Chapter 6 THERMOCHEMISTRY

1. Chapter 6THERMOCHEMISTRY • Energy: The capacity to do work or to produce heat • Law of Conservation of Energy: Energy can be converted from one form to another but can be neither created nor destroyed. The energy of the universe is constant (Euniv = constant)

2. The Two Types of Energy: Potential Energy: Energy due to position or composition - can be converted to work Kinetic Energy: Energy due to the motion of the object and depends on the mass of the object m and its velocity , K.E =1/2m2

3. Work: Defined as force acting over a distance. • Temperature V. Heat: Temperature reflects random motions of particles, therefore related to kinetic energy of the system. Heat involves a transfer of energy between two objects due to temperature difference. • State Function (Property): A property of the system that depends only on its present state. A state function does not depend on the system’s past or future. It does not depend on how the system arrived at the present state (independent of pathway).

4. System: The part of the universe on which we wish to focus attention. • Surroundings: Include everything else in the universe Universe = System + Surroundings • Exothermic Reaction: A reaction results in the evolution of heat. Energy flows out of the system. A reaction gives off heat, the surrounding become warmer. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + energy (heat)

5. The Combustion of Methane

6. Endothermic Reaction: Reactions that absorb energy (heat) from the surroundings. Heat flow is into a system, the surroundings become cooler. N2(g) + O2(g) + energy (heat)  2NO(g) • For exothermic process – energy gained by the surroundings must be equal to the energy lost by the system. For endothermic process the situation is reversed. • Thermodynamics: The study of energy and its interconversions.

7. The Energy Diagram for the Reaction of Nitrogen and Oxygen to Form Nitric Oxide

8. First Law of thermodynamics: The energy of the universe is constant. • Internal Energy (E): The internal energy (E) of a system can be defined as the sum of the kinetic and potential energies of all the particles in the system. The internal energy of a system can be changed by a flow of work, heat or both. E = q + w E = change in system’s internal energy q = heat w = work

9. Sign of heat flow: ‘q’ is positive when heat flows into the system from the surroundings. When q is positive, the process is called endothermic. ‘q’ is negative when heat flows out of the system to the surroundings. When q is negative the process is called exothermic.

10. Pressure is defined as force per unit area, P = F/A  F = P x A Work = force x distance Work = F x h (h = distance) = P x A x h Change in volume, V = final volume – initial volume = A x h work = P x A x h = P x V For an expanding gas, V is a positive quantity because the volume is increasing. V and w have opposite signs, which leads the equation, (work flows into the surrounding, w is negative) W = -PV Work = pressure x volume

11. The Volume of a Cylinder

12. Enthalpy and Calorimetry • Enthalpy (H): Defined as H = E + PV where, E = Internal energy, P = Pressure, V = Volume At constant pressure, E = qp + w E = qp - PV qp= E + PV where, qp = heat at constant pressure H = E + (PV) = E + PV Therefore, H = qp

13. At constant pressure, the change in enthalpy H of the system is equal to the energy flow as heat, i.e, flow of heat is a measure of the change in enthalpy. Enthalpy change, H = Hproducts – Hreactants exothermic reactions: H < 0 (negative) endothermic reactions: H > 0 (positive)

14. Rules of Thermochemistry • The magnitude of H (heat flow) is directly proportional to the amount of reactant or product (extensive property). • H for a reaction is equal in magnitude but opposite in sign to H for the reverse reaction. • If a reaction is multiplied or divided by a number the H of the reaction is also multiplied or divided by that number. • The values of H for a reaction is the same whether it occurs directly or in a series of steps. • The H of a reaction depends on the physical state of the reactants and products.

15. Example: When 1 mole of methane (CH4) is burned at constant pressure, 890 kJ of energy is released as heat. Calculate H for a process in which a 5.8 g sample of methane is burned at constant pressure. CH4(g) + 2O2(g) CO2(g) + 2H2O(l)H = -890 kJ • If we have 1.00 g of O2 then H = ?

