Physical Chemistry Lecture 3: Enthalpy of Reaction

Physical Chemistry (F/R-010) - Lecture 3

• Course Instructor: Prof. Dr. Hossieny Ibrahim
• Institution: Badr University in Assiut, School of Biotechnology
• Contact Information: [email protected], Bio-326

Unit 1: Enthalpy of a Reaction

• Thermochemical measurements can be taken at constant volume or constant pressure.
• The observed changes in these conditions differ.
• Change in internal energy (ΔE) is the heat change accompanying a chemical reaction at constant volume. This is because no external work is done.
• Most chemical reactions are conducted at constant pressure in the lab.
• To examine heat changes in reactions under constant pressure and temperature, enthalpy (H) is used.
• Enthalpy (H) is defined as the sum of internal energy (E) and the product of pressure (P) and volume (V). (H = E + PV)
• Enthalpy is a state function, meaning its value depends only on the current state of the system, not how it reached that state.
• Absolute enthalpy values cannot be measured, only changes in enthalpy (ΔH) can be determined.
• ΔH is calculated as the difference in enthalpy between products and reactants: ΔH = H_products - H_reactants = H_p - H_r

Unit 1: Enthalpy of a Reaction -Solids and Liquids

• For reactions involving solids and liquids, the change in volume (ΔV) is very small.
• Therefore, the term PΔV is negligible (ΔH ≈ ΔE).

Unit 1: Enthalpy of a Reaction - Gases

• For gas-based reactions, the change in volume (ΔV) is significant.
• Consequently, the term PΔV is appreciable.
• ΔH = ΔE + PΔV

Unit 1: Exothermic and Endothermic Reactions

• Reactions under constant pressure can either release or absorb heat.
• ΔH = 0: No heat absorption or release; products and reactants have the same enthalpy.
• ΔH < 0 (Negative): Products have lower enthalpy than reactants; heat is released (exothermic).
• ΔH > 0 (Positive): Products have higher enthalpy than reactants; heat is absorbed (endothermic).

Unit 1: Calculation of ΔH from ΔE

• ΔH = ΔE + PΔV
• Change in moles (Δn) = moles of products - moles of reactants.
• PΔV = ΔnRT (where R is the ideal gas constant and T is the temperature).
• Substituting, ΔH = ΔE + ΔnRT

Unit 1: Solved Problem - Ethylene Combustion

• The heat of combustion of ethylene at constant volume is -332.19 kcals at 17°C.
• Calculate the heat of combustion at constant pressure considering water as a liquid.
• Provided data: R = 2 cal degree⁻¹ mol⁻¹

Unit 1: Solved Problem - CO Combustion

• The heat of combustion of carbon monoxide at constant volume and 17°C is -283.3 kJ.
• Calculate the heat of combustion at constant pressure.

Unit 1: Thermochemical Equations

• A thermochemical equation indicates the heat change during a reaction.
• It must be balanced.
• Provide ΔE or ΔH values.
• Include physical states of reactants and products (e.g., (s) for solid, (l) for liquid, (g) for gas, (aq) for aqueous).

Unit 1: Thermochemical Equations -ΔH

• Is ΔH calculated negative or positive?
• Negative ΔH indicates an exothermic reaction (heat released) and positive ΔH represents an endothermic reaction (heat absorbed).

Unit 1: Thermochemical Equations Stoichiometry

• The stoichiometric coefficients in a thermochemical equation represent moles of substances.
• Reversing a reaction changes the sign of ΔH.
• Changing the equation's factors by a factor 'n' multiplies the ΔH value by 'n'.

Unit 1: Types of Heat (Enthalpy) of Reactions

• Heat of Formation (ΔH_f): Enthalpy change when one mole of a compound forms from constituent elements.
• Heat of Combustion (ΔH_c): Change in enthalpy when one mole of a substance undergoes complete combustion in excess air or oxygen.
• Heat of Solution (ΔH_solution): Change in enthalpy when one mole of a substance is dissolved in a specific quantity of solvent at a set temperature.

Example

• Heat of formation for ferrous sulfide (FeS) and acetylene (C₂H₂).
• Heat of formation for hydrogen chloride (HCl)