1. 3.3 – Enthalpy and Heat of Reaction

2. Basic Info: All chemical reactions involve energy changes, whether energy is being absorbed or given off. Where does this energy come from? All this energy is stored in chemical bonds. The bonds within a molecule holding the atoms together are called intramolecular bonds. The bonds holding molecules to each other are called intermolecular bonds. When bonds are broken, energy is absorbed. This would be an endothermic process as energy needs to be put in to break the bonds apart When bonds are formed, energy is released. This would be an exothermic process as energy is given off when the compound becomes more stable after a bond forms

3. Exothermic Reactions: Reactions in which there is a net release of energy. When energy is released, an energy term will appear on the product side of the equation. i.e.: The thermite reaction is a highly exothermic reaction between iron (III) oxide and aluminum, producing aluminum oxide and molten iron: Fe 2 O 3 (s) + 2Al (s) → Al 2 O 3(s) + 2Fe (l) + 847.6 kJ If you had this reaction occurring in your hand (don't try this at home!!!) your hand would feel warmer because the reaction is releasing heat.

4. Endothermic Reactions: Reactions that require a net input of energy. This is indicated by writing the energy term on the reactant side of the equation. i.e.: Consider the reaction when sulfur trioxide decomposes into sulfur trioxide and oxygen. This reaction requires a net input of 198 kJ of energy in order to occur: 2 SO 3 (g) + 198 kJ → 2 SO 2 (g) + O 2 (g) If you held this reaction in your hand, your hand would feel colder because the energy required for the reaction to occur would be taken from the surroundings, in this case your hand.

5. Enthalpy Thermodynamics is the study of heat. Specifically, we are studying the energy changes during chemical reactions. Enthalpy (H˚) is the heat content in a system, or, the total amount of potential and kinetic energy within a substance. This is energy, so it is measured in joules (J). Note: The ˚ symbol indicates a temperature of 25˚C and 101 kPa, which is standard temperature and pressure in thermodynamics (STP).

6. Heat of Reaction The enthalpy of a system then comes from both the bonds holding the substance together, and the motion of the particles within the system. Although we cannot measure the enthalpy of a system, we can calculate the enthalpy changes during a reaction (ΔH˚). We also call this change in enthalpy the heat of reaction.

7. Enthalpy and Exothermic Reactions Recall that an exothermic reaction releases more energy to form the bonds in the products than it did to break the bonds in the reactants. Ex) Cu(s) + Cl2(g)  CuCl2(g) + 220.1 kJ Exothermic reactions release heat because the reactants have a higher heat content (enthalpy). The heat released is the ‘excess’ heat. Since CuCl2 has a lower heat content, the extra heat is released to the surroundings.

8. Enthalpy and Exothermic Reactions Graphically, we can represent the potential energy for the reactants and products like:

9. Enthalpy and Exothermic Reactions The graph shows that the amount of potential energy in the reactants is more than the potential energy in the products. This is like rolling a boulder down a hill. As the boulder goes down, it loses potential energy. The drop in potential energy is the heat released, or the heat of reaction (ΔH). The value is negative since heat is being released.

10. Enthalpy and Exothermic Reactions We can then remove the energy from the equation and record it as a change in enthalpy: Cu(s) + Cl2(g)  CuCl2(g) ΔH˚ = -220.1 kJ So, the heat of reaction above is -220.1 kJ. We describe the reaction by saying that in the formation of 1 mole of CuCl2 from 1 mole of Cu and 1 mole of Cl2, 220.1 kJ was released.

11. Enthalpy and Endothermic Reactions Likewise, an endothermic reaction requires the addition of energy. This ‘excess’ energy is stored as potential energy in the bonds of the products. Ex) H2O(g) + C(s) + 132 kJ  CO2(g) + 2H2(g) Here, the potential energy is lower in the reactants. Therefore, the enthalpy is increased during the reaction.

12. Enthalpy and Endothermic Reactions

13. This is like pushing a boulder up a hill… it takes energy. We can now remove the energy from the equation and record it as a change in enthalpy. H2O(g) + C(s)  CO2(g) + 2H2(g) ΔH˚ = 132 kJ

14. Heat of Formation Standard heat of formation (ΔH˚f) is the amount of heat required or given off when 1 mole of a compound is formed from its component elements. The key here is 1 mole! For example, the formation of 1 mole of water: H2(g) + ½ O2(g)  H2O(l) + 285.8 kJ Therefore, ΔH˚f = -285.8 kJ for the formation of water.

15. Heat of Formation Remember that this value is only for the formation of 1 mole from its elements. This is only a special heat of reaction (when 1 mole of a substance is formed). Note the heat of formation of pure substances, like H2, Cl2, Al, etc. is 0 kJ/mole. When writing heat of formation reactions, remember the following: 1. The equation must be balanced so 1 mole of the product is produced. 2. The diatomic elements are: H, O, F, Br, I, N, and Cl. 3. The reactants must be elements, not polyatomic ions.

16. Examples: Write the heat of formation reactions for the following compounds: a) MgO -601.7 kJ/mol b) CaCO3 -1206.9 kJ/mol

17. Examples: If we look back at our equation for water, how much energy is released if two moles are created? H2(g) + ½ O2(g)  H2O(l) + 285.8 kJ ΔH˚f = -285.8 kJ/mol We can find the enthalpy of any number of moles of product formed by multiplying the enthalpy by the number of moles. For 2 moles: -285.8 kJ/mol x 2 mol = -571.6 kJ

18. Examples: How much heat is released when 25.0 g of water are produced?