Energetics

1. Energetics Exothermic and endothermic reactions Topic 5.1

2. Energy and Heat • Energy is the capacity to do work or to transfer heat. • Work (J) = force (N) x distance (m) • Energy is measured in Joules. You do 1 Joule of work when you exert a force of 1 N over a distance of 1 meter. • 1 Calorie = 4.184 Joules • Energy comes in many forms and includes heat, light, sound, electricity and chemical energy. • The most important type of energy in chemistry is heat.

3. Temperature and Heat • Heat, represented by q, is the energy that transfers from one object to another as a result of a temperature difference. • Temperature is a measure of the average kinetic energy of the particles in a substance. It is independent of the amount of substance present. • Ek = ½ mv2 • Adding heat to a substance causes its kinetic energy to increase and, consequently, causes an increase in the temperature of the substance. So heat is energy used to cause the temperature of an object to increase.

4. System and Surrounding • Chemical and physical changes take place in many different environments. It is important to distinguish between system and surroundings. • The system – the part of the universe we wish to focus our attention on. The system consists of those molecules which are reacting. • The surroundings - everything else; the rest of the universe • The system + surrounding = universe

5. Exothermic and Endothermic Reactions • Exothermic reactions – reactions that result in a transfer of heat from the system to the surroundings. • Most chemical reactions, including all combustion and neutralization reactions are exothermic. • Since heat is given out during the reaction- the products have less energy or heat content than the reactants.

6. Exothermic and Endothermic Reactions • Endothermic reactions – reactions in which energy is absorbed from the surroundings. • In endothermic reactions, the products have more heat content than the reactants.

7. Endothermic and Exothermic Reactions • Endothermic and Exothermic Reactions

8. Exothermic and Endothermic Reactions • The heat content of a substance is called its enthalpy, H. • The absolute value of the enthalpy of the reactants cannot be measured, nor can the enthalpy of the products, but what can be measured is the difference between them, Δ H. This quantity, H, is called the enthalpy of reaction, or the heat of reaction. • Δ H = Hproducts – Hreactants • Δ H is negative for exothermic reactions because the enthalpy of the products is less than the reactants. • Δ H is positive for endothermic reactions because the enthalpy of the products is greater than the reactants.

9. Exothermic and Endothermic Reactions When heat is released by the system to the surroundings, the process is exothermic When heat is absorbed by the system from the surroundings, the process is endothermic

10. Standard Enthalpy Change of a Reaction • The standard enthalpy change of a reaction is the enthalpy change that occurs in a system when one mole of matter is transformed by a chemical reaction under standard conditions. • The standard enthalpy change of a reaction is denoted by Δ Ho • The standard conditions for enthalpy changes are: • a temperature of 298 K • A pressure of 1 atm • Concentrations of 1 mol dm-3 for all solutions • All the substances are in their standard states

11. Thermochemical Equations • A thermochemical equation shows the reactants and products and the heat that is released or absorbed by a reaction. • For example, the combustion of methane can be described by a thermochemical equation: • CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) Δ H = - 890 kJ mol-1 1 mole of methane gas reacts with two moles of oxygen gas to give one mole of gaseous carbon dioxide and two moles of liquid water, and releases 890 kJ of heat energy • The thermochemical equation for photosynthesis can be represented as: • 6CO2(g) + 6H2O(l)  C6H12O6(aq) + 6O2(g) Δ H +2802.5 kJ mol-1

12. Thermochemical Equations • Keep in mind the following when using thermochemical equations: • 1. the coefficients in a balanced thermochemical equation represent the number of moles of reactants and products and never the numbers of molecules. So, the coefficients can be written as fractions when necessary. • 2. the physical state of the products and reactants involved must be specified. • 3. the change in enthalpy represented in a thermochemical equation is directly proportional to the number of moles of substances undergoing a change. If the moles are doubled, then the enthalpy change must also be doubled. • 4. the value of the enthalpy change is usually not significantly influenced by changing temperature. • 5. When a chemical equation is reversed, the magnitude of Δ H remains the same, but the sign changes.

13. Calorimetry • Calorimetry is the science of measuring the heat changes of chemical reactions or physical changes. Calorimetry is performed with a calorimeter. • So the enthalpy change for a reaction can be measured experimentally by using a calorimeter. • In a simple calorimeter, the heat evolved in an exothermic reaction is used to raise the temperature of a known mass of water. • For endothermic reactions, the heat transferred from the water to the reaction can be calculated by measuring the lowering of temperature of a known mass of water.

14. Calorimetry • In general, the increase in temperature when an object is heated depends on : • The mass of the object • The heat added • The nature of a substance • Different substances need different amounts of heat to increase their temperature; they have different specific heats – defined as the heat needed to raise the temperature of 1 gram of a substance by 1 degree Celsius. • So, heat change (q) = mass (m) x specific heat (s) x change in temperature (Δ T)

15. Enthalpy Changes • In chemical reactions, the chemicals are changing in a way that reduces their enthalpy, in this way they become more stable. So, we generally expect a reaction to occur if it leads to a reduction in enthalpy. However, this is not completely reliable since there is another factor that affects whether or not a reaction occurs spontaneously, the disorder of the system (called entropy) • So, some endothermic reactions, even though the products are less stable than the reactants, occur spontaneously because there is an increase in the disorder of the system. • For example: 6SOCl2(l) + FeCl3·6H2O(s)FeCl3(s) +6SO2(g) + 12 HCl(g) Δ H = +1271 kJ mol-1 • Even though this reaction is extremely endothermic, it occurs spontaneously because the production of gas increases the disorder of the system.