Thermodynamics of Reactions

ΔG, ΔH, and ΔS in Organic Chemistry

In organic chemistry, understanding a reaction's favorability and direction relies on three key concepts: Gibbs Free Energy (ΔG), Enthalpy (ΔH), and Entropy (ΔS). These are interrelated and tell us a story about the energy landscape of the reaction.

ΔG: The Driving Force (Gibbs Free Energy)

Imagine ΔG as the net usable energy in a reaction. A negative ΔG indicates a spontaneous reaction, meaning it can proceed without external input and tends to reach a state of equilibrium with more products. Conversely, a positive ΔG suggests a non-spontaneous reaction that wouldn't occur on its own.  A negative ΔG signifies an exoergonic reaction, while a positive 

ΔH: Heat Flow (Enthalpy) and Bond Energies

ΔH represents the heat exchange between the reaction and its surroundings. A negative ΔH signifies an exothermic reaction, which releases heat as the reaction progresses. An endothermic reaction, with a positive ΔH, absorbs heat from the surroundings.

We can estimate ΔH for a reaction using bond enthalpies, which are the average energy required to break a specific type of bond. For example, consider the combustion of methane:

CH₄ (g) + 2O₂ (g) → CO₂ (g) + 2H₂O (g)

Here, we look up the bond enthalpies (available in reference tables) for each bond broken (C-H, O=O) and formed (C=O, O-H).

ΔH ≈ Σ (Bond enthalpies of broken bonds) - Σ (Bond enthalpies of formed bonds)

Example Calculation:

Bond Enthalpy (kJ/mol) Change
C-H 436 -4(436) (4 C-H bonds broken in CH₄)
O=O 498 -1(498) (1 O=O bond broken in O₂)
C=O 799 -(799) (1 C=O bond formed in CO₂)
O-H 463 2(463) (2 O-H bonds formed in 2H₂O)

ΔH ≈ (-4(436) - 498) - (-799 + 2(463)) ≈ -502 kJ/mol

This negative ΔH indicates an exothermic reaction, releasing 502 kJ of energy per mole of methane combusted.

ΔS: The Disorder Factor (Entropy)

Entropy is a measure of randomness or disorder in a system.

A positive ΔS indicates an increase in disorder, often associated with:

  • Formation of more gas molecules: Gas molecules have more freedom of movement compared to solids or liquids, leading to higher entropy.

    For example, the decomposition of calcium carbonate: CaCO₃ (s) → CaO (s) + CO₂ (g) Here, the formation of a gas molecule (CO₂) increases the entropy (ΔS > 0).

  • Dissolving a solid in a liquid:  Solid particles are more ordered than dispersed particles in solution, so dissolving a solid increases entropy (ΔS > 0).

Conversely, a negative ΔS suggests a decrease in disorder, such as:

  • Formation of polymers from monomers:** Monomers have more freedom of movement than covalently linked polymer chains, resulting in a decrease in entropy (ΔS < 0).

The Interplay: ΔG Connects ΔH and ΔS

The magic lies in how ΔG considers both ΔH and ΔS:

ΔG = ΔH - TΔS (where T is temperature)

This equation tells us that a favorable reaction (negative ΔG) can result from either a negative ΔH (exothermic) or a positive ΔS (increased disorder), or ideally, a combination of both. Even an endothermic reaction (positive ΔH) might be spontaneous if the increase in entropy (positive ΔS) is significant enough at a particular temperature (T).

Understanding ΔG, ΔH, and ΔS empowers organic chemists to predict reaction feasibility, optimize reaction conditions, and design new synthetic pathways for organic molecules.

 

Equilibrium Constant (K): The Balancing Act

The equilibrium constant (K) is a temperature-dependent constant that reflects the ratio of product concentrations to reactant concentrations at equilibrium. For a general reaction:

aA + bB <=> cC + dD

K = [C]^c [D]^d / [A]^a [B]^b

where:

  • [A], [B], [C], and [D] represent the equilibrium molar concentrations of reactants and products, respectively.
  • a, b, c, and d are the stoichiometric coefficients from the balanced chemical equation.

A high K value (greater than 1) indicates a product-favored equilibrium, meaning there are more products than reactants at equilibrium. Conversely, a K value less than 1 signifies a reactant-favored equilibrium. If K is equal to 1, the concentrations of reactants and products are equal at equilibrium.

  • ΔG and K: Unveiling the Connection

The magic lies in the relationship between ΔG and K, described by the following equation:

ΔG = -RTln(K)

or

K = e(-RT/ΔG)

where:

  • R is the gas constant (8.314 J/mol*K)
  • T is the absolute temperature (Kelvin)
  • ln(K) is the natural logarithm of K

This equation tells us that the value of K is related tp the sign and magnitude of ΔG:

  • Negative ΔG (Spontaneous): When ΔG is negative, the reaction leans towards products, and ln(K) will be a positive value. Consequently, K will be greater than 1, indicating a product-favored equilibrium.
  • Positive ΔG (Non-spontaneous): A positive ΔG suggests the reaction favors reactants. ln(K) becomes negative, resulting in a K value less than 1, signifying a reactant-favored equilibrium.
  • ΔG = 0 (Equilibrium): At equilibrium, ΔG is zero, and ln(K) becomes zero. This translates to K = 1, indicating equal concentrations of reactants and products.

By calculating ΔG using thermochemical data or estimating it from bond enthalpies, and knowing the temperature, we can predict the equilibrium constant (K) and understand the extent of a reaction at equilibrium in organic chemistry.