Energy and Chemical Reactions Chapter 5 Learning Objectives

Energy and Chemical Reactions Chapter 5

Learning Objectives • Students understand • And apply the first law of thermodynamics

Learning Objectives • Students will be able to • Assess the transfer of energy as heat • Calculate P-V work done by a system and use energy as heat and work to calculate changes in internal energy • Calculate energy and enthalpy changes from calorimetry data • Calculate energy evolved or required using thermodynamic data

5. 1 Energy • the science of heat and work is called thermodynamics • The kinetic energy of molecules is referred to as thermal energy.

Energy • The Law of Conservation of Energy states that energy can be converted from one form to another, but can neither be created nor destroyed • restated mathematically as the first law of thermodynamics

Systems and Surroundings • A system is an object, or collection of objects, being studied. • The surroundings include everything outside the system that can exchange energy and/or matter with the system.

Temperature and Heat • heat is not the same as temperature • more thermal energy = more particle motion • total thermal energy is sum of individual energies of all particles

Thermal Equilibrium • thermal equilibrium means that two objects have reached the same temperature • heat always transfers from hotter object to cooler object • transfer continues until equilibrium is reached • heat lost = heat gained

Thermal Equilibrium • exothermic process occurs when heat transfers from system to surroundings • heat is released as a product • qsys < 0 • endothermic process occurs when heat transfers from surroundings into system • heat is required as a reactant • qsys > 0

5. 2 Specific Heat • quantity of energy transferred depends on • Quantity of material • Magnitude of the temperature change • Identity, including phase, of the material gaining or losing energy

5. 2 Specific Heat • specific heat is the quantity of heat required to raise 1 gram of substance 1 kelvin (J/g∙K) • q = CmΔT • Where ΔT = Tfinal - Tinitial

Practice Problem • In an experiment it was determined that 59. 8 J was required to change the temperature of 25. 0 g of ethylene glycol by 1. 00 K. Calculate the specific heat capacity of ethylene glycol.

Quantitative Aspects • In an isolated system, the sum of the energy changes must be zero! • Energy is CONSERVED!!

Practice Problem • A 15. 5 g piece of chromium, heated to 100. 0 o. C, is dropped into 55. 5 g of water at 16. 5 o. C. The final temperature of the metal and the water is 18. 9 o. C. What is the specific heat of chromium? Assume no heat is lost to container or air.

Homework • After reading sections 5. 1 and 5. 2, you should be able to do the following problems… • p. 217 (3 -4, 9 -16)

5. 3 Energy and Changes of State • Heat of fusion: solid to liquid (endothermic) • Heat of vaporization: liquid to gas (endothermic) • Heat of solidification: liquid to solid (exothermic) • Heat of condensation: gas to liquid (exothermic)

Energy and Changes of State • Temperature is constant during a state change!!

Practice Problem • How much heat must be absorbed to warm 25. 0 g of liquid methanol, CH 3 OH, from 25. 0 o. C to 64. 6 o. C and then to evaporate the methanol completely at that temperature? The specific heat of methanol is 2. 53 J/g∙K. ΔHvap = 2. 00 x 103 J/g

Practice Problem • To make a glass of iced tea, you pour 250 m. L of tea, whose temperature is 18. 2 o. C, into a glass containing 5 ice cubes. Each cube has a mass of about 15 g. What quantity of ice will melt, and how much ice will remain to float at the surface in this beverage? Iced tea has a density of 1. 0 g/cm 3 and a specific heat of 4. 2 J/g∙K. Assume no energy transferred between system and surroundings.

5. 4 First Law of Thermodynamics • the energy change for a system is the sum of the energy transferred as heat and the energy transferred as work • ΔU = q + w • Where • ΔU is change in internal energy • q is heat transfer to or from system • w is work transfer to or from system, called P-V (pressure-volume) work

P-V work • work done on or by the system equals the volume change that occurs against resisting external pressure • w = - P(ΔV) • If energy is transferred at a constant volume process, energy transferred as work is zero and the ΔU = q

Enthalpy • Enthalpy is heat content of a substance at constant pressure • Enthalpy change (ΔH) • negative ΔH and ΔU mean that energy is transferred from system to surroundings • positive signs mean that energy is transferred from surroundings to system

State Functions • changes in ΔH or ΔU depend only on final and initial states • the pathway to go between is not relevent (example: bank accounts)

Homework • After reading sections 5. 3 and 5. 4, you should be able to do the following problems… • p. 217 a (21 -30)

5. 5 Enthalpy Changes • enthalpy changes are specific to the reactants and products and their amounts (states of matter are important) • ΔH depends on the number of moles of reaction as written • ΔH is positive when endothermic and negative when exothermic (reverse reactions have opposite sign)

