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Chapter 5 “Thermochemistry” Credit to: Charles Page High School Stephen L. Cotton Mr Daniel : OHS AP Chemistry.

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Presentation on theme: "Chapter 5 “Thermochemistry” Credit to: Charles Page High School Stephen L. Cotton Mr Daniel : OHS AP Chemistry."— Presentation transcript:

1 Chapter 5 “Thermochemistry” Credit to: Charles Page High School Stephen L. Cotton Mr Daniel : OHS AP Chemistry

2 2 Section 11.1 The Flow of Energy – Heat and Work u OBJECTIVES: Explain how energy, heat, and work are related. Classify processes as either exothermic or endothermic. Identify the units used to measure heat transfer Distinguish between heat capacity and specific heat.

3 3 Energy Transformations u Thermochemistry – study of heat changes that occur during chemical and physical changes. u Energy - capacity for doing work or supplying heat weightless, odorless, tasteless Energy stored within chemical substances is called chemical potential energy Gasoline contains a significant amount of chemical potential energy

4 4 Energy Transformations u Heat (“q”) is energy that transfers from one object to another, because of a temperature difference between them. only changes of heat can be detected! flows from warmer  cooler object All chemical reactions and changes in physical state involve either: a)release of heat, or b)absorption of heat

5 Exothermic and Endothermic Processes u When studying heat changes there are two important categories: the system - the part of the universe on which you focus your attention (a test tube, a room, a molecule, the Earth etc.) the surroundings - includes everything else in the universe u Together, the system and it’s surroundings constitute the universe

6 Exothermic and Endothermic Processes u Thermochemistry is concerned with the flow of heat from the system to it’s surroundings, and vice-versa. 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.

7 Exothermic and Endothermic Processes

8 u Exothermic: Heat flowing out of a system into it’s surroundings defined as negative change  q has a negative value The system loses heat as the surroundings heat up u

9 9 Exothermic and Endothermic u Every reaction has an energy change associated with it. For example, the… Gummy Bear Sacrifice!! u ** Gummy Bear Sacrifice!! ** (next week!!!) u Exothermic reactions release energy, usually in the form of heat. u Endothermic reactions absorb energy Chemical Energy is stored in the form of bonds between atoms.

10 10 Units for Measuring Heat Flow 1)A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 o C. a Calorie, (written with a capital C), is 1000 calories and refers to the energy in food (100 Calorie candy bar is 100,000 calories) 1 Calorie = 1 kilocalorie = 1000 calories

11 11 Units for Measuring Heat Flow 2)The Joule is the SI unit of heat and energy and is related to the calorie: J = 1 cal Named after James Prescott Joule

12 Heat Capacity and Specific Heat u Heat Capacity (Cp) - the amount of heat needed to increase the temperature of an entire object exactly 1 o C u Cop = Joules/ o C Heat capacity is dependent on both the object’s mass and its chemical composition

13 13 Heat Capacity and Specific Heat u Specific Heat Capacity (c) - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 o C Often called simply “Specific Heat” c = Joules/ gram o C

14 Table of Specific Heats

15 15 Heat Capacity and Specific Heat u Water has a HUGE specific heat value, when is compared to other chemicals: C (H 2 O) = 4.18 J/(g o C), or C (H 2 O) = 1.00 cal/(g o C) It heats slowly and It cools slowly Consequently, water maintains a relatively constant temperature!

16 16 Heat Capacity and Specific Heat u Heat calculations for various materials are done with the formula:  q = mass (g) x c x  T  q = change in heat c = specific heat  T = change in temperature Units are either J/(g o C) or cal/(g o C)

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18 18 Section 11.2 Measuring and Expressing Enthalpy Changes u OBJECTIVES: Describe how calorimeters are used to measure heat flow. Construct thermochemical equations. Solve for enthalpy changes in chemical reactions by using heats of reaction.

