Thermochemistry

1. Thermochemistry Chapters 16 and 17 Thermodynamics – the study of energy and energy transfer. Thermochemistry – the study of energy involved in chemical reactions.

2. Law of Conservation Of Energy – the total energy in the universe is constant. - energy cannot be created or destroyed.  ∆ Energy Universe = 0 Energy CAN: - be transferred from one substance to another - or be converted into various forms.

3. In order to study energy changes – we need to define what part of the universe we are dealing with. System - the part of the universe that is being studied or observed. Surroundings – everything else in the universe. Ex. In this chemical rxn – the system is the reactants and products in the flask - the surroundings is the flask, the student, the air, and the hand holding the flask.

4. Therefore: Universe = System + Surroundings ∆ E universe = ∆ E system + ∆ E surroundings = 0 This relationship is known as the 1st Law of Thermodynamics: “Any change in the energy of a system is accompanied by an equal and opposite change in the energy of the surroundings” ∆ E system = - ∆ E surroundings the energy the system gains = the energy the surroundings loses OR - ∆ E system = ∆ E surroundings the energy the system loses = the energy the surroundings gains

5. Depending on how the system is separated from its surroundings, systems are defined in three different ways: Open System – system is open to the surroundings – both matter and energy can be exchanged between the two. Ex. A reaction in an open beaker. Closed System – a system where energy can move between the system and the surroundings – however, matter cannot move. Ex. A reaction in a stoppered Erlenmeyer flask. Isolated System – a system is completely isolated from its surroundings – neither matter nor energy is exchanged between the two. Ex. A reaction in a calorimeter.

6. Types Of Energy 1) Kinetic Energy - energy in motion (always accompanied by a temperature change) 2) Potential Energy – energy that is stored (changes in state and chemical rxns) • The SI unit for both types of energy is the joule (J). • One joule is equal to 1 kg∙m2/s2 • Kinetic Energy is measured in J. • Potential Energy is measured in kJ.

7. Temperature Change and Heat • Temperature, T, is a measure of the average kinetic energy of the particles that make up a system. • Temperature is measured in either Celsius degrees (oC) or kelvins (K). • Temperature change in a substance is an indication of a change in kinetic energy. - Symbolized by ∆T ∆T = Tf – Ti change in temp = temp. final – temp. initial

8. Transfer of Kinetic Energy • Heat, q, refers to the transfer of kinetic energy between objects with different temperatures. • It is measured in joules. • When a substance absorbs heat, the average kinetic energy of the particles increases – therefore the temperature increases. Reconsidering the 1st Law of Thermodynamics: q system = -q surroundings The heat the system gains is equal to the heat the surroundings loses. OR - q system = q surroundings The heat the system loses is equal to the heat the surroundings absorbs.

9. When substances with different temperatures come in contact, kinetic energy is transferred from the particles of the warmer substance to the particles of the cooler substance. Heat Transfers: - from hot chocolate to your mouth - from hot chocolate to the mug - from the mug to your fingers - from the hot chocolate to the air You blow on the hot chocolate - moves cool air over the hot chocolate and the hot chocolate transfers heat to the air.

10. Factors In Heat Transfer

11. Factors In Heat Transfer Temperature and Energy Transfer • Heat moves from greater to lower temperatures • The greater the heat difference the greater the energy transfer Mass and Energy Transfer • Mass is directly related to heat transfer. • Is used in calculations – designated symbol m. Types of Substances and Energy Transfer • In calculating heat transfer, we do not use the type of substance as a variable, but rather we use a variable that reflects the individual nature of different substances – the specific heat capacity.

12. Specific Heat Capacity (c) - the amount of energy required to change one gram of a substance by one degree Celsius. • This reflects how well a substance stores energy. • Units are J/g∙oC • A substance with a large specific heat capacity can absorb and release more energy that a substance with a smaller specific heat capacity. • Water has a very large specific heat capacity • = 4.184 J/g∙oC • See Table on Page 632

13. Calculating Heat Transfer Three variables determine heat transfer: 1) temperature 2) mass 3) specific heat capacity Mathematical Relationship:

14. Complete Practice Problems #1-4 Page 634

15. Basic Calorimetry Calorimetry • the technological process of measuring the changes in kinetic energy. • uses a calorimeter Calorimeter – a basic calorimeter contains water, a thermometer, and an isolated system. We will use a coffee-cup calorimeter or a Poor Man’s Calorimeter

16. Basic Principles of Calorimetry • The system is isolated – no heat is exchanged with the surroundings outside the calorimeter. • The amount of heat exchanged with the calorimeter itself is small enough to be ignored. • If something dissolves in or reacts with the calorimeter water, the solution still retains the properties of water. • specific heat capacity and density