Enthalpy Changes • Standard reaction enthalpy: enthalpy change accompanying a specific reaction • Standard state: most stable form of the substance in the physical state at a given pressure and temperature

Enthalpy Changes • change depends on molar amounts of substances • calculate moles of substance and then multiply that by heat transfer per mole

Practice Problem • What quantity of heat energy is required to decompose 12. 6 g of liquid water to the elements? • The combustion of ethane, C 2 H 6, has an enthalpy change of -2857. 3 k. J for the reaction. Calculate ΔH when 15. 0 g is burned. 2 C 2 H 6(g) + 7 O 2(g) 4 CO 2(g) + 6 H 2 O(g)

5. 6 Calorimetry • technique to determine heat transfer • constant pressure measures change in enthalpy • constant volume measures change in internal energy

Constant Pressure Calorimetry • can be used to determine heat gained or lost by solution • can be used to determine heat required or released by reaction • at constant pressure, the heat measured is ΔH • change in heat content of solution can be measure and used to calculate the heat of reaction • qrxn + qsoln = 0

Constant-Pressure Calorimetry

Practice Problem • Assume you mix 200. m. L of 0. 400 M HCl with 200. m. L of 0. 400 M Na. OH in a coffee-cup calorimeter. The temperature of the solutions before mixing was 25. 20 o. C; after mixing and allowing reaction to occur, the temperature is 27. 78 o. C. What is the molar enthalpy of neutralization of the acid? (assume densities of all solutions are 1. 00 g/m. L and their specific heats are 4. 20 J/g. K)

Constant Volume Calorimetry • used to calculate heats of combustion and caloric value of foods – use a “bomb” • constant volume, so energy transfer as work doesn’t occur and heat is therefore the change in internal energy • qr + qbomb + qwater = 0

Constant Volume Calorimetry

Practice Problem • A 1. 00 g sample of sucrose is burned in a bomb calorimeter. The temperature of 1. 50 x 103 g of water in the calorimeter rises from 25. 00 o. C to 27. 32 o. C. The heat capacity of the bomb is 837 J/K and the specific heat of water is 4. 20 J/g. K. Calculate (a) the heat evolved per gram of sucrose and (b) the heat evolved per mole of sucrose.

Homework • After reading sections 5. 5 and 5. 6, you should be able to do the following… • p. 217 a-b (33 -42)

5. 7 Hess’s Law • sometimes products immediately undergo other reactions and therefore calorimetry cannot be used • Hess’s Law states that if a reaction is the sum of two or more other reactions, ΔH for the overall process is the sum of the ΔH values of those reactions

Energy Level Diagrams • can visualize Hess’s Law by using diagrams to show the formation of different steps in a reaction as well as the enthalpy changes involved

Energy Level Diagrams

Practice Problem • What is the enthalpy change for the formation of ethane, C 2 H 6, from elemental carbon and hydrogen? 2 C(s) + 3 H 2(g) C 2 H 6(g)

Practice Problem • Use Hess’s Law to calculate enthalpy change for the formation of CS 2(l) from C(s) and S(s) from the following enthalpy values. C(s) + O 2(g) CO 2(g) ΔH = -393. 5 k. J S(s) + O 2(g) SO 2(g) ΔH = -296. 8 k. J CS 2(g) + 3 O 2(g) CO 2(g) + 2 SO 2(g) ΔH = -1103. 9 k. J C(s) + 2 S(s) CS 2(g) ΔH = ?

Standard Enthalpies • Standard molar enthalpy of formation, ΔHof, is the heat change for the formation of 1 mol of a compound directly from its component elements in their standard states. • See appendix L

Standard Enthalpies • Standard enthalpy of formation for an element in its standard state is zero. • Most values are negative • can compare thermal stability; more exothermic is more stable

Enthalpy Change for a Reaction • ΔHorxn = Σ[ΔHof(products)] - Σ[ΔHof(reactants)] • Find enthalpies in a table and then plug the values into the equation above

Practice Problem • Calculate the standard enthalpy of combustion for benzene, C 6 H 6. • C 6 H 6(l) + 7½O 2(g) 6 CO 2(g) + 3 H 2 O(l) • ΔHof[C 6 H 6(l)] = +48. 95 k. J/mol

Practice Problem • Case Study p. 212: The Fuel Controversy – Alcohol and Gasoline

5. 8 Favored Reactions • product-favored reactions go from left right • most reactions that are exothermic (have negative values of enthalpy) are product-favored • most reactions that are endothermic (have positive enthalpy change) are reactant-favored

Practice Problem • Calculate ΔHorxn for each of the following reactions and decide if the reaction may be product- or reactant-favored. • 2 HBr(g) H 2(g) + Br 2(g) • C(diamond) C(graphite)

Homework • After reading sections 5. 7 and 5. 8, you should be able to do the following problems… • p. 217 c-d (51 -52, 63 -64, 115)