pathway. It enables calculation of enthalpy changes for difficult or dangerous reactions.
6. How does the Second Law explain spontaneous processes?
   Answer: Total entropy of an isolated system (universe) must always increase ($\Delta S_{\text{universe}} > 0$).
7. Under what temperature conditions will a reaction with positive $\Delta H$ and positive $\Delta S$ be spontaneous?
   Answer: Only at high temperatures where the $T\Delta S$ term exceeds $\Delta H$.
8. Define bond enthalpy and explain why average values are used.
   Answer: Energy required to break one mole of covalent bonds in gas phase. Average values account for variations across different molecular environments.
9. What is the physical significance of Gibbs Free Energy ($\Delta G$)?
   Answer: Determines spontaneity ($\Delta G < 0$) and measures maximum non-expansion work obtainable from a system.
10. Calculate $\Delta U$ for a system that absorbs 250 J of heat and does 100 J of work on surroundings.
    Answer: $q = +250\text{ J}$, $w = -100\text{ J}$. $\Delta U = 250 + (-100) = +150\text{ J}$.

• Universe = System + Surroundings
• First Law: $\Delta U = q + w$
• Enthalpy: $H = U + pV$ ($\Delta H = q_p$)
• Gibbs Energy: $\Delta G = \Delta H - T\Delta S$ (Negative = Spontaneous)

Mnemonics: "Good Boys Go Spontaneous" (Gibbs $G < 0$ is Spontaneous). "Exit heat = Exothermic" (Negative $\Delta H$).