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Thermochemistry

Thermochemistry. The Nature of Energy. Kinetic Energy and Potential Energy Kinetic energy is the energy of motion: Potential energy is the energy an object possesses by virtue of its position. Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.

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Thermochemistry

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  1. Thermochemistry

  2. The Nature of Energy • Kinetic Energy and Potential Energy • Kinetic energy is the energy of motion: • Potential energy is the energy an object possesses by virtue of its position. • Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.

  3. The Nature of Energy • Units of Energy • SI Unit for energy is the joule, J: • sometimes the calorie is used instead of the joule: • 1 cal = 4.184 J (exactly) • A nutritional Calorie: • 1 Cal = 1000 cal = 1 kcal

  4. Thermochemistry Terminology • System: part of the universe we are interested in. • Surrounding: the rest of the universe. • Boundary: between system & surrounding. • Exothermic: energy released by system to surrounding. • Endothermic: energy absorbed by system from surr. • • Work ( w ): product of force applied to an object over a distance. • Heat ( q ): transfer of energy between two objects

  5. The First Law of Thermodynamics • Internal Energy • Internal Energy: total energy of a system. • Involves translational, rotational, vibrational motions. • Cannot measure absolute internal energy. • Change in internal energy, HyperChem

  6. The First Law of Thermodynamics • Relating DE to Heat(q) and Work(w) • Energy cannot be created or destroyed. • Energy of (system + surroundings) is constant. • Any energy transferred from a system must be transferred to the surroundings (and vice versa). • From the first law of thermodynamics:

  7. The First Law of Thermodynamics

  8. First Law of Thermodynamics • Calculate the energy change for a system undergoing an exothermic process in which 15.4 kJ of heat flows and where 6.3 kJ of work is done on the system. DE = q + w

  9. The First Law of Thermodynamics • Exothermic and Endothermic Processes • Endothermic: absorbs heat from the surroundings. • An endothermic reaction feels cold. • Exothermic: transfers heat to the surroundings. • An exothermic reaction feels hot.

  10. Endothermic Reaction Ba(OH)2•8H2O(s) + 2 NH4SCN(s)  Ba(SCN)2(s) + 2 NH3(g) + 10 H2O(l)

  11. The First Law of Thermodynamics • State Functions • State function: depends only on the initial and final states of system, not on how the internal energy is used.

  12. Enthalpy • Chemical reactions can absorb or release heat. • However, they also have the ability to do work. • For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work. • Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) • The work performed by the above reaction is called pressure-volume work. • When the pressure is constant,

  13. Enthalpy

  14. Enthalpy • Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. • Enthalpy is a state function. • If the process occurs at constant pressure,

  15. Enthalpy • Since we know that • We can write • When DH is positive, the system gains heat from the surroundings. • When DH is negative, the surroundings gain heat from the system.

  16. Enthalpy => Heat of Reaction

  17. Enthalpies of Reaction • For a reaction: • Enthalpy is an extensive property (magnitude DH is directly proportional to amount): • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ • 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) DH = -1604 kJ

  18. Enthalpies of Reaction • When we reverse a reaction, we change the sign of DH: • CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ • Change in enthalpy depends on state: • H2O(g)  H2O(l) DH = -44 kJ

  19. Calorimetry • Heat Capacity and Specific Heat • Calorimetry = measurement of heat flow. • Calorimeter = apparatus that measures heat flow. • Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). • Molar heat capacity = heat capacity of 1 mol of a substance. • Specific heat = specific heat capacity = heat capacity of 1 g of a substance.

  20. Table 5.2: Specific Heats (S) of Some Substances at 298 K

  21. If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is the final temperature of the aluminum if its initial temperature is 5.0oC? CyberChem - Pizza

  22. Calorimetry • Constant Pressure Calorimetry • Atmospheric pressure is constant!

  23. Calorimetry Constant Pressure Calorimetry

  24. Calorimetry

  25. Calorimetry Examples CyberChem - Pizza • In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, 1.500 g of Mg(s) was combined with 125.0 mL of 1.0 M HCl. The initial temperature was 25.0oC and the final temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction in terms of the number of moles of Mg(s) used. Ans: (a) –25.0 kJ (b) –406 kJ/mol • 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup calorimeter. After the mixing process, the thermometer reading was at 31.9oC. Calculate the energy involved in the reaction and the enthalpy per moles of hydrogen ions used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization for strong acid/base reactions]

  26. Hess’s Law • Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step. • For example: • CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ • 2H2O(g)  2H2O(l) H= - 88 kJ • CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ

  27. Another Example of Hess’s Law Given: C(s) + ½ O2(g)  CO(g) DH = -110.5 kJ CO2(g)  CO(g) + ½ O2(g) DH = 283.0 kJ Calculate DH for: C(s) + O2(g)  CO2(g)

  28. Enthalpies of Formation • If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof . • Standard conditions (standard state): Most stable form of the substance at 1 atm and 25 oC (298 K). • Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state. • Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.

  29. Enthalpies of Formation • If there is more than one state for a substance under standard conditions, the more stable one is used. • Standard enthalpy of formation of the most stable form of an element is zero.

  30. Enthalpies of Formation

  31. Enthalpies of Formation • Using Enthalpies of Formation to Calculate Enthalpies of Reaction • For a reaction • Note: n & m are stoichiometric coefficients. • Calculate heat of reaction for the combustion of propane gas giving carbon dioxide and water. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()

  32. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O() ∆Hof (kJ/mol):

  33. Foods and Fuels • Foods • 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal. • Energy in our bodies comes from carbohydrates and fats (mostly). • Intestines: carbohydrates converted into glucose: • C6H12O6 + 6O2 6CO2 + 6H2O, DH = -2816 kJ • Fats break down as follows: • 2C57H110O6 + 163O2 114CO2 + 110H2O, DH = -75,520 kJ • Fats contain more energy; are not water soluble, so are good for energy storage.

  34. Foods and Fuels • Fuels • Fuel value = energy released when 1 g of substance is burned. • Most from petroleum and natural gas. • Remainder from coal, nuclear, and hydroelectric. • Fossil fuels are not renewable. • In 2000 the United States consumed 1.03  1017 kJ of fuel. • In 2005 the United States consumed 1.05  1017 kJ of energy. • Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g. [ gasoline ≈ 35 kJ/g ]

  35. [8.1%] [6.0%] Foods and Fuels [40.4%] [22.9%] (% for 2000) [% for 2005] [22.6%]

  36. Thermochemistry

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