Thermodynamics is all about energy changes in chemical reactions. We'll look at enthalpy (heat flow), entropy (disorder), and Gibbs free energy (spontaneity) to understand why reactions happen.
These concepts help us predict if a reaction will occur on its own. We'll see how temperature affects reactions and learn to calculate energy changes using some handy equations.
Enthalpy Changes in Reactions
Defining Enthalpy and Its Properties
- Enthalpy (H) measures the total heat content of a system at constant pressure
- State function with units of energy (joules or kilojoules)
- The change in enthalpy (ΔH) is the heat absorbed or released by a system during a process at constant pressure
- Positive ΔH indicates an endothermic process (heat absorbed)
- Negative ΔH indicates an exothermic process (heat released)
Calculating Enthalpy Changes
- Hess's law states the overall enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps, regardless of the pathway
- Standard enthalpy of formation (ΔH°f) is the enthalpy change when one mole of a compound forms from its constituent elements in their standard states (1 atm, usually 25°C)
- Enthalpy changes can be calculated for various chemical processes:
- Phase changes (melting, vaporization)
- Chemical reactions
- Formation or breaking of chemical bonds
Entropy and Disorder
Understanding Entropy
- Entropy (S) measures the degree of disorder or randomness in a system
- State function with units of energy per temperature (joules per kelvin)
- Related to the number of possible particle arrangements in a system
- The second law of thermodynamics states the total entropy of the universe always increases during a spontaneous process
- The degree of disorder in the universe is constantly increasing
Entropy and Energy Dispersal
- Entropy is related to the dispersal of energy within a system
- As energy disperses more evenly, the entropy of the system increases
- The change in entropy (ΔS) for a process can be calculated using the equation: $ΔS = q/T$
- q is the heat absorbed or released by the system during a reversible process
- T is the absolute temperature in Kelvin
- Entropy changes during phase transitions:
- Increases when a system goes from a more ordered to a less ordered state (melting, vaporization)
- Decreases during the reverse processes (freezing, condensation)
Spontaneity of Processes
Gibbs Free Energy and Spontaneity
- Gibbs free energy (G) combines the effects of enthalpy and entropy to determine the spontaneity of a process at constant temperature and pressure
- Defined as: $G = H - TS$
- The change in Gibbs free energy (ΔG) for a process can be calculated using the equation: $ΔG = ΔH - TΔS$
- Negative ΔG indicates a spontaneous process
- Positive ΔG indicates a non-spontaneous process
- ΔG of zero indicates a system at equilibrium
Factors Influencing Spontaneity
- The spontaneity of a process depends on the relative magnitudes of the enthalpy and entropy changes:
- Negative ΔH (exothermic) and positive ΔS (increasing disorder) always spontaneous
- Positive ΔH (endothermic) and negative ΔS (decreasing disorder) always non-spontaneous
- Temperature plays a crucial role when ΔH and ΔS have opposite signs:
- At low temperatures, the TΔS term is small, and the sign of ΔG is determined by the sign of ΔH
- At high temperatures, the TΔS term becomes more significant, and the sign of ΔG is determined by the sign of ΔS
Gibbs Free Energy and Equilibrium
Relating Gibbs Free Energy to Equilibrium Constant
- The change in Gibbs free energy (ΔG) for a chemical reaction is related to the equilibrium constant (K) by the equation: $ΔG° = -RTlnK$
- ΔG° is the standard Gibbs free energy change
- R is the universal gas constant (8.314 J/mol·K)
- T is the absolute temperature in Kelvin
- At equilibrium, ΔG = 0, and the equation becomes: $0 = -RTlnK$
- Allows for the calculation of the equilibrium constant from the standard Gibbs free energy change
Interpreting Equilibrium Constants
- A negative ΔG° indicates a spontaneous reaction and corresponds to an equilibrium constant greater than 1 (K > 1)
- A positive ΔG° indicates a non-spontaneous reaction and corresponds to an equilibrium constant less than 1 (K < 1)
- The magnitude of the equilibrium constant provides information about the extent of the reaction at equilibrium:
- Large K value (K >> 1) indicates products are favored
- Small K value (K << 1) indicates reactants are favored
- The relationship between ΔG° and K can be used to:
- Predict the direction of a reaction
- Determine relative concentrations of reactants and products at equilibrium
- Calculate the standard Gibbs free energy change from experimentally determined equilibrium constants