16. Calorimetry • Calorimeter: The device used experimentally to determine the heat associated with a chemical reaction. • Calorimetry: The science of measuring heat, is based on observing temperature change when a body absorbs or discharges energy as heat. • Heat Capacity: The amount of energy required to raise the temperature of a system by1oC (J/oC) C = [heat absorbed(q)] / [increase in temperature(T)] q = c x T

17. A Coffee-Cup Calorimeter Made of Two Styrofoam Cups

18. A Bomb Calorimeter

19. Specific Heat Capacity: The energy required to raise the temperature of one gram of a substance by one degree celsius (J/oC.g). • Molar Heat Capacity: The energy required to raise the temperature of one mole of a substance by one degree celsius (J/oC.mole). • Energy released by the reaction = energy absorbed by the solution = specific heat capacity x mass of solution x increase in temperature = s x m x T

20. Example: If 100 J of energy is added to 50 g of Cu, initially at 25.0oC, what will be the final temperature? (Specific heat capacity = 0.382 J/oC.g )

21. Example: How much energy does it take to heat 120 g of water from 20oC to 45oC? q = m x s x T = 120 x 4.18 J / oC .g x (45oC - 20oC) = 12540 J = 12.540 kJ

22. Example: A 55.0 g piece of metal was heated to 99.8oC and dropped into a calorimeter which contains 225 mL of water (density = 1.00g / mL) and is at 21.0oC. The final temperature of the water and the metal is 23.1oC. What is the specific heat of the metal? Energy released by the metal (qmetal) = Energy absorbed by water (qwater) Qwater = m x s x T = 225 g x 4.18 J/oC.g x (23.1oC -21.0oC) = 1975 J Qmetal = -1975 J = m x s x T = 55.0 g x s x (23.1oC –99.8oC) s = -1975 J / -4218.5 goC = 0.468 J/oC.g

23. Hess’s Law • In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. Example: Reactants Products N2(g) + 2O2(g) 2NO2(g)H1 = 68 kJ

24. …Hess’s Law continued… This reaction also can be carried out in two distinct steps, N2(g) + O2(g) 2NO(g)H2 = 180 kJ 2NO(g) + O2(g) 2NO2(g)H3 = -112 kJ Net reaction: N2(g) + 2O2(g) 2NO2(g)H2+ H3 = 68 kJ H1 = H2+ H3 = 68 kJ

25. The Principle of Hess’s Law

26. Characteristics of Enthalpy Changes • If a reaction is reversed, the sign of H is also reversed Xe(g) + 2F2(g) XeF4(s) H = -251 kJ XeF4(s) Xe(g) + 2F2(g) H = +251 kJ

27. …Characteristics of Enthalpy Changes continued… • The magnitude of H is directly proportional to the quantities of reactants and products in a reaction. If the co-efficients in a balanced reaction are multiplied by an integer, the value of H is multiplied by the same integer. H is an extensive property, depends on the amount of substance reacting. Xe(g) + 2F2(g) XeF4(s)H = -251 kJ 2Xe(g) + 4F2(g) 2XeF4(s)H = 2(-251 kJ) = -502 kJ

28. Standard Enthalpies of Formation • The standard enthalpy of formation (Hof) of a compound is defined as the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states.

29. Standard States For a compound: • The standard state of a gaseous substance is a pressure of exactly 1 atmosphere. • For a pure substance in a condensed state (liquid or solid), the standard state is the pure liquid or solid. • For a substance present in a solution, the standard state is a concentration of exactly 1M. For an element: • The standard state of an element is the form in which the element exists under conditions of 1 atmosphere and 25oC.