17. Therefore: q system = -q surroundings The heat the system absorbs is equal to the heat that the surroundings released. OR - q system = q surroundings The heat the system releases is equal to the heat the surroundings absorbs. Sample Problem on Page 663 Page 664 - #1-4

18. 16.2 Enthalpy Changes Now we will consider changes in the potential energy of a system.... Enthalpy (∆H) • refers to the potential energy change of a system during a process such as a chemical or physical change. • measured at constant pressure • units are kJ/mol

19. Enthalpy Changes in Chemical Rxns • The enthalpy change of a chemical reaction represents the difference between the potential energy of the products and the potential energy of the reactants. • In chemical rxns, potential energy changes result from the breaking or the forming of bonds. • A chemical bond is caused by the attraction between electrons and the nuclei of the atoms. Energy is stored in bonds. • Breaking a bond will require energy. • Creating a bond will release energy.

20. Net Absorption of Energy = Endothermic Rxn Net Release of Energy = Exothermic Rxn

21. Chemists define the total energy of a substance at a constant pressure as its enthalpy, H. They use relative enthalpy of the reactants and the products to determine the change in enthalpythat accompanies the reaction. Consider: N2 + O2→ 2NO triple bonds double bonds single bonds This rxn absorbs energy -  less energy is released when NO bonds are formed than the energy required to break the double and the triple bonds

22. Representing Enthalpy Changes ∆H Rxn- enthalpy of reaction - enthalpy change of a rxn - dependent on conditions of temperature and pressure - it is the energy change associated with the reaction of 1 mole of a compound with another compound.  ∆HoRxn will represent the enthalpy of rxn at SATP (SATP = 1 bar, 25oC) ** people will say “heat of reaction” ** O means nought – means standard conditions or standard state

23. 3 Ways To Represent Enthalpy Changes 1. Thermochemical Equation - a balanced chemical equation that indicates the amount of heat absorbed or released in a chemical rxn. ENDOTHERMIC - because heat is absorbed in an endothermic rxn, the heat term is included on the reactant side of the equation. 117.3 kJ + MgCO3(s)→MgO(s) + CO2(g) EXOTHERMIC - because heat is released in an exothermic rxn, the heat term is included on the product side of the equation. H2 (g) + ½ O2 (g)→ H2O (l) + 285.8 kJ

24. 2. Separate Equation Method - for this method, the heat term is included on the right side of the equation. - the sign for the heat term must be included to show when the equation is endothermic or exothermic. ENDOTHERMIC MgCO3(s)→MgO(s) + CO2(g) ∆Ho = +117.3 kJ EXOTHERMIC H2 (g) + ½ O2 (g)→ H2O (l) ∆Ho = - 285.8 kJ

25. You should always include: • Title • Correct graph • reactants • products • y-axis • x-axis • ∆H Enthalpy Diagrams Enthalpy Diagram for an Exothermic Rxn

27. Types of Reactions • In thermochemistry: - formation - combustion - melting - freezing - condensation - vaporization - solution

28. Standard Molar Enthalpy Of Formation • In a formation reaction, a substance is formed from its elements in their standard states. ∆Hoform = standard molar enthalpy of formation - the quantity of energy that is absorbed or released when 1 mol of a substance is formed directly from its elements in their *standard states. - it is a value – a number. * The standard state of an element is its most stable form at SATP.

29. Formation Equation for water (liquid) Thermochemical Equation: 1 H2 (g) + ½ O2 (g)→ 1 H2O (l) Separate Equation Method: 1 H2 (g) + ½ O2 (g)→ 1 H2O (l) + 285.8 kJ ∆Hoform = -285.8 kJ

30. Standard Molar Enthalpy of Combustion • In a combustion equation, a substance burns in oxygen to form carbon dioxide and water. ∆Hocomb= the molar enthalpy of combustion - the quantity of energy that is absorbed or released when 1 mol of a compound burns in oxygen to produce carbon dioxide and water. - reactants and products should be in their standard states.

