Thermochemistry

• Thermochemistry is the the heat (energy) absorbed or released during a chemical reaction at constant pressure and temperature when chemical binds are being broken or made.

Defining Enthalpy

• If a system releases heat, it is an exothermic process.
• In an exothermic reaction, the value of ΔH is less than zero, indicated by a negative sign.
• When a system absorbs heat, it is an endothermic process.
• In an endothermic reaction, the value of ΔH is more than zero, marked by a positive sign.

ΔH of Phase Changes

• Phase changes occur under constant pressure.
• The heat flow during a phase change matches the enthalpy change.
• During phase change, temperature remains constant as the heat flow goes towards forming or breaking intermolecular attractive forces
• The heat needed to change a liquid to a gas is the heat of vaporization and is represented as ΔHvap.
  • ΔHvap is endothermic with a positive value.
• The heat required to convert a gas to a liquid is the heat of condensation which is represented as ΔHcond
  • ΔHcond is exothermic with a negative value.
• ΔHcond = -ΔHvap
• Enthalpy changes in opposite directions share equal numeric values but differ in their signs only.
• Enthalpy change magnitude depends on the substance.
• The value of ΔH for a phase change is compound-specific and is measured in kJ/mol.
• The heat flow calculation uses the number of substance moles (n), and the value of the enthalpy change (ΔH = n × ΔHphase change)

Heat of Reaction

• Enthalpy changes calculations are applicable to chemical reactions as well as temperature changes and phase transitions.
• Enthalpy change is commonly referred to as the heat of reaction.

Bonds and Energy

• The estimation of the enthalpy change for a reaction uses bond energies.
• During a chemical reaction, reactant bonds are broken and product bonds are made.
  • Breaking bonds requires energy.
  • Making bonds releases energy.
• When energy released in product bond formation exceeds the energy needed to break reactant bonds, the reaction is exothermic. If less energy is released than required, the reaction is endothermic.
• The combustion of methane includes breaking of 4 C-H bonds and 2 O=O bonds and the making of 2 C=O bonds and 4 O-H bonds.
• The accuracy of enthalpy changes derived from tabulated bond energies is not very good - the bond energies used are averages.
• The bond energy method is suitable in estimating enthalpy changes for reactions involving compounds where no thermochemical data is available.
• A thermochemical equation conveys the overall energetics for a chemical reaction, where the sign on ΔH shows if a reaction is endothermic or exothermic.
• In the combustion of methane (CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)), ΔH = −890.4 kJ.
• The combustion of methane is an exothermic reaction, releasing 890.4 kJ of heat energy per mole of methane with 2 moles of oxygen.
• When stoichiometric coefficients in thermochemical equations are multiplied by a factor, the heat of reaction must be multiplied by the same factor.

Heat of Formation

• Heat of formation, ΔHf, defines the heat of reaction for the formation of substances.
• C(s) + ½ O2(g) → CO(g)
• Fractional coefficients are allowed in formation reactions as only one mole of product can be formed.
• A formation reaction is the chemical reaction in which one mole of a compound forms from its elements in their standard states.
• The standard state refers to the most stable state of an element at room temperature (25°C) and a pressure of 1 atm, denoted by a superscript.
• ΔHf° = 0 for elements in their standard state.

Hess's Law and Heats of Reaction

• Direct calorimetric determination of some reactions may be too difficult or dangerous to perform.
• An indirect method is therefore needed to obtain heats of reaction.
• Hess's law says the enthalpy change in any process is independent of the method by which the process is carried out.
• Enthalpy is a state function: a variable whose value depends only on the system's state, not the history of that state.
• Hess's law notes that enthalpy is a sate function, so any path is suitable for calculating the enthalpy change.
• For methane combustion, CH4 converts to CO, which then converts to CO2. This process uses ΔH for steps to determine the full ΔH .

Sample Problem & Formation Reactions

• SO3 originates from sulfur combustion in coal at the equation S(s) + 3/2 O2(g) → SO3(g).
• For thermochemical information for this, determine the heat of reaction via:
  • S(s) + O2(g) → SO2(g) (ΔH° = −296.8 kJ)
  • 2SO2(g) + O2(g) → 2SO3(g) (ΔH° = −197.0 kJ)
• Using tabulated formation enthalpies to calculate the enthalpy change for a chemical reaction, is expressed conceptually in a diagram.
• The calculation of the enthalpy change for a reaction, using Hess's law and heats of formation, uses the equation: ΔH°=∑νiΔHf°(products)i − ∑ vjΔHf°(reactants)j .

Variation of enthalpy with temperature

• Defined as Kirchoff's Equation:
  • Cp = (∂H/∂T)p
  • dH = CpdT
  • ∫dH = ∫T1=298.15T2 CpdT.
  • ΔHT2 = ΔHT1 + ∫Cp (net) dT
  • T₁ and T₂ are the system temperatures, with T₁ being the standard temperature (25 °C).
  • Cp (net) = ∑(viCp)products - ∑(vjCp)reactants = △Cp
• Cp approximation can be done as average heat capacity at constant pressure over a temperature range.
  • ΔH = CpΔT
  • ΔHT2 - ΔHT1 = Cp(T2 - 298.15K)
  • ΔHT2 = ΔHT1 + Cp(T2 - 298.15K)
• Cp variation with temperature is commonly represented as a power series in temperature.
  • Cp(T) = a + bT + cT² + dT³ + ...

Enthalpy Changes for Other Processes

• Enthalpy of combustion (heat of combustion) - enthalpy change when 1 mole of a compound is reacted completely with excess O2.
  • For example, CH₄(g) + 2 O2(g) → CO2(g) + 2H2O(l)
• Bond dissociation enthalpy (bond energies) represents the energy per mole needed to break a chemical bond.
  • For example, CH₄(g) → C(g) + 4 H(g)
• Enthalpy of solution (dilution) is the enthalpy change for dissolving 1 mole of a compound in a large excess of a pure solvent.
  • For example, HCl(g) + H₂O(excess) → HCl(aq)
• Enthalpy of neutralization - the reaction between a strong/weak acid and another strong/weak base to form salt and water, forming a neutralization reaction.
  • Enthalpy decreases with acid and base strength; it is exothermic (-ΔH) and the enthalpy value is calculated per mole of water produced.
  • Generic chemical equation: Acid + Base → Salt + Water