30. Pathway for the Combustion of Methane

31. A Schematic Diagram of the Energy Changes for the Reaction CH4(g) + 2O2(g) ® CO2(g) + 2H2O(l)

32. Change in Enthalpy • The enthalpy change for a given reaction can be calculated by subtracting the enthalpies of formation of the reactants from the enthalpies of formation of the products. Horeaction = np Hof (products) -  nr Hof (reactants) • Example: 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(l) 4(-46) 5(0) 4(+90) 6(-242) Hrxn = 360 + (-1452) – (-184) = -1092 + 184 = -908 kJ

33. Present Sources of Energy Energy Sources Used in the United States

34. Petroleum is a thick, dark liquid composed mostly of compounds called hydrocarbons that contain carbon and hydrogen. Example, Gasoline (C5 – C10), Kerosene and Jet fuel (C10 – C18), Diesel fuel, Heating oil, Lubricating oil (C15 – C25), Asphalt (>C25 ). • Natural gas consists mostly of methane (CH4), but it also contains ethane (C2H6), propane (C3H8) and butane (C4H10). • Coal was formed from the remains of plants that were buried and subject to high pressure and heat over long periods of time.

35. Effects of Carbon dioxide on Climate The earth receives a tremendous quantity of radiant energy from the sun. Some of this energy is absorbed by plants for photosynthesis and some by the oceans to evaporate water but most of it is absorbed by soil, rocks, and water, increasing the temperature of the earth’s surface. Molecules in the atmosphere, principally H2O and CO2, strongly absorb infrared radiation and radiate it back toward the earth so a net amount of thermal energy is retained by the earth’s atmosphere, causing the earth to be much warmer than it would be without its atmosphere.

36. The Earth’s Atmosphere

37. Atmospheric CO2 Concentration

38. New Energy Sources • As we search for the energy sources of the future, we need to consider economic, climatic, and supply factors. • There are several potential energy sources: the sun (solar), nuclear processes (fission and fusion), biomass (plants), and synthetic fuels. • Direct use of the sun’s radiant energy to heat our homes and run our factories and transportation systems seems a sensible long-term goal.

39. Coal Conversion One alternative energy source involves using a traditional fuel – coal – in new ways. • Since transportation costs for solid coal are high, more energy-efficient fuels are being developed from coal. • One possibility is to produce a gaseous fuel. • To convert coal from a solid to a gas requires reducing the size of the molecules; the coal structure must be broken down in a process called coal gasification.

40. Hydrogen as a Fuel • The combustion reaction is H2(g) + 1/2O2(g) H2O(l)Ho = -286 kJ The heat of combustion of per gram is approximately 2.5 times that of natural gas. • In addition, hydrogen has a real advantage over fossil fuels in that only product of hydrogen combustion is water; Fossil fuels also produce carbon dioxide. • However, even though it appears that hydrogen is a very logical choice as a major fuel for the future, there are three main problems: the cost of production, storage, and transport.

41. …hydrogen as a Fuel continued… • Although hydrogen is very abundant on earth, virtually none of it exits as the free gas. The main source of hydrogen gas is from the treatment of natural gas with steam: CH4(g) + H2O(g) 3H2(g) + CO(g) This reaction is highly endothermic; treating methane with steam is not an efficient way to obtain hydrogen for fuel. It would be much more economical to burn the methane directly.

42. A virtually inexhaustible supply of hydrogen exists in the waters of the world’s oceans. However, the reaction H2O(l)  H2(g) + 1/2O2(g), requires 286 kJ of energy per mole of liquid water, large-scale production of hydrogen from water is not economically feasible. However, several methods for such production are currently being studied: electrolysis of water, thermal decomposition of water, thermochemical decomposition of water and biological decomposition of water. • A current research goal is to find a system for which the required temperatures are low enough that sunlight can be used as the energy source.

43. Summary • Energy: Potential Energy and Kinetic Energy • Law of Conservation of Energy: energy can’t be created or destroyed • Heat: A means for the transfer of energy • Universe = System + Surroundings • Exothermic Reaction and Endothermic Reaction • First Law of Thermodynamics: Euniv = Constant • Internal Energy, E = q + w • Enthalpy H = E + PV • Rules of Thermochemistry

44. …Summary continued… • Calorimetry: Science of measuring heat • Heat Capacity, specific Heat Capacity, and Molar Heat Capacity: • Energy changed = m x s x T • Energy Released = Energy Absorbed • Hess’s Law: One step or series of step H same • Change in Enthalpy: Horeaction = np Hof (products) -  nr Hof (reactants) • Energy Sources