31. Combustion of Propane Thermochemical Equation: 1 C3H8 (g) + 5 O2 (g)→ 3 CO2 (g) + 4 H2O (l) + 2323.7 kJ Separate Equation Format: 1C3H8 (g) + 5 O2 (g)→ 3 CO2 (g)+ 4 H2O (l) ∆Hocomb= -2323.7 kJ Combustion Of Butane Thermochemical Equation: 1 C4H10 (g) + O2 (g)→ 4 CO2 (g) + 5 H2O (l) + 3003.0 kJ Separate Equation Format: 1C4H10 (g) + O2 (g)→ 4 CO2 (g)+ 5 H2O (l) ∆Hocomb= -3003.0 kJ Page 643 #15-18

32. Calculating Enthalpy Changes • Use Stoichiometry!!! How much heat is released when 50.00 g of methane forms from its elements? 50.00 g CH4 change from use the molar enthalpy grams to moles (obtained from the chart)  232.4 kJ was released.

33. How much heat is released when 50.00 g of methane combusts completely? 50.00 g CH4 Page 645 #19-23 Remember that: at SATP: 1 mol of a gas = 22.4 L of the gas 1L = 1000 mL

34. Enthalpy Changes and Changes in State • Changes in state involve changes in the potential energy of a system. • Changes that you know: melting, freezing, etc. • The temperature of the system undergoing the change remains constant but because the energy is being absorbed or released as heat, the temperature of the surroundings often changes.

35. In general, the energy change associated with changes in state are much smaller than those of chemical changes. • Why? Intermolecular forces (L.D., D.D., H.B.) are easier to break than intramolecular forces (ionic, covalent, metallic bonds). • State changes for water animation • Solid changing state

36. Representing Enthalpy Changes For Changes In State • ∆Homelt= the molar enthalpy of melting (fusion) - the quantity of energy that is absorbed when 1 mol of a compound changes state from a solid to a liquid. Ex. ∆Homelt = 27.2 kJ 27.2 kJ + NaCl(s)→ NaCl(l ) • ∆Hofre= the molar enthalpy of freezing - the quantity of energythat is released when 1 mol of a compound changes state from a liquid to a solid. Ex. ∆Hofre = - 27.2 kJ NaCl(l )→ NaCl (s) + 27.2 kJ

37. ∆Hovap= the molar enthalpy of vaporization - the quantity of energy that is absorbed when 1 mol of a compound changes state from a liquid to a gas. 207 kJ + NaCl (l )→ NaCl(g) • ∆Hocond= the molar enthalpy of condensation - the quantity of energy that is released when 1 mol of a compound changes state from a gas to a liquid. NaCl(g)→ NaCl(l ) + 207 kJ Note: ∆Homelt= -∆Hofre ∆Hovap= - ∆Hocond See Table 16-4 on Page 647

38. ∆Hosoln= the molar enthalpy of solution - the quantity of energy that is absorbed or released when 1 mol of a compound dissolves in a solvent. • Dissolving can be exothermic or endothermic. • Manufacturers use this information to make hot packs and/or cold packs. • Hot Packs – process of dissolving is exothermic • Cold Packs – process of dissolving is endothermic Ex. 25.7 kJ + NH4NO3 (s)→ NH4NO3 (aq) Ex. CaCl2 (s)→ CaCl2 (aq) + 82.8 kJ

39. Calorimetry for Potential Energy Changes • We can re-look at the theory for calorimetry when considering state changes and chemical rxns. Consider Melting Ice: When a measured amount of hot water is placed in a calorimeter and ice is added to it... • heat travels from the water into the ice. • the energy the ice absorbs is used to melt the ice.  - q water = + q ice The heat the water releases is equal to the heat the ice absorbs Better: The heat the water releases is equal to the enthalpy the ice absorbs

40. Water undergoes a kinetic energy change (there is a temperature change) •  we can use q = mc∆T • Then, - q water = + q ice • Ice undergoes a potential energy change (no temperature change) •  we can use stoichiometry. • Note: We do not mass the ice in an analytical balance. Take the mass of water when the lab is finished and subtract the initial mass of water and you will obtain the mass of ice that melted.

41. Sample Problem: A student takes 100.0 mL of water at 51.8oC and places it in a calorimeter. Ice is added to the water until it reaches 0.0oC. At that point, all the ice is removed and the water is re-measured in a graduated cylinder. The final volume of water is 165.7 mL. Calculate the molar enthalpy of melting.

