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 GOAL: To be able to define energy and to recognize different types of energy

3. First Law of Thermodynamics • The total energy of the universe is constant • Energy is conserved GOAL: To understand that energy is not created or destroy, but transferred between different places and between different types of energy

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 GOAL: To understand the different between heat and temperature, and to use the two terms correctly

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 GOAL: To be able to recognize exothermic and endothermic in a variety of situations

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 GOAL: To use the unit Joules to solve energy problems

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) GOAL: To know the factors that determine temperature change when heat is applied or removed from an object

8. Consequences of Specific Heat Capacity • Objects with a large c value take more energy to change temperature (compare seat belt buckle with bottle of water, or sand with water at the beach) • If objects with different c values are allowed to reach thermal equilibrium, the object with the large c value contains more heat energy (aluminum foil vs baked potato) GOAL: To understand how specific heat capacity affects temperature change as heat is applied or removed from an object

9. q = m c DT • Use to find Heat when Temperature is Changing! • 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 GOAL: To be able to calculate heat lost or gained when a temperature change occurs

10. 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 GOAL: To understand these assumptions and apply them to solve Calorimetry problems (begins on slide 14)

12. 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) GOAL: To be able to calculate the energy lost or gained in a state change

13. For a State Change: q = m (Hfus) or q = m (Hvap) EXAMPLE: The heat required to turn 10 degree water into 120 degree steam = (heat to raise temp of water to 100 degrees) + (heat to change state) + (heat to raise temp of steam to 120 degrees)

14. Observe: Solid-Liquid Equilibrium • Discussion • Use these terms: heat/energy, temperature, kinetic energy, potential energy, states, solid, liquid, state change, equilibrium • White boards: Draw particle view of a solid Draw particle view of a liquid Draw particle view of the melting/freezing phase change

15. GOAL: To be able to label a heating/cooling curve with: states (solid, liquid, gas), state/phase changes (boiling, evaporating, condensing, solidifying), types of energy changes (kinetic or potential), how heat is calculated at each segment (heat of fusion or vaporization, q=mcDT)

17. 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 GOAL: To understand that heat lost by one substance equals heat gained by another substance within a closed system; and to be able to use this concept to solve constant pressure calorimetry problems

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

19. 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) GOAL: To be able to assign signs to heat and work and solve for DE

20. 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 GOAL: To be able to define state function and recognize that DE and DH are state functions while q and w are not

21. 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

22. 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 GOAL: To be able to solve for DHrxn using both the indirect and direct methods

23. C(s) + O2(g)  CO2(g) DH=-393.5 kJ/mol CO(g) + 1/2O2(g) CO2(g) DH=-283.0 kJ/mol _________________________________ _________________________________ C(s) + 1/2O2(g)  CO(g) DH=?

24. C(s) + O2(g)  CO2(g)DH=-393.5 kJ/mol H2(g) + 1/2O2(g)  H2O(l)DH=-285.8 kJ/mol 2C2H2(g) + 5O2(g)  4CO2(g) + 2H2O(l) DH=-2598.8 kJ/mol __________________________________ __________________________________ 2C(s) + H2(g)  C2H2(g)DH=?

25. 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

26. 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

27. 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

28. 2H2S(g) + 3O2(g)  2H2O(l) + 2SO2(g) DH=? SubstanceDHf(kJ/mol) H2S (g) -20.15 H2O (l) -285.8 SO2 (g) -296.1

29. Application of Enthalpy 2 Al(s) + Fe2O3(s)  Al2O3(s) + 2 Fe(l) DHrxn= -822.8 kJ • How much heat is released if 1 mole of Al is used? • How much heat is released if 4.2 moles of Al is used? • How much heat is released if 150. g of Al is used?