Understanding Enthalpy Change (∆H°)

Enthalpy Change (∆H°)

• Enthalpy change represents the change in heat energy, symbolized as ∆H°.
• Enthalpy is measured under standard conditions: 100 kPa pressure and a specified temperature, usually 298 K.
• During a reaction, energy is absorbed to break bonds and released when bonds are formed. The overall energy change relies on the energy transfer in these processes.
• An enthalpy change is positive when energy is absorbed from the surroundings.
• An enthalpy change is negative when energy is released into the surroundings.
• The overall enthalpy change (∆H) is calculated by adding the energy needed to break bonds (positive value) to the energy released when making bonds (negative value).
• Energy level diagrams visually represent enthalpy changes in a reaction, indicating whether the reaction is endothermic or exothermic.
• Endothermic reactions require more energy to break bonds than to form them, resulting in a positive overall ∆H and heat being absorbed from the surroundings.
• Exothermic reactions release heat and have a negative ∆H, requiring more energy to form new bonds than break existing ones.

Measuring Enthalpy Change

• The enthalpy change of reaction (∆rHo) refers to the enthalpy change when quantities of substances in standard states react completely under standard conditions.
• The enthalpy change of formation (∆fHo) is the enthalpy change when one mole of a substance is produced from its elements under standard conditions.
• The enthalpy change of combustion (∆cHo) refers to the enthalpy change when one mole of a substance is burned completely in oxygen under standard conditions.
• The enthalpy change of neutralization (∆neutHo) is the enthalpy change when solutions of acid and alkali react together under standard conditions to produce one mole of water.
• The enthalpy change of atomization (∆atHo) refers to the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.

Calorimetry

• Calorimetry serves as an experimental method to determine enthalpy change by measuring temperature change over time.
• Plotted data can be extrapolated to provide a precise temperature change value at the beginning of the reaction.
• The measured change in temperature, ∆T, is proportional to the energy change, represented by the equation q = mc∆T, where q is energy change (J), m is mass (g), c is specific heat capacity (J g-1 °C-1), and ∆T is temperature change (°C).
• The specific heat capacity refers to the energy required to raise 1g of a substance by 1K without changing its state.
• Enthalpy change per mole (J mol-1) is calculated using the formula: ΔH = q / moles.
• Calorimetry-derived ∆H values are never completely accurate because energy is easily lost from the system.
• Heat loss can occur through conduction, convection, or temperature measurement inaccuracies.
• Using a lid on the calorimeter and insulating it with a material like polystyrene helps minimize heat loss to the surroundings.
• The specific heat capacity of the solution is taken to be 4.18 kJ mol-1, which is the value for water and not the actual solution.

Hess's Law

• Energy in a reaction system is conserved, thus the overall enthalpy change for a reaction remains the same regardless of the pathway.
• Hess's Law helps determine enthalpy changes for reactions not directly measurable experimentally, using a triangular cycle method with an intermediate product.
• Hess’s Law uses a triangular cycle method where the direction of the arrows indicates whether the values should be added or taken away, being treated like vectors.
• In triangular diagrams for enthalpies of formation, arrows point up from the central product (C) as both A and B are formed from the elements at C.
• In triangular diagrams for enthalpies of combustion, arrows point towards the central product (always H2O and CO2) as both A and B burn to form products at C.

Bond Enthalpies

• Bond enthalpy data is an averaged value representing the energy needed to break one mole of a bond in a gaseous state under standard conditions.
• Different covalent bonds require different amounts of energy to break, values for which can be found experimentally using calorimetry methods.
• Data book values are approximate averages, as bond enthalpy values calculated in this way will vary in each situation.
• Mean bond enthalpy values indicate the energy needed to break a specific bond, averaged across compounds containing that bond.
• Mean bond enthalpy values relate to bond strength and predict which bonds break first in a chemical reaction, with lower values indicating easier breakage.
• Reactions requiring the breaking of many strong bonds or that are endothermic overall are likely to occur slowly at room temperature.
• Mean bond enthalpy values calculate the overall enthalpy change for a reaction using the formula: ΔHreaction = ΣH(bonds broken) - ΣH(bonds formed).
• The enthalpy change for a reaction equals the sum of the bond enthalpies for the bonds broken minus the sum of the bond enthalpies for the bonds formed.
• Bond enthalpy values are often presented in tables to help identify bonds broken and formed when calculating overall enthalpy change.
• For example, the enthalpy change for methane complete combustion in oxygen is calculated using given bond enthalpies, where bonds formed and broken are tallied up to find the net energy change.

Bond Enthalpy Example

• Bond enthalpies in kJ mol-1: C-H is 413, O=O is 498, O-H is 464, and C=O is 799.
• For methane combustion: CH4 + 2O2 → CO2 + 2H2O.
• Total energy to form bonds: C=O x 2 = 1598 kJ/mol, H-O x 4 = 1856 kJ/mol.
• Total energy to break bonds: C-H x 4 = 1652 kJ/mol, O=O x 2 = 996 kJ/mol.
• Enthalpy change = (1652 + 996) - (1856 + 1598) = -806 kJ/mol.