1. (11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to:
(A) understand energy and its forms, including kinetic, potential, chemical, and thermal energies;
(B) understand the law of conservation of energy and the processes of heat transfer;
(C) use thermochemical equations to calculate energy changes that occur in chemical reactions and classify reactions as exothermic or endothermic;
(D) perform calculations involving heat, mass, temperature change, and specific heat; and
(E) use calorimetry to calculate the heat of a chemical process.

2. Energy and Chemical Change

3. Energy- the ability to do work or produce heat 2 Forms:
Potential Energy
Kinetic Energy

4. Potential Energy
Potential Energy -energy due to the composition or position of an object.
Ex: water stored behind a dam
depends on composition:
1. the type of atoms
2. the number and type of chemical bonds joining the atoms
3. the way the atoms are arranged.

5. Kinetic Energy
Kinetic Energy – is the energy of motion
Ex: water flows from the dam

6. Chemical systems contain both potential and kinetic energy
Potential Kinetic

7. Heat- represented by symbol Q- energy that is in the process of flowing from a warmer object to a cooler object

8. Chemical Potential Energy -
the energy stored in a substance because of its composition.
Composition is the type, number, and arrangement of atoms and bonds.

9. Thermal energy
the energy created by moving particles inside a substance.
more movement of particles = more thermal energy

10. Heat is Thermal energy that is transferred
Heat is Transferred in 3 ways
Conduction – the way heat moves through solids. (direct transfer)
Vibrating molecules pass on heat from molecule to molecule.

11. Convection – the way heat moves through gases and liquids.
Heated molecules move AWAY from the heat and cooler molecules take their place.
Ex: Hot air rises and cool air sinks

12. Radiation
Radiation – the way heat moves through empty space.
Does not need atoms or molecules to work.
Electromagnetic radiation – light and heat from the sun, visible light, microwaves, X-rays, etc.

13. Forms of Energy

15. Wednesday

16. Phase Changes

18. Specific Heat –is the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius.
each substance has its own specific heat
Table 16-2 pg 492

19. Heat of Vaporization
The amount of heat required to convert unit mass of a liquid into the vapor without a change in temperature.

20. Heat of Fusion
The amount of heat required to convert unit mass of a solid into the liquid without a change in temperature.

21. Measuring HEAT!!!

22. Two units for measuring heat
calorie - the amount of heat required to raise the temperature of one gram of pure water by one degree Celsius
Joule - SI unit of heat and energy

23. 1 calorie = 4.184 joules 1000 calorie = 1 Calorie 1J = 0.2390 calories
Table 16-1 Conversion factors and relationships pg 491

24. Calories are nutritional or food Calories
1 Calorie = 1000 calories
1Calorie = 1 kilocalorie
approximates the energy needed to increase the temperature of 1 kilogram of water by 1 °C.
The small calorie or gram calorie (symbol: cal)[2] approximates the energy needed to increase the temperature of 1 gram of water by 1 °C. This is about 4.2 joules.
The large calorie, kilogram calorie, dietary calorie, or food calorie (symbol: Cal)[2] approximates the energy needed to increase the temperature of 1 kilogram of water by 1 °C. This is exactly 1,000 small calories or about 4.2 kilojoules. It is also called the nutritionist's calorie.

25. Calculating Specific Heat
Q = m x c x ΔT
Q = heat absorbed or released
m = mass of the sample in grams
c = specific heat of the substance
ΔT = difference between final temperature and initial temperature, or Tfinal- Tinitial

26. 16.2 Heat in Chemical Reactions and Processes
Measuring Heat
Heat changes are measured with a calorimeter

27. Lab and worksheet
The temperature of a sample of iron has a mass of 10.0g changed from 50.4oC to 25.0oC with the release of 114 J of heat. What is the specific heat of iron? Q = mc∆T 114 = 10 x c x ( ) 114 = 254c C = 114/254 = J/goC

28. Calorimeter – an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.
Data is the change in temperature of this mass of the substance.