42. qw = mwcw∆Tw qw = (100.0 g)(4.184 J/g∙oC)(-51.8oC) = -21673.12 J = -21670 J - q water = + q ice -21670 J = +21.67 kJ ∆Hmelt= You need the moles of ice: 165.7 mL – 100.0 mL = 65.7 mL 65.7 mL x

43. Calculating the Molar Enthalpy of Melting Problems 1. A student adds 7 ice cubes to 101.6 mL of water 38.4oC. When the ice melted, the temperature of the water is 0.0oC. The final volume of water is 150.3 mL. Calculate the molar enthalpy of melting of water. = 6.04 kJ/mol 2. A student takes 225.4 mL of water at 52.6oC and places it in a calorimeter. Ice is added to the water until it reaches 0.0oC. At that point, all the ice is removed and the water is re-measured in a graduated cylinder. The final volume of water is 371.2 mL. Calculate the molar enthalpy of melting. = 6.13kJ/mol

44. Dissolving A Substance • When a measured amount of water is placed in a calorimeter and a solid substance is dissolved in it... • Energy will travel from the water to the solid or from the solid to the water. Exothermic Dissolving • the heat the water gains is equal to the enthalpy the process of dissolving releases. Endothermic Dissolving • the heat the water loses is equal to the enthalpy the process of dissolving absorbed.

45.  Exothermic Dissolving +qwater = -qsoln  Endothermic Dissolving -qwater = +qsoln Use stoichiometry – using the mass of the substance dissolved – to determine the molar enthalpy of solution for the substance.

46. When a 4.25 g sample of ammonium nitrate dissolves in 60.0 g of water in a calorimeter, the temperature drops from 22.0oC to 16.9oC. Calculate the ∆Hsoln for ammonium nitrate. Water q=mc∆T = (60.0g)(4.184J/goC)(-5.1oC) = -1280 J - qw=qsoln -1280 J = 1.280 kJ ∆Hsoln= kJ = 1.280 kJ = 24 kJ/mol mol 0.05309 mol Ammonia Nitrate 4.25 g x 1 mol NH4NO3 = 0.05309 mol 80.06 g NH4NO3

47. Calculating the Molar Enthalpy of Solution 1. When a 1.80 g sample of magnesium hydroxide dissolves in 108.9 g of water in a calorimeter, the temperature rises from 22.0oC to 83.9oC. Calculate the ∆Hsoln for magnesium hydroxide. = -914 kJ/mol 2. A 9.96 g sample of lithium chloride was dissolved in 101.2 mL of water at 22.2oC. When the salt is dissolved, the temp. of the soln was 41.3oC. Calculate the molar enthalpy of solution for lithium chloride. = -34.4 kJ/mol • A student dissolves 1.96 g of NaOH in 99.7 mL of water at 23.4oC. After the NaOH dissolves, the temp. of the water rises to 28.7oC. What is the molar enthalpy of solution? = -45 kJ/mol

48. Calculating Molar Enthalpy of Rxn • Coffee-cup calorimetry can be used to study changes in dilute aqueous solutions. • The water in the calorimeter absorbs (or provides) the energy that is released (or absorbed) by the chemical rxn.  +qsoln = -qrxnor -qsoln = +qrxn Same calculations as melting and solution... We just need to determine the limiting reagent when calculating the molar enthalpy.

49. Sample Problem Page 665 • Copper sulfate, CuSO4, reacts with sodium hydroxide, NaOH, in a double displacement reaction. A precipitate, Cu(OH)2, is formed. • CuSO4 + 2 NaOH→ Cu(OH)2 + Na2SO4 50.0 mL of 0.300 mol/L CuSO4 reacts with 51.8 mL of 0.600 mol/L NaOH. The initial temperature of both solutions is 21.4oC. After mixing the solutions in a coffee-cup calorimeter, the highest temperature reached is 24.6oC. Determine the enthalpy change for the reaction and then write a thermochemical equation.

50. Page 667 • Consider the following rxn: • HCl(aq) + NaOH(aq)→ NaCl(aq) + H2O (l) • The chemist uses a coffee-cup calorimeter to neutralize completely 61.1 mL of 0.543 mol/L HCl with 62.5 mL of 0.543 mol/L of NaOH. The initial temperature of both solutions is 17.8oC. After neutralization, the highest recorded temperature is 21.6oC. • 2. Consider the following rxn: • Mg (s) + 2 HCl(aq)→ MgCl2(aq) + H2 (g) • The chemist uses a coffee-cup calorimeter to react completely 0.500 g of magnesium ribbon with 100.0 mL of 1.00 mol/L of HCl. The initial temperature of the HCl is 20.4oC. After the reaction is complete, the highest recorded temperature is 40.7oC. • 3. Consider the following rxn: • HNO3(aq) + KOH (aq)→ KNO3(aq) + H2O (l) ∆Hrxn = -53.4 kJ/mol • The chemist uses a coffee-cup calorimeter to neutralize completely 55.0 mL of 1.30 mol/L of HNO3 with 58.9 mL of 1.50 mol/L of KOH. The initial temperature of both solutions is 21.4oC. What is the final temperature of the mixture?