1. Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed. by Steven S. Zumdahl & Donald J. DeCoste University of Illinois

2. Chapter 10 Energy

3. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 3 Energy and Energy Changes Energy: ability to do work or produce heat –Chemical, mechanical, thermal, electrical, radiant, sound, nuclear –Potential and kinetic Energy may affect matter. –e.g. Raise its temperature, eventually causing a state change, or cause a chemical change such as decomposition All physical changes and chemical changes involve energy changes.

4. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 4 Energy and Energy Changes Potential Energy: energy due to composition or position Kinetic Energy: energy due to motion – - ½ mv 2

5. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 5 Energy and Energy Changes (cont.) Law of Conservation of Energy: energy can be converted from one form to another, but cannot be created or destroyed

6. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 6 Work and Energy Work: force acting over a distance –w = f d –Work done on a system will increase the energy of the system, whereas work done by the system will decrease the energy of the system State function: a property that changes independent of pathway

7. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 7 Temperature and Heat Heat: a flow of energy due to a temperature difference Temperature: a measure of the random motions of the components of a substance

8. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 8 Temperature and Heat (cont.)

9. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 9 Exothermic vs. Endothermic System: that part of the universe that we wish to study Surroundings: everything else in the universe Exothermic process: is any process that gives off heat – transfers thermal energy from the system to the surroundings. Example: when a match is struck, it is an exothermic process because energy is produced as heat. and 2H 2 (g) + O 2 (g) 2H 2 O (l) + energy Endothermic process: absorbs heat -Example: melting ice to form liquid water is an endothermic process because the ice absorbs heat in order to melt energy + H 2 O (s) H 2 O (l)

10. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 10 Exothermic Process

11. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 11 Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.  H = H (products) – H (reactants)  H = heat given off or absorbed during a reaction at constant pressure

12. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 12 Thermochemical Equations H 2 O (s) H 2 O (l)  H = 6.01 kJ Is  H negative or positive? System absorbs heat Endothermic  H > 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm. 6.3

13. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 13 H 2 O (s) H 2 O (l)  H = 6.01 kJ The stoichiometric coefficients always refer to the number of moles of a substance Thermochemical Equations If you reverse a reaction, the sign of  H changes H 2 O (l) H 2 O (s)  H = - 6.01 kJ If you multiply both sides of the equation by a factor n, then  H must change by the same factor n. 2H 2 O (s) 2H 2 O (l)  H = 2 x 6.01 = 12.0 kJ

14. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 14 H 2 O (s) H 2 O (l)  H = 6.01 kJ The physical states of all reactants and products must be specified in thermochemical equations. Thermochemical Equations H 2 O (l) H 2 O (g)  H = 44.0 kJ How much heat is evolved when 266 g of white phosphorus (P 4 ) burn in air? P 4 (s) + 5O 2 (g) P 4 O 10 (s)  H = -3013 kJ 266 g P 4 1 mol P 4 123.9 g P 4 x 3013 kJ 1 mol P 4 x = 6470 kJ

15. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 15 Different enthalpies Heat of reaction (  H r or  H rxn )- heat energy absorbed or released during a reaction. Heat of formation (  H f  )- heat energy absorbed or released during synthesis of one mole of a compound from its elements at 298 K and 1 atm pressure (STP- standard temp and pressure). Heat of solution (  H sol )- heat energy absorbed or released when a substance dissolves in a solvent. Heat of combustion (  H comb )- heat energy released when a substance reacts with oxygen to form CO 2 and H 2 O.

16. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 16 Heat of fusion (  H fus )= Energy needed to melt one mole (solid to liquid) Heat of vaporization (  H vap )=Energy needed to boil one mole (liquid to gas) In a phase change graph, it is possible to calculate the total energy involved as well as the energy consumed in each step. Note that water has different values for sp. heat depending upon its physical state.

17. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 17 Thermodynamics The Law of Conservation of Energy is also known as The First Law of Thermodynamics. It can be stated as “the energy of the universe is constant.” Internal Energy (E) = kinetic energy + potential energy ΔE = q + w = change in internal energy q = heat absorbed by the system w = work done on the system

18. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 18 Units of Energy One calorie = amount of energy needed to raise the temperature of one gram of water by 1°C –kcal = energy needed to raise the temperature of 1000 g of water 1°C joule –4.184 J = 1 cal In nutrition, calories are capitalized. –1 Cal = 1 kcal

19. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 19 Example - Converting Calories to Joules Convert 60.1 cal to joules.

20. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 20 Energy & Temperature of Matter The amount the temperature of an object increases depends on the amount of heat added (q). –If you double the added heat energy the temperature will increase twice as much. The amount the temperature of an object increases when heat is added depends on its mass –If you double the mass it will take twice as much heat energy to raise the temperature the same amount.

21. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 21 Specific Heat Capacity Specific heat (s): the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius Amount of Heat = Specific Heat x Mass x Temperature Change Q = s x m x  T

22. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 22 Specific Heat Capacity

23. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 23 Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C. Example #1:

24. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 24 Mass = 7.40 g Temperature change = 46.0°C – 29.0°C = 17.0°C Q = s m  T Example #1 (cont.) Specific heat of water = 4.184 C g J 

25. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 25 A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold? Example #2

26. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 26 Example #2 Table 10.1 lists the specific heat of gold as 0.13 Therefore the metal cannot be pure gold.

27. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 27 Enthalpy Change in enthalpy (ΔH p = q p ): the amount of heat exchanged when heat exchange occurs under conditions of constant pressure Enthalpy is a state function ΔH is independent of the path taken

28. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 28 Hess’s Law: in going from a set of reactants to a set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. Hess’s Law

29. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 29 Hess’ Law (cont.) ΔH reaction = ∑Δh steps If the direction of a reaction is reversed, the sign of ΔH is reversed. ΔH forward = -Δh reverse Magnitude of ΔH α quantities of reactants and products

30. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 30 Overall reaction: N 2 + 2O 2 2NO 2 ΔH = 68 kJ This reaction can be carried out in 2 steps: N 2 + O 2 2NO ΔH = 180 kJ 2NO + O 2 2NO 2 ΔH = -112 kJ -------------------------------------------------------- N 2 + 2O 2 2NO 2 ΔH = 68 Kj Note: the sum of the two reactions gives the overall reaction and the same is true for the sum of the enthalpy change values. Hess’s Law (cont.)

31. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 31 Calorimetry The amount of heat flow transferred during a reaction is determined from temperature measurements made in a calorimeter. Heat loss is minimized by having insulation. A simple calorimeter can be made in the lab by stacking 2 styrofoam cups. A calorimeter minimizes heat exchange between the system and the surroundings. Amount of heat produced is calculated by measuring the temp change in the surrounding water.  H=  t (H20) X m H2O XCp H2O

32. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 32 Calorimetry (cont.)

33. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 33 Energy Quality & Quantity While the total amount or quantity of energy in the universe is constant (1st Law) the quality of energy is degraded as it is used. Burning of petroleum: High grade concentrated energy Low grade energy (heat)

34. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 34 Fuels Petroleum –A fossil fuel composed mainly of hydrocarbons Natural gas –Consists largely of methane –Also contains ethane, propane, and butane

35. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 35 Fuels (cont.)

36. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 36 Fuels (cont.)

37. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 37 Fuels (cont.) Coal –Matures geologically through stages

38. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 38 Global Warming

39. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 39 Global Warming (cont.)

40. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 40 Energy Use and Sources

41. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 41 Energy as a Driving Force Most processes that occur spontaneously involve an “energy spread.” –Heat flows from high to low temperature and “spreads” …or a “matter spread” –Salt dissolves or “spreads” in water

42. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 42 Entropy Entropy (S) is a measure of disorder or randomness. –As a system becomes more disordered, ΔS >0 Second Law of Thermodynamics: the entropy of the universe is always increasing. Tendency in nature is to increase disorder (unless external forces counteract). Ex—messy room, throwing a puzzle -Entropy in solids< liquids<gases

43. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 43 Entropy=S (absolute Entropy for a substance=Entropy at absolute zero temp -273K). Units J/K.mol - change in Entropy=  S  S=S products -S reactants –-Entropy can be the driving force behind reactions. Ex- reactions leading to the formation of gases (from solids) are favored. –-Higher temp=higher Entropy (due to more KE of particles) –-Lower temp=lower Entropy (less KE) –-Ideal conditions for a spontaneous reaction Increase Entropy (disorder)  Decrease enthalpy of products 

44. Copyright © Houghton Mifflin Company. All rights reserved. 10 | 44 - Reactions can thus occur even if not entropically favored. Ex- 2H 2 (g)+O 2 (g)  2H 2 O (g) 3 molecules  2 molecules (decrease in entropy- same physical state) But reaction is highly exothermic (ie. Much lower enthalpy of products)