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Energy and Chemical Change

Energy and Chemical Change

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Energy and Chemical Change

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  1. Energy and Chemical Change Chapter 16

  2. Energy Section 16.1

  3. Energy • The ability to do work or produce heat. • It exists in two basic forms, potential energy and kinetic energy: • Potential: energy due to the composition or position of an object • Kinetic: energy of motion • Chemical potential energy is the energy stored in a substance because of its composition.

  4. Energy flow • Exothermic: when heat flows from the object to its surroundings; heat is released • Endothermic: when heat flows from the surroundings to the object; heat is absorbed • Was making ice cream exothermic or endothermic?

  5. Law of conservation of energy • In any chemical reaction or physical process, energy can be converted from one form to another, but is neither created nor destroyed.

  6. Heat (q) • Energy that is in the process of flowing from a warmer object to a cooler object. • The amount of heat required to raise one gram of water by one degree Celsius is defined as a calorie (cal). • 1 Calorie = 1 kcal = 1000 cal • The SI unit of heat and energy is the joule (J).

  7. Converting energy units • 1 J = 0.2390 cal 1 cal = 4.184 J • A granola bar contains 142 Calories. Convert this to joules.

  8. Specific heat • The amount of heat required to raise the temperature of one gram of that substance by one degree Celsius. • q = m · c ·ΔT • q = the heat absorbed or released • m = the mass of the sample in g • c = the specific heat of the substance • ΔT = the change in temperature

  9. Calculations with specific heat • The temperature of a sample of iron with a mass of 10.0g changed from 50.4˚C to 25.0˚C with the release of 114J heat. What is the specific heat of iron? • If the temperature of 34.4g of ethanol increases from 25.0˚C to 78.8˚C, how much heat has been absorbed by the ethanol? The specific heat of ethanol is 2.44J/(g·˚C)

  10. Homework • After learning about section 16.1, you should be able to do “Heat And Its Measurement” • On the back: • Pick a favorite food that has a nutrition label • Write down what the food is and the Calorie content for one serving • Convert this to cal, J, kJ

  11. Problem Solving Lab (p. 503) • Create the graph and answer questions 1-3

  12. Heat in Chemical Reactions and Processes Section 16.2

  13. Thermochemistry • The study of heat changes that accompanies chemical reactions and phase changes. • System: the specific part of the universe that is being studied (chemical reaction) • Surroundings: everything else • Universe = system + surroundings

  14. Measuring heat • A calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

  15. Bomb Calorimeter

  16. Enthalpy (H) • The heat content of a system at constant pressure • Enthalpy of reaction(ΔHrxn) is the change in enthalpy for a reaction • ΔHrxn= Hfinal – Hinitial • ΔHrxn= Hproducts - Hreactants • If ΔH is positive, the reaction is endothermic • If ΔH is negative, the reaction is exothermic

  17. Think-Pair-Share • Compare energy changes in chemical reactions to profits and losses in a business. • Each month the business has receipts (positive dollar amounts) and expenses (negative dollar amounts). • If the total receipts exceed expenditures, a positive dollar amount or profit occurs. • If expenditures are greater than receipts, money is lost and a negative dollar amount results.

  18. Practice Problem • A 75.0g sample of a metal is placed in boiling water until its temperature is 100.0˚C. A calorimeter contains 100.00g of water at a temperature of 24.4˚C. The metal sample is removed from the boiling water and immediately placed in the water in the calorimeter. The final temperature of the metal and water in the calorimeter is 34.9˚C. Assuming that the calorimeter provides perfect insulation, what is the specific heat of the metal?

  19. Homework • P. 500 (14-18) • P. 525 (79-81)

  20. Thermochemical Equations Section 16.3

  21. Thermochemical Equation • A balanced equation that includes the physical states and energy change. • Example: • 4Fe(s) + 3O2(g) → 2Fe2O3(s) ΔH= -1625 kJ

  22. Enthalpy (heat) of combustion (ΔHcomb) • The enthalpy change for the complete burning of one mole of a substance • Example: • C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l) ΔHcomb=-2808kJ

  23. Changes of State • Molar enthalpy (heat) of vaporization (ΔHvap) • H2O(l) → H2O(g) ΔHvap= 40.7kJ • H2O(g) → H2O(l) ΔHcond= -40.7kJ • Molar enthalpy (heat) of fusion (ΔHfus) • H2O(s) → H2O(l) ΔHfus= 6.01kJ • H2O(l) → H2O(s) ΔHsolid= 6.01kJ • What would be the molar enthalpy of sublimation?

  24. Practice Problems • Calculate the heat required to melt 25.7 g of solid methanol at its melting point. ΔHfus=3.22 kJ/mol • How much heat is evolved when 275 g of ammonia gas condenses to a liquid at its boiling point? ΔHvap=23.3 kJ/mol • What mass of methane must be burned in order to liberate 12,880 kJ of heat? ΔHcomb=-891 kJ/mol

  25. Homework • P. 505 (26-27) • P. 525 (82-84)

  26. Calculating Enthalpy Change Section 16.4

  27. Hess’s Law • Read section on p. 506 • Add the following equations: • A + B → C • C + D → E + B

  28. Example • Ex) 2S(s) + 3O2(g) → 2SO3(g) ΔH= ? • S(s) + O2(g) → SO2(g) ΔH= -297kJ • 2SO3(g) → 2SO2(g) + O2(g) ΔH= 198kJ

  29. Reaction Spontaneity 16.5

  30. Spontaneous Process • A spontaneous process is a physical or chemical change that occurs with no outside intervention. • Ex. Iron forming rust, combustion of methane, melting of ice cream • A non-spontaneous change is a change that occurs only when driven e.g. forcing electric current through a metal block to heat it • What makes a reaction spontaneous?

  31. What determines if a reaction is spontaneous? • Entropy: Measure of the disorder or randomness of the particles that make up a system. • Spontaneous processes always proceed in such a way that the entropy of the universe increases. (ΔS) ΔSuniverse= ΔSsystem + ΔSsurroundings • Larger entropy value, larger degree of randomness. Law of Disorder (second law of thermodynamics): Spontaneous processes always proceed in such a way that the entropy of the universe increases.

  32. Law of Disorder (Second Law of Thermodynamics) • Law of disorder: Spontaneous processes always proceed in such a way that the entropy of the universe increases. • Entropy gas > entropy liquid > entropy solid • Dissolving of a gas in a solvent always results in a decrease in entropy • An increase in temperature results in an increase in entropy. • Entropy increases when the number of product particles is greater than the number of reactant particles. • 2SO3(g)2SO2(g) + O2(g) ΔSsystem > 0

  33. Entropy, the Universe and Free Energy • For any spontaneous reaction • ΔSuniverse > 0 • ΔSuniverse is positive when, • 1. the reaction is exothermic • 2. The entropy of the system increases, so ΔSsystem is positive Gibbs Free Energy (G) or Free energy: -the energy that is available to do work. ΔGsystem= ΔHsystem -TΔSsystem

  34. Practice Problem • For a process, ΔHsystem= 145 kJ and ΔSsystem=322 J/K. Is the process spontaneous at 382 K? • Answer: 22,000 J