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(11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: (A) understand energy and.

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Presentation on theme: "(11) Science concepts. The student understands the energy changes that occur in chemical reactions. The student is expected to: (A) understand energy and."— Presentation transcript:

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 Ch. 16 Energy and Chemical Change

3 16.1 Energy  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 PotentialKinetic

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

14

15 Wednesday

16 Phase Changes  ture=player_embedded&v=YG77v1 PwQNM ture=player_embedded&v=YG77v1 PwQNM

17

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 = joules  1000 calorie = 1 Calorie  1J = 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.

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 T final - T initial

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 50.4 o C to 25.0 o C 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/g o C

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 C 50 g metal at C 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 x 0 C) x 125 g x ( C – C) q water = J/(g x 0 C) x 125 g X C q water = 1900 J

31 Calculating Specific Heat  Example  50 g metal at C 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)(115 0 C – C)  c metal = ____1900 J_____ (50.0 g)( C)  c metal = 0.44 J/(g x 0 C) 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  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? H 2 S + 4F 2  2HF + SF 6

38  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 + NH 4 NO 3  NH 4 + NO 3  NH 4 NO 3  NH 4 + NO 3 ΔH = +27 kJ  Energy required to break the bonds in a reactant is less than released after the bonds in the product is formed Endothermic Reaction

39 Exothermic Reaction  If the ∆H is shown on the products side, it is exothermic (losing energy)  The heat of the reaction will be negative.  4 Fe + 3O 2  2 Fe 2 O kJ (energy)  4 Fe + 3O 2  2 Fe 2 O 3 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  H products < H reactants  Endothermic reactions have a positive enthalpy  H products > H reactants

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

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  ΔH vap  Endothermic (positive enthalpy)

45 Molar Enthalpy (heat) of Fusion  The heat required to melt one mole of a solid substance  ΔH fus  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 + 3O 2  2 Fe 2 O kJ (energy)  4 Fe + 3O 2  2 Fe 2 O 3 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 + NH 4 NO 3  NH4 + NO 3  NH 4 NO 3  NH 4 + NO 3 Heat of rxn = 27 kJ


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