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Energy and Heat

Energy and Heat. Definitions. Thermochemistry: the study of the energy changes that accompany chemical reactions Energy : A property of matter describing the ability to do work Kinetic Energy: related to movement of an object

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Energy and Heat

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  1. Energy and Heat

  2. Definitions • Thermochemistry: the study of the energy changes that accompany chemical reactions • Energy: A property of matter describing the ability to do work • Kinetic Energy: related to movement of an object • Potential Energy: based on position of an object (stored energy)

  3. Review1st Law of Thermodynamics • 1st Law of Thermodynamics: The law of conservation of energy • Energy cannot be created or destroyed, but transformed from one form to another • The Kinetic Molecular Theory • Molecules are constantly moving (kinetic energy) • When heat is added, the molecules move faster and their kinetic energy increases

  4. Energy • Units: • Joule (kg*m2/s2) • 1 calorie = 4.184 J • Heat energy (q): energy transferred between substances during a physical or chemical change

  5. Temperature vs. Thermal Energy • Temperature: A measure of the average kinetic energy of the molecules in a substance (measured) • The intensity of heat • Units: oC or K • Thermal Energy: a form of kinetic energy that results from the motion of molecules (cannot be measured)

  6. Systems • Systems: Where a physical or chemical change occurs • Surroundings: the rest of the universe outside the system • Open System: Both matter and energy can freely flow into the surroundings • Closed System: Energy can flow to surroundings but matter cannot • Isolated System: Neither energy nor matter can flow to the surroundings

  7. Types of Reactions • Endothermic: When heat is transferred from the surroundings to the system • Q of the system (+) • Q of the surroundings (-) • Exothermic: When heat is transferred from the system to the surroundings • Q of the system (-) • Q of the surroundings (+)

  8. Heat Capacity (C) • The amount of heat energy required to raise the temperature of a substance (of any mass) by 1 oC • (Units: J/oC) • Specific heat capacity: Amount of heat energy required to raise the temperature of 1g of a substance by 1oC • (Units: J/(g•°C)) • Molar heat capacity: Amount of heat energy required to raise 1 mole of a substance by 1oC • Units: J/(mol•°C))

  9. Practice • You are provided with a glass of milk and a swimming pool of milk • Which will have the higher heat capacity? • Which will have the higher specific heat capacity? • Answer: • The swimming pool of milk will have the higher heat capacity, because it is dependent on the amount and there is more in the pool • The glass and swimming pool of milk have that same specific heat capacity because that is based on 1g of the substance

  10. Factors Affecting Heat Capacity • Mass (m) – the greater the number of molecules, the more heat required • Temperature change (∆T) - the greater the temperature increase, the more heat required • Type of substance – each substance differs in ability to absorb heat (c)

  11. Calculating Quantities of Heat q=mc⧋t • q = quantity of heat transferred (J or kJ) • If q is negative, the reaction was exothermic • If q is positive, the reaction was endothermic • m = mass of substance (grams) • C = specific heat capacity (J/(g•°C)) • ⧋t =t2-t1(oC)

  12. Practice When 600g of water in an electric kettle is heated from 20oC to 85oC to make a cup of tea, how much heat flows into the water?

  13. Practice When 600g of water in an electric kettle is heated from 20oC to 85oC to make a cup of tea, how much heat flows into the water? Given: m=600g C = 4.184 J/(g•°C) t1 = 20°C t2 = 85°C q=mc⧋t q=600g(4.184 J/(g•°C))(85°C-20°C) =1.63x105J or 163kJ Therefore, the amount of heat that flowed into the water was 163kJ. Unknown: q=? Note: 1000J = 1kJ

  14. Enthalpy • Impossible to measure the sum of kinetic and potential energy • Instead we measure enthalpy change • Enthalpy Change (⧋H): the energy absorbed or released to the surroundings when a system changes from reactants to products • ⧋H = q(when pressure kept constant) • Units: J or kJ

  15. Endothermic: Heat enters the system from the surroundings ( + ΔH) Exothermic: Heat leaves the system into the surroundings( - ΔH)

  16. Making and Breaking Bonds • It takes energy to break bonds and energy is released when bonds are formed • Exothermic Reactions: Less energy is required to break bonds in the reactants than is released by formation of new bonds in the products • Endothermic Reactions: More energy is required to break bonds in the reactants than is released by formation of new bonds in the products

  17. Molar Enthalpy • Molar Enthalpy: the enthalpy change per mole of a substance undergoing a change • They are specific to a reaction • ⧋Hx • x = type of change that is occurring

  18. Molar Heat of Fusion (ΔHfus) • The amount of heat necessary to melt 1.0 mole of a substance at its melting point. • Units: J/molor kJ/mol • Specific to each substance *No temperature change involved* • You add heat to the substance, but the substance remains the same temperature • All the heat is being converted into energy that is being used to break some of the intermolecular forces (to convert the substance from solid to liquid) q (J) = n (mol) x ΔHfus (J/mol)

  19. Question • 11.0 kJ are required to melt 4.2 mol of sodium (Na). What is the molar enthalpy of fusion? • q = 11.0 kJ • n = 4.2 mol • ΔHfus = ? Note: You should be able to do a problem like this if you were given the amount in grams instead of moles n=m/M q = n x ΔHfus ΔHfus= q / n ΔHfus = 11.0 kJ / 4.2 mol ΔHfus = 2.6 kJ/mol

