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Thermochemistry

Thermochemistry

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Thermochemistry

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  1. Thermochemistry Chapter 8

  2. 900C 900C Energy Changes in Chemical Reactions Which one has more thermal energy? Coffee cup? Bathtub?

  3. 900C 400C Energy Changes in Chemical Reactions Which one has more thermal energy? What additional information is needed? How does it relate to heat?

  4. Energy Changes in Chemical Reactions Temperature = Thermal Energy Heat: the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. Measure of the average kinetic energy

  5. Energy • Energy:the capacity to do work • Radiant energy comes from the sun and is earth’s primary energy source • Thermal energy is the energy associated with the random motion of atoms and molecules • Chemical energy is the energy stored within the bonds of chemical substances • Nuclear energy is the energy stored within the collection of neutrons and protons in the atom • Potential energy is the energy available by virtue of an object’s position. E = mgh • Kinetic Energy: energy of motion E = ½ mu2

  6. Energy Changes SURROUNDINGS SYSTEM The system is the specific part of the universe that is of interest in the study. closed isolated open energy nothing Exchange: mass & energy

  7. Energy Changes Detected by generation of Heat: energy that flows from a body at higher temperature to a body of lower temperature Represented by the letter q. Work: can be done on the system by surroundings or vice versa. Examples: hammering a nail into a wooden block; gasoline burning in an engine; tossing a baseball (or any other ball) T ∝ KE KE = ½ mv2 Units of Energy: Joules calories : 1 calorie = 4.18 J Calories = food calorie = 1000 calories

  8. Temperature Degree of hotness or coldness. Measures the average KE of the molecules. Thermometers operate on the following principles: 1. expansion of liquids 2. expansion of gas 3. radiation properties of substances 4. electrical properties (digital thermometers) Temperature scales: Kelvin Celsius Fahrenheit

  9. Velocity or speed Temperature The higher the temperature, the greater the average speed of the molecules. T µ KE KE =½ mv2 T µv

  10. Energy Changes State Properties: depend only on the initial and final states of the system as it is defined by volume, temperature, pressure, number of moles.

  11. Heat: Endothermic/Exothermic 2H2(g) + O2(g) 2H2O (l) + energy H2O (g) H2O (l) + energy energy + 2HgO (s) 2Hg (l) + O2(g) energy + H2O (s) H2O (l) Exothermic process:is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process: is any process in which heat is supplied to the system from the surroundings.

  12. Enthalpy: Enthalpy Diagrams ΔHrxn= ΔHproducts – ΔHreactants

  13. Enthalpy Enthalpy: a thermodynamic quantity used to describe heat changes at constant pressure(most reactions occur at constant pressure). Heat absorbed or released at const. P during chemical reaction. Equation: ΔHreaction = ΔHproducts - ΔHreactants

  14. Enthalpy: Examples Indicate the sign of enthalpy change in the following processes carried out under atmospheric pressure, and indicate whether the process is exothermic or endothermic: • An ice cube melts • 1 g of butane (C4H10) is combusted in sufficient oxygen to give complete combustion to CO2 and H2O • A bowling ball is dropped from a height of 8 ft into a bucket of sand.

  15. Thermochemical Equations Find the differences between the following equations: • CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = -890.4 kJ • CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔH = -802.4 kJ Thermochemical Equations: describe the reaction (includes the state of the substances) and energy changes that occur during the chemical reaction.

  16. Thermochemical Equations H2O (s)→ H2O (l) DH = 6.01 kJ Is DH negative or positive? System absorbs heat Endothermic DH > 0 6.01 kJ are absorbed for every 1 mole of ice that melts at 00C and 1 atm. 6.3

  17. Thermochemical Equations DH = -890.4 kJ CH4(g) + 2O2(g)→ CO2(g) + 2H2O (l) Is DH negative or positive? System gives off heat Exothermic DH < 0 890.4 kJ are released for every 1 mole of methane that is combusted at 250C and 1 atm. 6.3

  18. Rules of Thermochemistry ΔH directly proportional to the amount of reactants or products (stoichiometric relationship). ΔH forward = - ΔH reverse • The value of ΔH for the reaction is the same whether it occurs in one step or in series of steps. (Hess’ Law- AP Chem) • The physical states of all reactants and products must be specified in a thermochemical equation.

  19. Rule #1: ΔH is directly proportional to the amount of reactants or products (stoichiometric relationship). C6H6(l) → C6H6(g) ΔH = 7.36 kCal 2C6H6(l) ΔH = ½ C6H6(l) ΔH = Given the thermochemical equation C (s) + O2(g) → CO2(g) ΔH = -94.1 kCal Calculate (a) ΔH when 1.00 g of C burns. (b) ΔH when 1.8 g of CO2 is produced.

  20. Thermochemical Equations: Rules How much heat is evolved when 266 g of white phosphorus (P4) burn in air? P4(s) + 5O2(g) P4O10(s)DH = -3013 kJ x -3013 kJ 1 mol P4 x 1 mol P4 123.9 g P4 [ -6470 kJ] 266 g P4

  21. Example Given the thermochemical equation: C8H18(l) + 12.5 O2(g) → 8CO2(g) + 9H2O(l) ΔH = -1308 kCal How many grams of octane have to be burned to evolve 1 kilocalorie of heat? Fractions are allowed

  22. Thermochemical Equations: Reverse Reaction When a reaction is reversed, the magnitude of ΔH remains the same, but its sign changes. H2O(s) → H2O(l) ΔH = 1.44 kCal CaCO3(s) → CaO(s) + CO2(g) ΔH = 42.5 kCal C10H8(s) + 12 O2(g) → 10CO2(g) + H2O(l) ΔH = -1232.5 kCal

  23. Thermochemical Equations: Rules • The physical states of all reactants and products must be specified in thermochemical equations. H2O(s)→ H2O(l)ΔH = 6.00 kJ H2O(l) → H2O(g)ΔH = 44.0 kJ • What will be the ΔH when 1 mole of ice at 0ºC is changed into one mole of steam at the boiling point (100 0ºC)?

