Comprehensive Study Guide: Enthalpy and Enthalpy Change

Energy in Chemical Systems and Thermodynamics

• First Law of Thermodynamics: During a chemical reaction, energy cannot be created or destroyed, but it can be transferred from one form to another.

• Chemical Energy: Composed of two fundamental energy components:

  • Kinetic Energy: A measure of the motion of particles (atoms, molecules, or ions) within a substance.

  • Potential Energy: A measure of how strongly particles interact with one another through both attractive and repulsive forces.

• Heat Energy:

  • The portion of potential energy and kinetic energy in a substance responsible for its temperature.

  • Directly proportional to the absolute temperature measured in Kelvin ($K$).

• Distinction Between Heat Energy and Heating:

  • Heating is the transfer of heat energy from one system to another, resulting in a temperature change.

  • Example: When water at $60\,^\circ\text{C}$ is placed in contact with air at $20\,^\circ\text{C}$, heat energy is transferred from the water to the air. This causes an increase in the temperature of the air and a corresponding decrease in the temperature of the water, explaining why hot beverages cool down when left in room air.

Enthalpy and Enthalpy Change

• Enthalpy ($H$):

  • A measure of the total energy of a system.

  • Mathematically defined by the equation:     $H = U + pV$     where $H$ is enthalpy, $U$ is internal energy (the energy required to create the system), $p$ is the pressure of the system, and $V$ is the volume of the system.

  • Enthalpy combines internal energy with the amount of energy required to make room for the system by displacing its surrounding environment and establishing its volume and pressure.

• System and Surroundings:

  • System: Refers to the chemical reaction mixture.

  • Surroundings: Refers to everything outside the system, which in laboratory practice is the surrounding air in the room.

• Enthalpy Change ($\Delta H$):

  • The absolute enthalpy ($H$) of a system cannot be determined directly.

  • Enthalpy change ($\Delta H$) can be measured during physical or chemical changes.

  • Definition: The heat energy transferred between the system and the surroundings measured at constant pressure.

Exothermic and Endothermic Processes

• Exothermic Processes:

  • Definition: A process or reaction in which heat energy is transferred from the system to the surroundings.

  • Sign Convention: Indicated by a negative sign ($-\Delta H$), representing heat energy transferred away from the system.

  • Temperature Effect: Produces an immediate increase in the temperature of the reaction mixture or surroundings.

  • Bond Formation: Bond formation is an exothermic process. For example, when water vapour condenses, hydrogen bonds form between water molecules, releasing heat.

  • Specific Reaction Example:     $\text{HCl}(aq) + \text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l) \quad \Delta H = -57.1\,kJ\,mol^{-1}$

  • Comprehensive List of Exothermic Examples:

  • Freezing water

  • Condensing water vapour

  • Dissolving sodium hydroxide in water

  • Reaction between dilute hydrochloric acid and aqueous sodium hydroxide

  • Combustion of petrol

  • Burning natural gas in oxygen (producing a hot flame)

  • Burning candle wax

  • Formation of snow in clouds

• Endothermic Processes:

  • Definition: A process or reaction in which heat energy is transferred from the surroundings to the system.

  • Sign Convention: Indicated by a positive sign ($+\Delta H$), representing heat energy transferred from the surroundings into the system.

  • Temperature Effect: Produces a decrease in temperature of the reaction mixture, or requires continuous heating from an external source for the reaction to take place.

  • Specific Reaction Examples:

  • Reaction between citric acid and solid sodium hydrogencarbonate:       $\text{C}_6\text{H}_8\text{O}_7(aq) + 3\text{NaHCO}_3(s) \rightarrow \text{C}_6\text{H}_5\text{O}_7^{3-}(aq) + 3\text{Na}^+(aq) + 3\text{CO}_2(g) + 3\text{H}_2\text{O}(l) \quad \Delta H = +70\,kJ\,mol^{-1}$

  • Thermal decomposition of calcium carbonate into calcium oxide and carbon dioxide:       $\text{CaCO}_3(s) \rightarrow \text{CaO}(s) + \text{CO}_2(g) \quad \Delta H = +178\,kJ\,mol^{-1}$

  • Comprehensive List of Endothermic Examples:

  • Melting ice

  • Evaporating water

  • Dissolving ammonium nitrate in water

  • Reaction between dilute ethanoic acid and solid sodium hydrogencarbonate

  • Reaction between citric acid and solid sodium hydrogencarbonate

  • Photosynthesis

  • Cooking an egg

  • Baking bread

  • Forming a cation from an atom in the gas phase

  • Sherbet reaction in saliva: Sherbet is a solid mixture of sodium hydrogencarbonate and citric acid. When mixed with saliva in the mouth, the solid mixture dissolves and undergoes an endothermic reaction, absorbing heat energy from the mouth and causing a cold sensation.

Standard Conditions and Measurement

• IUPAC Standard Conditions:

  • Recommended in 1982 by the International Union of Pure and Applied Chemistry (IUPAC).

  • Standard Pressure: $100\,kPa$

  • Standard Temperature: A stated temperature, most commonly $298\,K$

  • Historical / Non-IUPAC Standard Pressure: Some textbooks or data books quote standard enthalpy changes at $1\,\text{atm} = 101.325\,kPa$. This is not in agreement with IUPAC recommendations, which specify $100\,kPa$ as standard for thermochemical measurements.

  • Notation: Standard enthalpy change measured under these conditions is denoted as $\Delta H_{298\text{K}}$ or simply $\Delta H$

Standard Enthalpy Change of Reaction

• Standard Enthalpy Change of Reaction ($\Delta_r H$):

  • Definition: The enthalpy change which occurs when equation quantities of materials react under standard conditions.

  • Equation Specificity: Standard enthalpy change of reaction is calculated for the reaction strictly as written in the equation.

  • Interpretation of "per mole" ($kJ\,mol^{-1}$): Refers to one mole of the reaction equation as written, not to one mole of any specific reactant or product.

  • Comparison of Stoichiometric Variations:

  • Reaction written for formation of two moles of ammonia:       $\text{N}_2(g) + 3\text{H}_2(g) \rightarrow 2\text{NH}_3(g) \quad \Delta_r H = -92\,kJ\,mol^{-1}$

  • Reaction written for formation of one mole of ammonia:       $\frac{1}{2}\text{N}_2(g) + \frac{3}{2}\text{H}_2(g) \rightarrow \text{NH}_3(g) \quad \Delta_r H = -46\,kJ\,mol^{-1}$

  • State Symbols: State symbols ($(s)$, $(l)$, $(g)$, $(aq)$) must always be included when writing thermochemical equations.