Energy and Chemical Reactions

1. Energy and Chemical Reactions

2. Thermochemistry (Enthalpy) • Remember some reactions are exothermic (produce heat energy) and others are endothermic (absorb heat energy). • Chemists determine exactly how much energy is transferred using a energy function called enthalpy, which is designated by H.

3. When a reaction occurs under conditions of constantpressure, the change in enthalpy (∆H) is equal to the energy that flows as heat. • ∆Hp = heat • the subscript p indicates the process has occurred under constant pressure • ∆ means “change in” • The enthalpy change for a reaction under constant pressure is same as the heat (q) for that reaction.

4. Calorimetry • A calorimeter is a device used to determine the heat associated with a chemical reaction. • The reaction is run in the calorimeter and the temperature change is observed. • Knowing the temperature change in the calorimeter and the heat capacity of the calorimeter the heat energy absorbed by released by the reaction can be calculated. • From this information, the ∆H value for the reaction can be calculated.

5. Hess’s Law • Enthalpy is a state function, meaning the change in enthalpy for a given process is independent of the pathway for the process. • Hess’s law states that in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step of a series of steps.

6. To illustrate Hess’s law the oxidation of nitrogen to produce nitrogen dioxide will be examined. • The overall reaction can be written in one step, where the enthalpy change is represented by ∆H1. • N2 (g) + 2O2 (g) → 2NO2 (g) ∆H1 = 68 kJ • The reaction is carried out in two distinct steps, with the enthalpy changes being designated as ∆H2 and ∆H3. • N2 (g) + O2 (g)→ 2NO (g) ∆H2 = 180 kJ • 2NO (g) + O2 (g) → 2NO2 (g) ∆H3 = -112 kJ • ______________________________________ • Net: N2 (g) + 2O2 (g) → 2NO2 (g) ∆H2 + ∆H3 = 68 kJ

7. Note that the sum of the two steps gives the net, or overall, reaction and that • ∆H1 = ∆H2 + ∆H3 = 68 kJ • The importance of Hess’s law is that it allows us to calculate heats of reaction that might be difficult to measure directly in a calorimeter.

8. To use Hess’s law to compute enthalpy changes for reactions, it is important to understand two characteristics of ∆H for a reaction: • If the reaction is reversed, the sign of ∆H is also reversed. • The sign of ∆H indicates the direction of heat flow: • Exothermic reaction = - ∆H value • Endothermic reaction = + ∆H value • Xe (g) + 2F2 (g) → XeF4 (s) ∆H = -251 kJ • The negative sign of ∆H indicates an exothermic reaction. • If the reaction is reversed • XeF4 (s) → Xe (g) + 2F2 (g) ∆H = +251 kJ • the sign of ∆H is reversed because it becomes endothermic.

9. 2nd characteristic: • The magnitude of ∆H is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ∆H is multiplied by the same integer. • For example, since 251 kJ of energy is evolved the following reaction • Xe (g) + 2F2 (g) → XeF4 (s) ∆H = -251 kJ • then for a preparation involving twice the quantities of reactants and products, or • 2Xe (g) + 4F2 (g) → 2XeF4 (s) ∆H = 2(-251 kJ) = • -502 kJ • twice as much heat is evolved.

10. Standard Enthalpies of Formation • Since the change in enthalpy for some reactions cannot be determined using a calorimeter another method is to use standard enthalpies of formation. • The standard enthalpy of formation (∆Hof) of a compound is defined as the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states. • A degree symbol on a thermodynamics function, for example, ∆Ho , means the process has been carried out at standard conditions (certain pressure and temperature conditions).

11. The enthalpy change for a given reaction can be calculated by subtracting the enthalpies of formation of the reactants from the enthalpies of formation of the products. • Remember to multiply the enthalpies of formation by the coefficients in the balanced equation. • Elements are not included in the calculation because elements require no change in form.