Chemical reactions involve the rearrangement of atoms and molecules, forming new substances. Every reaction involves an energy change as bonds are broken and formed. The enthalpy of reaction, symbolized as ΔH, represents the heat change during a chemical reaction at constant pressure. Understanding this energy change helps predict how much heat a reaction will absorb or release.
Understanding Enthalpy of Reaction
Enthalpy of reaction measures the heat exchanged between a chemical system and its surroundings under constant pressure. A positive enthalpy indicates an endothermic reaction, where the system absorbs heat from its surroundings. Conversely, a negative enthalpy signifies an exothermic reaction, where the system releases heat into its surroundings.
Enthalpy change is typically expressed in kilojoules per mole (kJ/mol). Enthalpy is a state function, meaning the overall change depends only on the initial and final states of the system, not on the specific pathway or steps taken. This allows for various methods of calculating enthalpy changes.
Calculating Enthalpy from Standard Formation Data
One approach to determine reaction enthalpy uses standard enthalpies of formation (ΔH°f). ΔH°f is the enthalpy change when one mole of a compound forms from its constituent elements in their most stable physical state under standard conditions (typically 25°C and 1 atmosphere of pressure). Elements in their standard states, such as O₂(g) or C(s, graphite), have a ΔH°f of zero.
The standard enthalpy of a reaction is calculated using: ΔH°reaction = ΣnΔH°f(products) – ΣmΔH°f(reactants). Here, ‘n’ and ‘m’ are the stoichiometric coefficients from the balanced chemical equation for the products and reactants. This method allows for calculating reaction enthalpies by using tabulated standard formation values.
For example, consider the combustion of methane: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l). To find its enthalpy, use the standard enthalpies of formation for carbon dioxide, liquid water, and methane. Oxygen gas has a ΔH°f of zero. If ΔH°f for CO₂(g) is -393.5 kJ/mol, for H₂O(l) is -285.8 kJ/mol, and for CH₄(g) is -74.8 kJ/mol, the calculation involves summing the heats of formation of the products (1 mol CO₂ + 2 mol H₂O) and subtracting the sum for the reactants (1 mol CH₄ + 2 mol O₂). This provides a precise value for the overall heat released or absorbed during the combustion process under standard conditions.
Calculating Enthalpy Using Hess’s Law
Hess’s Law states that the total enthalpy change for a chemical reaction is independent of the pathway taken, whether the reaction occurs in one step or in a series of steps. This law is useful when it is impractical or impossible to measure the enthalpy change of a reaction directly. It allows chemists to determine the enthalpy of a target reaction by algebraically combining the enthalpy changes of other reactions that sum up to the target reaction.
Applying Hess’s Law involves manipulating known chemical equations and their corresponding enthalpy changes. If a reaction equation is reversed, the sign of its enthalpy change must also be reversed. If a reaction equation is multiplied by a factor, its enthalpy change must be multiplied by the same factor. By strategically adding and subtracting known reactions, intermediate species can cancel out, leaving only the reactants and products of the desired overall reaction.
For example, to find the enthalpy change for the formation of carbon monoxide (C(s) + ½O₂(g) → CO(g)), which is difficult to measure directly, one could use the known enthalpies for the combustion of carbon to carbon dioxide (C(s) + O₂(g) → CO₂(g), ΔH₁ = -393.5 kJ/mol) and the combustion of carbon monoxide to carbon dioxide (CO(g) + ½O₂(g) → CO₂(g), ΔH₂ = -283.0 kJ/mol). By reversing the second reaction (CO₂(g) → CO(g) + ½O₂(g), ΔH₂’ = +283.0 kJ/mol) and adding it to the first reaction, the CO₂ cancels out, yielding the desired reaction and its enthalpy: C(s) + ½O₂(g) → CO(g), ΔH = ΔH₁ + ΔH₂’ = -393.5 kJ/mol + 283.0 kJ/mol = -110.5 kJ/mol. This demonstrates how complex reactions can be broken down into simpler, measurable steps.
Calculating Enthalpy from Bond Energies
Another method for estimating reaction enthalpy uses average bond energies. Chemical reactions involve breaking existing bonds in reactants and forming new bonds in products. Energy is required to break bonds (an endothermic process), while energy is released when new bonds form (an exothermic process). The difference between the energy absorbed and the energy released approximates the reaction’s enthalpy change.
The approximate enthalpy of reaction using bond energies is: ΔH°reaction ≈ Σ(bond energies of bonds broken) – Σ(bond energies of bonds formed). This method provides an estimate rather than an exact value because bond energies are typically average values from many different molecules, and a specific bond’s actual energy can vary depending on its molecular environment.
For example, consider H₂(g) + Cl₂(g) → 2HCl(g). To estimate the enthalpy change, sum the energy to break one H-H bond (approximately 436 kJ/mol) and one Cl-Cl bond (approximately 242 kJ/mol). Then, subtract the energy released by forming two H-Cl bonds (each approximately 431 kJ/mol). This calculation provides an insight into the energy balance of bond transformations during the reaction, indicating whether the reaction is likely to release or absorb heat.
Determining Enthalpy Experimentally
Reaction enthalpy can also be determined experimentally through calorimetry. Calorimetry measures the heat absorbed or released by a reaction by observing the temperature change of a known mass of a substance, usually water or a solution, within a device called a calorimeter. The basic principle relies on the idea that any heat released by the reaction is absorbed by the calorimeter and its contents, and vice-versa.
The quantity of heat (q) absorbed or released by the surroundings is calculated using q = mcΔT, where ‘m’ is the mass of the substance, ‘c’ is its specific heat capacity (the amount of heat required to raise the temperature of one gram of the substance by one degree Celsius), and ‘ΔT’ is the measured change in temperature. Once ‘q’ is determined, it can be related to the enthalpy change (ΔH) of the reaction. For reactions conducted at constant pressure, the heat absorbed or released by the solution in the calorimeter (q_solution) is equal in magnitude but opposite in sign to the enthalpy change of the reaction (q_reaction = -q_solution). This experimental approach provides direct, real-world measurements of energy changes in chemical processes.