  20. Molar Heat of Vaporization(ΔHvap) • The amount of heat necessary to boil 1.0 mole of a substance at its boiling point. • Unit: J/mol or kJ/mol • Specific to each substance *No temperature change involved* • You add heat to the substance, but the substance remains the same temperature • All the heat is being converted into energy which is being used to break all of the intermolecularforces (to convert the substance from liquid to gas) q (J) = n (mol) x ΔHvap (J/mol)

  21. Question • The ΔHvapof H2O is 40 700 J/mol. How many Joules are needed to completely boil 35.0 g of water at its boiling point? Given: • ΔHvap= 40 700 J/mol • m = 35.0 g • n = ? * Remember: n = m/M (molar mass) • M = 18.0 g/mol (calculated from atomic masses) • q = ? q = n x ΔHvap q = m/M x ΔHvap(J/mol) q = (35.0 g/18.0 g/mol) x 40 700 (J/mol) q = 7.91 x 104 J

  22. Heat energy is continuously being added to H2O. • However, the temperature does not increase at the point of fusion and vaporization Time-Temperature Graph

  23. Question: • How much energy is required to raise the temperature of 45 g of H2O from 30 °C to 150 °C?

  24. Answer: 1) (liquid) ΔT =T2 –T1 = 100 °C – 30 °C = 70 °C q1 = mc ΔT q1 = 45 gx 4.18 J/g•°Cx 70 °C q1 = 1.32 x 104 J

  25. Answer: 2) (phase change: Boiling ) Since ΔT = 0 °C q2 =nΔHvap q2 = m/M xΔHvap q2 = 45 g/18.0 g/mol x 40 800J/mol q2 = 1.02 x 105 J

  26. Answer: 3) (gas) T2 –T1 = 150 °C – 100 °C = 50 °C q3 = mc ΔT q3 = 45 gx 2.07 J/g•°Cx 50 °C q3 = 4.7 x 103 J

  27. Final Answer: qtotal = q1 + q2 + q3 = 1.2 x 105 J

  28. CALORIMETRY

  29. Calorimetry • The science of measuring the change in heat of chemical reactions or physical changes • A calorimeter: an insulated reaction vessel in which a reaction can occur and where the change in temperature of the system can be measured ∆Etotal = ∆Esystem + ∆Esurroundings = 0 ∆Esystem = -∆Esurroundings

  30. Calorimetry Assumptions • No heat is being transferred between the calorimeter and the outside environment (isolated system) • Any heat absorbed or released by the calorimeter itself is negligible (unless information is given to you in the question!) • A dilute aqueous solution is assumed to have a density and specific heat capacity equal to that of water

  31. Coffee-Cup Calorimeter • Reaction takes place in the water (aqueous) • Measure ΔT of water • Assume m solution = m H2O • Assume C solution = C of H2O • q=mcΔT • ΔHrxn = -q water • Limitations: • Cannot be used for reactions involving gases • Cannot be used for high temperature reactions

  32. Bomb Calorimeter • Sealed metal container • Temperature difference of the water is measured • ΔHrxn = -q water

  33. Problem • Hot copper was added to a calorimeter. Given the following calorimeter data, calculate the specific heat of copper. • Calorimeter: • Mass of water 500.0 g • Initial temp. of water 20.0 °C • Final temp. of water 26.2 °C • Copper sample: • Mass of copper 482.00 g • Initial temp. of copper 98.8 °C

  34. Answer: • qgained = - qlost • q = mcΔT • At temperature equilibrium, • mwcwΔTw= -(mcuccuΔTcu) • (500.0 g) (4.18 J/g•°C) (6.2 °C) = -(482.00 g) (ccu) (-72.6 °C) • ccu = 3.7 x 10-1 J/g•°C

  35. Flame Calorimetry • Measures enthalpy of combustion (ΔHcomb) • Fuel being tested is burned under a can of water, heating both the water and the can • q=mcΔT • q total = q water + q can • ΔHcomb=-q total

  36. Representing ⧋H

  37. 1. Thermochemical Equations Ex:2Na(s) + 2 H2O(l) 2NaOH(aq) + H2(g) ⧋ H = -368.6 kJ • Exothermic reactions in one direction become endothermic reactions in the reverse direction (the sign of ⧋Hois reversed)

  38. 2. Thermochemical Equations with Energy Terms “Energy is a reactant” “Energy is a product”

  39. 3. Molar Enthalpies of Reaction • The enthalpy change associated with 1 mole of a substance • The particular reactant or product must be specified

  40. 4. Potential Energy Diagrams Exothermic Endothermic

  41. Example 1: A chemist wants to determine empirically the enthalpy change for the following reaction: Mg(s) + HCl(aq)  MgCl2(aq) + H2 (g) The chemist uses a coffee-cup calorimeter to react 0.50 g of Mg ribbon with 100 mL of 1.00mol/L HCl(aq). The initial temperature of the HCl is 20.4˚C. After neutralization, the highest recorded temperature is 40.7˚C. Calculate the enthalpy change (⧋H), in kJ/mol of Mg Write the thermochemical equation Draw a potential energy diagram State any assumption you made

  42. Example 2: An aluminum-can calorimeter was used to determine the heat of combustion (kJ/mol) of isopropyl alcohol. It was found that combustion of 1.37 g of the alcohol raised the temperature of the can and water from 19.0˚C to 28.6˚C. The mass of the aluminum-can was found to be 45.73g and the mass of the aluminum-can + water was 202.46g. Hint: cAl=0.897J/g˚C Calculate the heat of combustion (kJ/mol) of isopropyl alcohol (C3H8O). Write the thermochemical equation with energy terms

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