  24. Hess’ Law • Heat of reaction is an algebraic sum of heats of chemical reactions that when added give the required equation. • Allow calculations of ΔH for reactions that cannot be easily experimentally determined. The value of the reaction is the same whether it occurs in one step or in series of steps. • Example: • H2(g) + ½ O2(g) → H2O(g) ΔH = -57.8 kCal • H2(g) + ½ O2(g) → H2O(l) ΔH = -68.3 kcal • Needed: • H2O(g) → H2O(l) , cannot be easily measured

  25. Calorimetry

  26. Calorimetry Heat thermal energy transferred from a hot object to a cold object. The heat transferred is proportional to the mass specific heatcapacity temperature change Heat has the symbol q and is calculated using … q = mcΔT

  27. Calorimetry • Quantity of heat depends on: 1. The mass of the substance 2. Mass of the calorimeter 3. Specific heat of the substance and calorimeter. Heat capacity, C: quantity of heat needed to raise the temperature of a given substance by 1ºC. Extensive or Intensive? ___________ Specific heat, c or cp: heat needed to raise the temperature of 1.0 g of substance by 1ºC. Extensive or Intensive? ___________

  28. specific heat capacity Quantity of heat temperature change mass q = m c DT Law of Conservation of Energy

  29. Examples • 1. How much heat is needed to raise the temperature of 25.6 grams of water from 20.0 C to 50.0 C? The specific heat of water is equal to 4.18 J/(g ºC) • Answer: 3210 J • How much heat is lost by the system when a solid Al ingot with a mass of 4110 g cools from 60.0 ºC to 25.0 ºC? • (-2.35 x 106 J) • 3. What is the final temperature of 27.0 grams of liquid water, initially at 0 ºC, after it absorbs 700.0 J of energy? • (6.20 ºC)

  30. Examples 4. A 40.0g of glass was heated from 0.0 ºC to 41.0 ºC and was found to have absorbed 32 J of heat. (a) What is the speicfic heat of this type of glass? (b) How much heat did the same glass sample gain when it was heated from 41.0 3. A 40.0g of glass was heated from 41.0 ºC to 70.0 ºC? 0.20 J/(g ºC); 23 J

  31. Examples 5. When 2.80 g CaCl2 dissolves in 100.0 g of water, the temperature of the awater risaes from 20.5 ºC to 25.4 ºC. Assume that all the heat is being absorbed by the water (cp = 4.18 J/(g ºC). (a) Write a balance equation of the solution process. (b) What is q for the solution process? (c) Is the process endo or exo? (d) How much heat is absorbed by the water if 1.00 mole of CaCl2 dissolved?

  32. Phase Changes Phase– homogeneous part of system in contact with other parts of system, separated by well-defined boundary e.g., ice in water, subliming dry ice, evaporating isopropanol

  33. Phase Changes Hess’ Law for phase changes.

  34. Phase Changes and Heating Curves

  35. Is heat is absorbed or released during a phase change? How could you measure the heat absorbed or released as substances change phase?

  36. Consider ice melting in water. • Does the temperature of the water change? • Is the water absorbing or releasing heat? • Does ice absorb heat or release heat as it melts?

  37. Consider ice melting in water. • Does the temperature of the water change? • Is the water absorbing or releasing heat? • Does ice absorb heat or release heat as it melts? No Releasing heat Absorb heat

  38. Consider ice melting in water. The word fusion means “melting”. How could you design an experiment to measure the heat of fusion of ice?

  39. Phase Changes and Heating Curves

  40. Is heat is absorbed or released during a phase change? How could you measure the heat absorbed or released as substances change phase?

  41. Consider ice melting in water. • Does the temperature of the water change? • Is the water absorbing or releasing heat? • Does ice absorb heat or release heat as it melts?

  42. Consider ice melting in water. • Does the temperature of the water change? • Is the water absorbing or releasing heat? • Does ice absorb heat or release heat as it melts? No Releasing heat Absorb heat

  43. Consider ice melting in water. The word fusion means “melting”. How could you design an experiment to measure the heat of fusion of ice?

  44. Consider ice melting in water. You could measure the heat lost by some water as it cools. Ice That should equal the heat gained by the ice as it melts.

  45. We now know that heat is absorbed or released during a phase change. Heat is absorbed as solids melt, or liquids vaporize.

  46. Heat is absorbed by the ice. Ice And melts.

  47. Heat is absorbed by the ice. One gram of ice at 0C absorbs 334 J as it melts to form water at 0C. … making liquidwater

  48. Heat is released by the water as it freezes. 334 joules is released when one gram of water freezes at 0C. Ice water

  49. Ice absorbs 334 J per gram as it melts at 0C Ice Water releases 334 J per gram as it freezes at 0C

  50. Heat is absorbed by the water as it vaporizes. Hotplate