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Thermochemistry Heat and Chemical Change Charles Page High School Dr. Stephen L. Cotton.

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Presentation on theme: "Thermochemistry Heat and Chemical Change Charles Page High School Dr. Stephen L. Cotton."— Presentation transcript:

1 Thermochemistry Heat and Chemical Change Charles Page High School Dr. Stephen L. Cotton

2 2 Section 11.1 The Flow of Energy - Heat u OBJECTIVES: Explain the relationship between energy and heat.

3 3 Section 11.1 The Flow of Energy - Heat u OBJECTIVES: Distinguish between heat capacity and specific heat.

4 4 Energy and Heat u Thermochemistry - concerned with heat changes that occur during chemical reactions u Energy - capacity for doing work or supplying heat weightless, odorless, tasteless if within the chemical substances- called chemical potential energy

5 5 Energy and Heat u Gasoline contains a significant amount of chemical potential energy u Heat - represented by q, is energy that transfers from one object to another, because of a temperature difference between them. only changes can be detected! flows from warmer cooler object

6 Exothermic and Endothermic Processes u Essentially all chemical reactions, and changes in physical state, involve either: release of heat, or absorption of heat

7 Exothermic and Endothermic Processes u In studying heat changes, think of defining these two parts: the system - the part of the universe on which you focus your attention the surroundings - includes everything else in the universe

8 Exothermic and Endothermic Processes u Together, the system and its surroundings constitute the universe u Thermochemistry is concerned with the flow of heat from the system to its surroundings, and vice-versa. Figure 11.3, page 294

9 Exothermic and Endothermic Processes u The Law of Conservation of Energy states that in any chemical or physical process, energy is neither created nor destroyed. All the energy is accounted for as work, stored energy, or heat.

10 Exothermic and Endothermic Processes u Fig. 11.3a, p heat flowing into a system from its surroundings: defined as positive q has a positive value called endothermic –system gains heat as the surroundings cool down

11 Exothermic and Endothermic Processes u Fig. 11.3b, p heat flowing out of a system into its surroundings: defined as negative q has a negative value called exothermic –system loses heat as the surroundings heat up

12 Exothermic and Endothermic u Fig. 11.4, page on the left, the system (the people) gain heat from its surroundings (the fire) this is endothermic u On the right, the system (the body) cools as perspiration evaporates, and heat flows to the surroundings this is exothermic

13 13 Exothemic and Endothermic u Every reaction has an energy change associated with it u Exothermic reactions release energy, usually in the form of heat. u Endothermic reactions absorb energy u Energy is stored in bonds between atoms

14 14 Heat Capacity and Specific Heat u A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 o C. Used except when referring to food a Calorie, written with a capital C, always refers to the energy in food 1 Calorie = 1 kilocalorie = 1000 cal.

15 15 Heat Capacity and Specific Heat u The calorie is also related to the joule, the SI unit of heat and energy named after James Prescott Joule J = 1 cal u Heat Capacity - the amount of heat needed to increase the temperature of an object exactly 1 o C

16 16 Heat Capacity and Specific Heat u Specific Heat Capacity - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 o C (abbreviated C) often called simply Specific Heat Note Table 11.2, page 296 u Water has a HUGE value, compared to other chemicals

17 17 Heat Capacity and Specific Heat u For water, C = 4.18 J/(g o C), and also C = 1.00 cal/(g o C) u Thus, for water: it takes a long time to heat up, and it takes a long time to cool off! u Water is used as a coolant! Note Figure 11.7, page 297

18 18 Heat Capacity and Specific Heat u To calculate, use the formula: u q = mass (g) x T x C u heat abbreviated as q u T = change in temperature u C = Specific Heat u Units are either J/(g o C) or cal/(g o C) u Sample problem 11-1, page 299

19 19 Section 11.2 Measuring and Expressing Heat Changes u OBJECTIVES: Construct equations that show the heat changes for chemical and physical processes.

20 20 Section 11.2 Measuring and Expressing Heat Changes u OBJECTIVES: Calculate heat changes in chemical and physical processes.

21 Calorimetry u Calorimetry - the accurate and precise measurement of heat change for chemical and physical processes. u The device used to measure the absorption or release of heat in chemical or physical processes is called a Calorimeter

22 Calorimetry u Foam cups are excellent heat insulators, and are commonly used as simple calorimeters Fig. 11.8, page 300 u For systems at constant pressure, the heat content is the same as a property called Enthalpy (H) of the system

23 Calorimetry u Changes in enthalpy = H u q = H These terms will be used interchangeably in this textbook u Thus, q = H = m x C x T u H is negative for an exothermic reaction u H is positive for an endothermic reaction (Note Table 11.3, p.301)

24 Calorimetry u Calorimetry experiments can be performed at a constant volume using a device called a bomb calorimeter - a closed system u Figure 11.9, page 301 u Sample 11-2, page 302

25 25 C + O 2 CO 2 Energy ReactantsProducts C + O 2 C O 2 395kJ kJ

26 26 In terms of bonds C O O C O O Breaking this bond will require energy. C O O O O C Making these bonds gives you energy. In this case making the bonds gives you more energy than breaking them.

