Chapter 16 Notes: Reaction Energy and Driving Forces

BIG IDEA: Energy in Chemical Reactions

• Energy is released or absorbed during chemical reactions.

• Nature favors reactions proceeding towards the lowest energy state.

Section 1: Thermochemistry

• Thermochemistry: Study of energy transfers as heat accompanying chemical reactions and physical changes.

• Law of Conservation of Energy: Energy is neither created nor destroyed, only changes forms.

• Key Terms:

  • Thermochemistry.

  • Heat.

  • Thermochemical equation.

  • Calorimeter.

  • Specific heat.

  • Molar enthalpy of formation.

  • Temperature.

  • Enthalpy change.

  • Enthalpy of combustion.

  • Joule.

  • Enthalpy of reaction.

  • Hess’s law.

Temperature and Heat

• Temperature: Measure of average kinetic energy of particles.

• Heat: Energy transferred between matter samples due to temperature difference.

• Temperature Scales:

  • Celsius and Kelvin are used in thermochemistry.

  • $K = 273.15 + °C$ (often rounded to $K = 273 + °C$).

• Joule (J): SI unit of heat and energy, derived from mass, distance, and time. $J = \frac{kg \cdot m^2}{s^2}$.

Energy Transfer

• Heat transfer occurs from higher to lower temperature.

• Energy transfer stops when temperatures equalize.

• Specific Heat (cp): Energy to raise 1 gram of a substance by 1°C or 1 K.

• Units: J/(g • °C) or J/(g • K) or cal/(g • °C).

• Water has a high specific heat, crucial for temperature regulation.

Specific Heat Equation:

• $c_p = \frac{q}{m \cdot \Delta T}$, where:

  • $c_p$ is specific heat at constant pressure,

  • $q$ is energy lost or gained,

  • $m$ is mass,

  • $\Delta T$ is temperature change.

Energy Lost or Gained Equation:

• $q = c_p \cdot m \cdot \Delta T$

Enthalpy Change (ΔH)

• Represents energy absorbed as heat during a chemical reaction at constant pressure.

• Only changes in enthalpy can be measured, not absolute enthalpy.

• $\Delta H = H<em>{\text{products}} - H</em>{\text{reactants}}$

Enthalpy of Reaction

• Quantity of energy transferred as heat during a chemical reaction (heat of reaction).

Exothermic Reactions

• Energy is released; products have less energy than reactants.

• Example:

  • $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(g) + 483.6 \text{ kJ}$

• Thermochemical equation: Includes energy released or absorbed as heat.

• Coefficients represent moles, not molecules.

Endothermic Reactions

• Energy is absorbed; products have higher enthalpy than reactants.

• Example:

  • $2H<em>2O(g) + 483.6 \text{ kJ} \rightarrow 2H</em>2(g) + O_2(g)$

• Physical states are important in thermochemical equations.

Thermochemical Equations

• Represent the value of $\Delta H$.

• Exothermic: $\Delta H$ is negative (energy lost).

• Endothermic: $\Delta H$ is positive (energy gained).

• Example:

  • $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(g) \quad \Delta H = -483.6 \text{ kJ}$

  • $2H<em>2O(g) \rightarrow 2H</em>2(g) + O_2(g) \quad \Delta H = +483.6 \text{ kJ}$

• Coefficients represent moles, physical states are crucial, $\Delta H$ is proportional to the number of moles, and $\Delta H$ isn't significantly affected by temperature changes.

Molar Enthalpy of Formation

• Enthalpy change when one mole of a compound is formed from its elements in their standard state (25°C and 1 atm).

• Standard state: State at atmospheric pressure and room temperature.

• Symbol: $\Delta H_f^0$

• Elements in standard states have $\Delta H_f^0 = 0$.

• Large negative $\Delta H_f^0$ means the compound is stable.

• Positive $\Delta H_f^0$ implies instability (e.g., hydrogen iodide).

• Mercury fulminate, $\text{HgC}<em>2\text{N}</em>2\text{O}_2$, has a very large enthalpy of formation of $+270 \text{ kJ/mol}$, and is very unstable

Enthalpy of Combustion

• Enthalpy change during complete combustion of one mole of a substance.

• Symbol: $\Delta H_c$

• Defined in terms of one mole of reactant.

• Measured using a combustion calorimeter.

Hess’s Law

• Overall enthalpy change equals the sum of enthalpy changes for individual steps.

• Energy difference between reactants and products is independent of the pathway.

• Used to calculate reaction enthalpies that are difficult to measure directly.

• Combining thermochemical equations:

  1. If a reaction is reversed, the sign of $\Delta H$ is reversed.

  2. Multiply coefficients and $\Delta H$ to match the desired thermochemical equation.

• Total Enthalpy Change:

  • $\Delta H^0 = \sum [(\Delta H<em>f^0 \text{ of products}) \times (\text{mol of products})] - \sum [(\Delta H</em>f^0 \text{ of reactants}) \times (\text{mol of reactants})]$

Hess’s Law - Example

• Formation of methane gas from its elements:

  • $C(s) + 2H<em>2(g) \rightarrow CH</em>4(g) \quad \Delta H_f^0 = ?$

• Combustion reactions:

  • $C(s) + O<em>2(g) \rightarrow CO</em>2(g) \quad \Delta H_c^0 = -393.5 \text{ kJ}$

  • $H<em>2(g) + \frac{1}{2} O</em>2(g) \rightarrow H<em>2O(l) \quad \Delta H</em>c^0 = -285.8 \text{ kJ}$

  • $CH<em>4(g) + 2O</em>2(g) \rightarrow CO<em>2(g) + 2H</em>2O(l) \quad \Delta H_c^0 = -890.8 \text{ kJ}$

Section 2: Driving Force of Reactions

Key Terms

• Entropy.

• Free energy.

• Free-energy change.

Reactions and Energy States

• Reactions proceed towards a lower-energy state.

• Exothermic reactions release energy; products are more stable.

• Endothermic reactions absorb energy; products are less stable.

• Something other than enthalpy change determines whether a reaction will occur.

Entropy (S)

• Measures randomness or disorder in a system.

• Tendency in nature to increase randomness.

  • $2NH<em>4NO</em>3(s) \rightarrow 2N<em>2(g) + 4H</em>2O(l) + O_2(g)$

• Solids have low entropy; gases have high entropy.

• Entropy of pure crystalline solid at absolute zero is zero.

Entropy Change

• $\Delta S > 0$: Increase in entropy.

• $\Delta S < 0$: Decrease in entropy.

• Mixing increases entropy.

Free Energy (G)

• Combines enthalpy and entropy factors to predict spontaneity.

• Natural processes lower free energy.

• Free-Energy Change ($\Delta G$):

  • $\Delta G^0 = \Delta H^0 - T\Delta S^0$

  • Where T is in Kelvin and standard states are assumed.

  • If $\Delta G < 0$, the reaction is spontaneous.

• $\Delta H$ and $\Delta S$ can be positive or negative, leading to four possible combinations.

Relating Enthalpy, Entropy, and Free-Energy Changes to Reaction Occurrence

• Negative $\Delta H$ (exothermic), positive $\Delta S$ (more random) always results in negative $\Delta G$ (spontaneous)

• The reaction rate is too slow for the spontaneous change to be observed.