Energy and Chemical Reactions

1. Energy and Chemical Reactions Energy is transferred during chemical and physical changes, most commonly in the form of heat

2. Energy • Energy can be kinetic – associated with motion, such as thermal, mechanical, electric, sound • Energy can be potential – associated with an object’s position, such as chemical, gravitational, electrostatic • Energy is converted from one form to another

3. First Law of Thermodynamics • The total energy of the universe is constant • Energy is conserved

4. Temperature and Heat • Temperature is a measure of the average kinetic energies of the particles in a substance • Heat is energy that can be transferred between substances that are at different temperatures • Heat will transfer between two objects in contact until thermal equilibrium occurs

5. Heat transfer • The quantity of heat lost by a hotter object and the quantity of heat gained by a cooler object when they are in contact are numerically equal (but opposite direction) • Exothermic – heat is transferred from the system to the surroundings • Endothermic – heat is transferred from the surroundings to the system

6. Energy Units • Joule is the SI unit for thermal energy • 1 J = 1 kg.m2/s2 • Kilojoules are also commonly used • The calorie is an older unit for heat; 1 cal = 4.184 J • Dietary Calories are actually 1000 calories

7. Specific Heat Capacity and Heat Transfer • The quantity of heat transferred to or from an object when its temperature changes depends on: • Quantity of the material • Size of the temperature change • Identity of the material • Specific heat capacity – the quantity of heat required to raise the temperature of 1.00g of a substance by one kelvin (J/g.K)

8. q = m c DT • q is heat in joules • m is mass in grams • DT = Tfinal – Tinitial • Water has a particularly high specific heat; metals have low specific heats

9. Assumptions • Heat transfers until both substances are at the same temperature • We assume no heat is transferred to warm the surroundings (though this is not accurate) • The heat that is lost by one substance is equal and opposite in sign to the heat that is gained by the other substance

10. Energy and Changes of State • Heat of fusion – energy to convert a substance from solid to liquid (J/g) • Heat of vaporization – energy to convert a substance from liquid to gas (J/g) • The energy required for a change of state is determined by the type of substance and its quantity (mass)

12. Calorimetry • Constant pressure calorimetry measures DH • Constant pressure calorimetry can be done with a coffee-cup calorimeter • A reaction changes the temperature of the solution in the calorimeter; measuring the change in the solution allows calculation of the change in the reaction • qrxn + qsolution = 0

13. Calorimetry • Constant volume calorimetry measures DE • A bomb calorimeter is used for constant volume calorimetry • qrxn +qbomb +qwater = 0

14. Thermodynamics – the study of heat and work • DE = q + w • DE is the change in kinetic and potential energies of the system • Positive q is heat going into the system • Negative q is heat leaving the system • Positive w is work done on the system • Negative w is work done by the system • Work (of a gas): w = - P(DV)

15. State Functions • A quantity that is the same no matter what path is chosen in going from initial to final • Changes in internal energy and enthalpy for chemical or physical changes are state functions • Neither heat nor work individually are state functions, but their sum is

16. Enthalpy Changes for Chemical Reactions • Measures the change in heat content • Enthalpy changes are specific to the identity and states of reactants and products and their amounts • DH is negative for exothermic reactions and positive for endothermic reactions • Values of DH are numerically equal but opposite in sign for chemical reactions that are the reverse of each other • Enthalpy change depends on molar amounts of reactants and products

17. 2 Methods to find DHrxnHess’s Law (indirect method) • If a reaction is the sum of two or more other reactions, DH for the overall process is the sum of the DH values of those reactions

18. Standard Enthalpies of Formation • The standard molar enthalpy of formation (DHfo) is the enthalpy change for the formations of 1 mol of a compound directly from its component elements in their standard states • The standard state of an element or a compound is the most stable form of the substance in the physical state that exists at standard atmosphere at a specified temperature

19. Standard Enthalpy of Formation • The standard enthalpy of formation for an element in its standard state is zero • Most enthalpies of formation values are negative, indicating an exothermic process • The most stable compounds have the largest exothermic values

20. Enthalpies of Formation • Enthalpy change for a reaction can be calculated from the enthalpies of formation of the products and reactants (direct method): • S [DHfo(products)] – S [DHfo(reactants)] = DHrxno • Reactions with negative values of DHrxno are generally product-favored, while positive DHrxno usually indicates a reactant-favored reaction