27 27 Exothermic u The products are lower in energy than the reactants u Releases energy

28 28 CaCO 3 CaO + CO 2 Energy ReactantsProducts CaCO 3 CaO + CO kJ CaCO kJ CaO + CO 2

29 29 Endothermic u The products are higher in energy than the reactants u Absorbs energy u Note Figure 11.11, page 303

30 30 Chemistry Happens in MOLES u An equation that includes energy is called a thermochemical equation CH 4 + 2O 2 CO 2 + 2H 2 O kJ u 1 mole of CH 4 releases kJ of energy. u When you make kJ you also make 2 moles of water

31 31 Thermochemical Equations u A heat of reaction is the heat change for the equation, exactly as written The physical state of reactants and products must also be given. Standard conditions for the reaction is kPa (1 atm.) and 25 o C

32 32 CH O 2 CO H 2 O kJ u If grams of CH 4 are burned completely, how much heat will be produced? g CH g CH 4 1 mol CH kJ = 514 kJ

33 33 CH O 2 CO H 2 O kJ u How many liters of O 2 at STP would be required to produce 23 kJ of heat? u How many grams of water would be produced with 506 kJ of heat?

34 Summary, so far...

35 35 Enthalpy u The heat content a substance has at a given temperature and pressure u Cant be measured directly because there is no set starting point u The reactants start with a heat content u The products end up with a heat content u So we can measure how much enthalpy changes

36 36 Enthalpy u Symbol is H Change in enthalpy is H (delta H) u If heat is released, the heat content of the products is lower H is negative (exothermic) u If heat is absorbed, the heat content of the products is higher H is positive (endothermic)

37 37 Energy ReactantsProducts Change is down H is <0

38 38 Energy ReactantsProducts Change is up H is > 0

39 39 Heat of Reaction u The heat that is released or absorbed in a chemical reaction Equivalent to H C + O 2 (g) CO 2 (g) kJ C + O 2 (g) CO 2 (g) H = kJ u In thermochemical equation, it is important to indicate the physical state H 2 (g) + 1/2O 2 (g) H 2 O(g) H = kJ H 2 (g) + 1/2O 2 (g) H 2 O(l) H = kJ

40 40 Heat of Combustion u The heat from the reaction that completely burns 1 mole of a substance u Note Table 11.4, page 305

41 41 Section 11.3 Heat in Changes of State u OBJECTIVES: Classify, by type, the heat changes that occur during melting, freezing, boiling, and condensing.

42 42 Section 11.3 Heat in Changes of State u OBJECTIVES: Calculate heat changes that occur during melting, freezing, boiling, and condensing.

43 43 Heats of Fusion and Solidification u Molar Heat of Fusion ( H fus ) - the heat absorbed by one mole of a substance in melting from a solid to a liquid u Molar Heat of Solidification ( H solid ) - heat lost when one mole of liquid solidifies

44 44 Heats of Fusion and Solidification u Heat absorbed by a melting solid is equal to heat lost when a liquid solidifies Thus, H fus = - H solid u Note Table 11.5, page 308 u Sample Problem 11-4, page 309

45 45 Heats of Vaporization and Condensation u When liquids absorb heat at their boiling points, they become vapors. u Molar Heat of Vaporization ( H vap ) - the amount of heat necessary to vaporize one mole of a given liquid. u Table 11.5, page 308

46 46 Heats of Vaporization and Condensation u Condensation is the opposite of vaporization. u Molar Heat of Condensation ( H cond ) - amount of heat released when one mole of vapor condenses u H vap = - H cond

47 47 Heats of Vaporization and Condensation u Note Figure 11.5, page 310 u The large values for H vap and H cond are the reason hot vapors such as steam is very dangerous You can receive a scalding burn from steam when the heat of condensation is released!

48 48 Heats of Vaporization and Condensation u H 2 0 (g) H 2 0 (l) H cond = kJ/mol u Sample Problem 11-5, page 311

49 49 Heat of Solution u Heat changes can also occur when a solute dissolves in a solvent. u Molar Heat of Solution ( H soln ) - heat change caused by dissolution of one mole of substance u Sodium hydroxide provides a good example of an exothermic molar heat of solution:

50 50 Heat of Solution NaOH (s) Na 1+ (aq) + OH 1- (aq) H soln = kJ/mol u The heat is released as the ions separate and interact with water, releasing kJ of heat as H soln thus becoming so hot it steams! u Sample Problem 11-6, page 313 H 2 O (l)

51 51 Section 11.4 Calculating Heat Changes u OBJECTIVES: Apply Hesss law of heat summation to find heat changes for chemical and physical processes.

52 52 Section 11.4 Calculating Heat Changes u OBJECTIVES: Calculate heat changes using standard heats of formation.

53 53 Hesss Law u If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction. Called Hesss law of heat summation u Example shown on page 314 for graphite and diamonds

54 54 Why Does It Work? u If you turn an equation around, you change the sign: If H 2 (g) + 1/2 O 2 (g) H 2 O(g) H= kJ then, H 2 O(g) H 2 (g) + 1/2 O 2 (g) H = kJ u also, u If you multiply the equation by a number, you multiply the heat by that number: 2 H 2 O(g) H 2 (g) + O 2 (g) H = kJ

55 55 Why does it work? u You make the products, so you need their heats of formation u You unmake the products so you have to subtract their heats. u How do you get good at this?

56 56 Standard Heats of Formation The H for a reaction that produces 1 mol of a compound from its elements at standard conditions u Standard conditions: 25°C and 1 atm. u Symbol is u The standard heat of formation of an element = 0 u This includes the diatomics

57 57 What good are they? u Table 11.6, page 316 has standard heats of formation u The heat of a reaction can be calculated by: subtracting the heats of formation of the reactants from the products H o = (Products) -(Reactants)

58 58 Examples CH 4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O(g) CH 4 (g) = kJ/molO 2 (g) = 0 kJ/molCO 2 (g) = kJ/molH 2 O(g) = kJ/mol H= [ (-241.8)] - [ (0)] H= kJ

59 59 Examples u Sample Problem 11-7, page 317


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