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Thermochemistry

Learn about kinetic and potential energy, essential thermodynamic terminology, the first law of thermodynamics, and the concept of enthalpy. Explore state functions, calorimetry, and enthalpies of reaction in this comprehensive guide.

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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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