Thermodynamics and Energy in Chemical Reactions

Thermodynamics in Chemistry

Fundamental Laws Governing Chemical Change

Chemical changes are governed by fundamental conservation laws and energetic principles:

• Law of Conservation of Mass: The total mass of reactants equals the total mass of products in a chemical reaction.

• Law of Conservation of Energy: Energy is neither created nor destroyed in a chemical reaction, but it can be converted between kinetic energy (KE) and potential energy (PE).

Key energetic aspects of chemical reactions include:

• Thermodynamics: Studies energy changes (e.g., heat and enthalpy) and spontaneity of reactions.

• Kinetics: Examines the speed of reactions (not covered in detail here).

Thermodynamics

Heat of Reaction (ΔH) and Enthalpy

When reactions occur, energy is exchanged between the system and surroundings, often as heat (q). The enthalpy change (ΔH) is the heat transferred at constant pressure.

• ΔH > 0: Endothermic reaction (absorbs heat from surroundings)

• ΔH < 0: Exothermic reaction (releases heat to surroundings)

Exothermic Reactions

• Produce heat; energy is released to the surroundings.

• Example: Combustion of sucrose (a carbohydrate):

$\mathrm{C_{12}H_{22}O_{11} (s) + 12\ O_2 (g) \rightarrow 12\ CO_2 (g) + 11\ H_2O (l)}$

$\Delta H = -1342\ \mathrm{kcal}$ (energy released; surroundings get warmer)

Endothermic Reactions

• Absorb heat; energy is taken from the surroundings.

• Example: Dissolving ammonium nitrate in water:

$\mathrm{NH_4NO_3 (s) \rightarrow NH_4^+ (aq) + NO_3^- (aq)}$

$\Delta H > 0$ (energy absorbed; surroundings get colder)

Energy Transformations

• During chemical reactions, potential energy stored in chemical bonds is converted to kinetic energy (motion, heat).

• Exothermic: PE → KE (energy released)

• Endothermic: KE → PE (energy absorbed)

Entropy (ΔS)

Definition and Trends

Entropy (ΔS) is a measure of randomness or disorder in a system. Entropy increases when a system becomes more disordered.

• Solid → Liquid: $\Delta S > 0$ (entropy increases)

• Liquid → Gas: $\Delta S > 0$ (entropy increases)

• Solid → Gas: $\Delta S > 0$ (largest increase in entropy)

• Solid → Aqueous solution: $\Delta S > 0$ (dissolving increases disorder)

Higher entropy corresponds to lower energy and greater randomness.

Gibbs Free Energy (ΔG)

Definition and Equation

The Gibbs free energy (ΔG) determines whether a reaction is spontaneous. It combines enthalpy and entropy changes:

$\Delta G = \Delta H - T\Delta S$

• ΔG < 0: Reaction is spontaneous (exergonic); does not require extra energy once started.

• ΔG > 0: Reaction is nonspontaneous (endergonic); requires energy input to proceed.

Exergonic vs. Endergonic Reactions

• Exergonic: $\Delta G < 0$, often exothermic ($\Delta H < 0$), spontaneous.

• Endergonic: $\Delta G > 0$, often endothermic ($\Delta H > 0$), nonspontaneous.

Example: ATP Hydrolysis

$\mathrm{ATP \rightarrow ADP + Energy}$, $\Delta G < 0$

This reaction releases energy used for cellular processes, such as moving ions across membranes.

Energy Diagrams

Visualizing Energy Changes

Energy diagrams show how energy changes during a reaction:

• Exergonic (Spontaneous): Products have lower energy than reactants; $\Delta G < 0$.

• Endergonic (Nonspontaneous): Products have higher energy than reactants; $\Delta G > 0$.

• Activation Energy (Ea): The energy barrier that must be overcome for a reaction to proceed.

Exothermic reactions: $\Delta H < 0$ Endothermic reactions: $\Delta H > 0$

Energy Content of Food

Caloric Values and Calculations

The nutritive value of food is the amount of energy it produces when burned (measured in Calories or kilojoules).

Food Type	Calories per gram (Cal)	Energy per gram (kJ)
Carbohydrate	4	17
Protein	4	17
Fat	9	38

Example Calculation: Estimate the energy content in Calories of 2 eggs (2 g carbohydrates, 10 g fat, 12 g protein):

• 2 g carbohydrates × 4 Cal/g = 8 Cal

• 10 g fat × 9 Cal/g = 90 Cal

• 12 g protein × 4 Cal/g = 48 Cal

• Total = 8 + 90 + 48 = 146 Cal

*Additional info: The calculation assumes the standard caloric values for each macronutrient and does not account for digestibility or metabolic differences.*