19 Calorimetry u Calorimetry - the precise measurement of the heat into or out of a system 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

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

21 A foam cup calorimeter – here, two cups are nestled together for better insulation

22 Calorimetry u Changes in heat = Changes in enthalpy =  H u  q =  H These terms are often used interchangeably u Thus,  q =  H = m x c x  T u  H is negative for an exothermic process u  H is positive for an endothermic process

23 Calorimetry u Calorimetry experiments can be performed at a constant volume using a device called a “bomb calorimeter” - a closed system Used by nutritionists to measure energy content of food

24 A Bomb Calorimeter A bomb calorimeter

25 25 C + O 2 → CO 2 Energy Reactants Products  C + O 2 CO 2 395kJ ∆ H= -395 kJ kJ

26 26 Exothermic u In an exothermic rxn the products are lower in energy than the reactants and energy is released. u ΔH rxn = -395 kJ The negative sign indicates that energy is released.

27 27 CaCO 3 → CaO + CO 2 Energy Reactants Products  CaCO 3 CaO + CO kJ CaCO kJ → CaO + CO 2 ∆H = kJ

28 28 Endothermic u The products are higher in energy than the reactants and energy is absorbed. u ΔH rxn = +176 kJ The positive sign indicates that energy is absorbed

29 29 Heat of Reaction u A heat of reaction is the heat change for the equation. The physical state of reactants and products must also be given. Standard conditions for the reaction is kPa (1 atm.) and 25 o C (different from STP)

30 30 Thermochemical Equations An equation that includes energy is called a thermochemical equation  CH 4 + 2O 2  CO 2 + 2H 2 O kJ 1 mole of CH 4 releases kJ of energy. When you make kJ you also make 2 moles of water

31 31 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 Thus, ΔH = kJ for this reaction

32 Summary, so far...

33 33 Enthalpy u Enthalpy: The heat content a substance has at a given temperature and pressure u Can’t 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  H (delta H) is the measure of change in heat

34 34 Enthalpy u Enthalpy: Amount of heat in a system (H) u.u.  If heat is released,  H is negative (exothermic)  If heat is absorbed,  H is positive (endothermic)

35 35 Energy ReactantsProducts  Change is down ΔH is <0 Exothermic

36 36 Energy ReactantsProducts  Change is up ΔH is > 0 Endothermic

37 37 Heat of Reaction  Heat of Reaction: 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 a 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

38 38 Heat of Combustion u The heat from the reaction that completely burns 1 mole of a substance: C + O 2 (g)  CO 2 (g) kJ C + O 2 (g)  CO 2 (g)  H = kJ

39 39 Section 17.3 Heat in Changes of State u OBJECTIVES: Classify the enthalpy change that occurs when a substance melts, freezes, boils, condenses, or dissolves.

40 40 Section 17.3 Heat in Changes of State u OBJECTIVES: Solve for the enthalpy change that occurs when a substance melts, freezes, boils condenses or dissolves.

41 41 Heat in Changes of State u Heat of Fusion (  H fus ) : the heat absorbed by a substance in melting from a solid to a liquid u For H 2 O:  H fus = 334 J/ g u Molar  H fus = 6.01 kJ/ mol u q = mass x  H fus (there is no temperature change)

42 u Molar Heat of Solidification (  H solid ) = heat lost when one mole of liquid solidifies (or freezes) to a solid u For H 2 O:  H fus = -334 J/ g u Molar  H fus = kJ/ mol u q = mass x  H solid (no temperature change)

43 43 Heat in Changes of State u Heat absorbed by a melting solid is equal to heat lost when a liquid solidifies Thus,  H fus = -  H solid u Note Table 17.3, page 522

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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 q = mass x  H vap (no temperature change) u For H 2 O:  H fus = 2260 J/ g u Molar  H fus = 40.7 kJ/ mol

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 to a liquid u  H vap = -  H cond

47 47 Heats of Vaporization and Condensation u Note Figure 17.10, page 523 u The large values for water  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! H 2 0 (g)  H 2 0 (l)  H cond = kJ/mol

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49 49 Heat of Solution (Ch 11) 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 (by dissolving) and interact with water, releasing kJ of heat as  H soln -thus becoming so hot it steams! H 2 O (l)

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