29. Determining Specific Heat
Place a hot metal into water.
Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal.
The heat gained by the water is equal to the heat lost by the metal

30. Calculating Heat Example
125 g water with an Initial temperature of 25.60C
50 g metal at 1150C is placed in the water.
Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal. Both have a final temperature of C. Calculate the Heat gained by the water.
Example Part A:
q = c x m x /\T
q water = J/(g x0C) x 125 g x (29.30C – 25.60C)
q water = J/(g x0C) x 125 g X 3.7 0C
q water = 1900 J

31. Calculating Specific Heat
Example
50 g metal at 1150C is placed in the water.
Heat flows from the hot metal to the cooler water until the temperature of the metal and water are equal. Both have a final temperature of C.
Water absorbed 1900 J of heat.
Example Part B: Calculate the Specific Heat of the Metal
c = q___
m x /\T
c metal = 1900 J
m x /\T
c metal = _______1900 J_________
(50.0 g)(1150C – C)
c metal = ____1900 J_____
(50.0 g)(85.700C)
c metal = 0.44 J/(g x 0C) specific heat of the metal
Look at pg 492 at the table. What is this metal?

32. Thursday- Lab

33. Friday- Practice worksheet

34. Monday

35. 16.3 and 16.4 Enthalpy and Enthalpy Changes
Enthalpy- (H) the heat content of a system at a constant pressure

36. A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change expressed as the change in enthalpy, ∆H.

37. Use the table on pg. 510 in your textbook
You can’t measure actual enthalpy, but you can measure change in enthalpy, which is called enthalpy (heat) of reaction (ΔH rxn)
Use the table on pg. 510 in your textbook
ΔH rxn = H final – H initial or
ΔH rxn = H products – H reactants
Example:
What is the heat of reaction for the following reaction? H2S + 4F2  2HF + SF6

38. Endothermic Reaction
If the ∆H is shown on the reactants side, it is endothermic (gaining energy)
The heat of the reaction will be positive.
(energy) 27 kJ + NH4NO3  NH4 + NO3
NH4NO3  NH4 + NO3 ΔH = +27 kJ
Energy required to break the bonds in a reactant is less than released after the bonds in the product is formed

39. Exothermic Reaction The heat of the reaction will be negative.
If the ∆H is shown on the products side, it is exothermic (losing energy)
The heat of the reaction will be negative.
4 Fe + 3O2  2 Fe2O kJ (energy)
4 Fe + 3O2  2 Fe2O3 ΔH = kJ
Energy needed to break the bond in the reactant is more than energy released after the bonds in the products are formed

40. END

41. Sign of the Enthalpy of Reaction
Exothermic reactions have a negative enthalpy
Hproducts < Hreactants
Endothermic reactions have a positive enthalpy
Hproducts > Hreactants

42. 16.3 Thermochemical Equations
Enthalpy (heat) of combustion- enthalpy change for the complete burning of one mole of the substance
ΔHcomb

43. Entropy
Measure of the disorder or randomness of the particles that make up a system
Symbolized by S

44. Molar Enthalpy (heat) of Vaporization
Heat required to vaporize one mole of a liquid
ΔHvap
Endothermic (positive enthalpy)

45. Molar Enthalpy (heat) of Fusion
The heat required to melt one mole of a solid substance
ΔHfus
Endothermic (positive enthalpy)

46. 16.5 Reaction Spontaneity
Spontaneous process- physical or chemical change that occurs with no outside intervention

47. Law of Disorder
States that spontaneous processes always proceed in such a way that the entropy of the universe increases

48. Chemical Energy and the Universe
Thermochemistry – the study of heat changes that accompany chemical reactions and phase changes.

49. system – the specific part of the universe that contains the reaction or process you wish to study.
surroundings – everything in the universe other than the system

50. universe – the system plus the surroundings
universe = system + surroundings

51. Example: Using a heat pack to warm your hands
Heat flows from the heat pack (the system) to your cold hands (surroundings)
Exothermic - If energy is shown as a product it means that heat is released. The heat of the reaction will be negative.
4 Fe + 3O2  2 Fe2O kJ (energy)
4 Fe + 3O2  2 Fe2O3 Heat of rxn = kJ

52. Example: Using a cold pack on an injured knee
Heat flows from the knee (the surroundings) to the cold pack (the system)
Endothermic – If energy is shown as a reactant it means that energy is absorbed.
The heat of the reaction will be positive.
(energy) 27 kJ + NH4NO3  NH4 + NO3
NH4NO3  NH4 + NO3 Heat of rxn = 27 kJ