chem 1411 ch5

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chem 1411 ch5
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  1. 5.1 The Nature of Energy
    • Thermodynamics is the study of energy and its transformations.
    • Thermochemistry is the study of the relationships between chemical reactions and energy changes involving heat.
  2. • Definitions:
    • • Energy is the capacity to do work or to transfer heat.
    • • Work is energy used to cause an object with mass to move.
    • w = F × d
  3. ch5
    • Heat is the energy used to cause the temperature of an object to increase.
    • • A force is any kind of push or pull exerted on an object.
    • • The most familiar force is the pull of gravity.

    • 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 or composition.
    • • Electrostatic energy is an example.
    • • It arises from interactions between charged particles.
    • • Potential energy can be converted into kinetic energy.
  4. ch 5
  5. Units of Energy
    • • SI unit is the joule (J).
    • • Traditionally, we use the calorie as a unit of energy.
    • • 1 cal = 4.184 J (exactly)
    • • The nutritional Calorie, Cal = 1,000 cal = 1 kcal.
  6. ch5
  7. System and Surroundings
    • • A system is the part of the universe we are interested in studying.
    • • Surroundings are the rest of the universe (i.e., the surroundings are the portions of the universe that are not involved in the system).
    • • Example: If we are interested in the interaction between hydrogen and oxygen in a cylinder, then the H2 and O2 in the cylinder form a system.
  8. ch 5
  9. Transferring Energy: Work and Heat
    • • From physics:
    • • Force is a push or pull exerted on an object.
    • • Work is the energy used to move an object against a force.
    • w = F × d
  10. ch 5
  11. Heat is the energy transferred from a hotter object to a colder one.
    • • Energy is the capacity to do work or to transfer heat.
    • • Energy in J (and kJ) to energy of thermodynamic calculations throughout Chapters 19, 14
    • (Arrhenius equation in section 14.5), and Chapter 20 (Gibbs free energy vs. cell potential in
    • section 20.5).
  12. 5.2 The First Law of Thermodynamics
    • The first law of thermodynamics states that energy cannot be created or destroyed.
    • The first law of thermodynamics is the law of conservation of energy.
  13. • That is, the energy of system + surroundings is constant.
    • • Thus, any energy transferred from a system must be transferred to the surroundings (and vice
    • versa).
  14. ch 5
  15. Internal Energy2
    • • The total energy, E, of a system is called the internal energy.
    • • It is the sum of all the kinetic and potential energies of all components of the system.
    • • Absolute internal energy cannot be measured, only changes in internal energy.
    • • Change in internal energy: ΔE = Efinal – Einitial.
    • • Thermodynamic quantities such as ΔE always have three parts:
    • • a number,
    • • a unit, and
    • • a sign that gives direction.
  16. ch 5
  17. Example: A mixture of H2(g) and O2(g) has a higher internal energy than H2O(g).
    • • Going from H2(g) and O2(g) to H2O(g) results in a negative change in internal energy,
    • indicating that the system has lost energy to the surroundings:
    • H2(g) + O2(g) 􀃆 2H2O(g) ΔE < 0
    • • Going from H2O(g) to H2(g) and O2(g) results in a positive change in internal energy,
    • indicating that the system has gained energy from the surroundings:
    • 2H2O 􀃆 H2(g) + O2(g) ΔE > 0
  18. ch 5
  19. Relating ΔE to Heat and Work3
    • • From the first law of thermodynamics:
    • • When a system undergoes a physical or chemical change, the change in internal energy is given
    • by the heat added to or liberated from the system plus the work done on or by the system:
    • ΔE = q + w
  20. ch 5
  21. Endothermic and Exothermic Processes
    • • An endothermic process is one that absorbs heat from the surroundings.
    • • An endothermic reaction feels cold.
    • • An exothermic process is one that transfers heat to the surroundings.
    • • An exothermic reaction feels hot.
  22. ch 5
  23. State Functions
    • • A state function depends only on the initial and final states of a system.
    • • Example: The altitude difference between Denver and Chicago does not depend on whether you
    • fly or drive, only on the elevation of the two cities above sea level.
    • • Similarly, the internal energy of 50 g of H2O(l) at 25 °C does not depend on whether we cool 50 g of H2O(l) from 100 °C to 25 °C or heat 50 g of H2O(l) at 0 °C to 25 °C.
    • • A state function does not depend on how the internal energy is used.
    • • Example: A battery in a flashlight can be discharged by producing heat and light. The same
    • battery in a toy car is used to produce heat and work. The change in internal energy of the battery
    • is the same in both cases.
  24. ch 5
  25. 5.3 Enthalpy
    • Chemical and physical changes that occur around us occur under essentially constant pressure of Earth’s atmosphere.• Changes may be accompanied by work done by or on the system
  26. ch 5
  27. Changes may involve the release or absorption of heat.
    • • We will focus much of our discussion on what we can learn from measurement of heat flow.
    • • Enthalpy (H) is defined as the internal energy (E) plus the product of the pressure and volume of the system.
    • H = E + PV
    • • Enthalpy is useful for investigating heat flow in events that occur under constant pressure.
    • • Again, we can only measure the change in enthalpy, ΔH.
    • • Mathematically,
    • ΔH = Hfinal – Hinitial = ΔE + PΔV
    • w = –PΔV; ΔE = q + wΔH = ΔE + PΔV = (qp + w) – w = qp
  28. ch 5
  29. • For most reactions PΔV is small thus ΔH = ΔE
    • • Heat transferred from surroundings to the system has a positive enthalpy (i.e., ΔH > 0 for an
    • endothermic reaction).
    • • Heat transferred from the system to the surroundings has a negative enthalpy (i.e., ΔH < 0 for an
    • exothermic reaction).
    • • Enthalpy is a state function.
    • • If the reaction is carried out under constant pressure,
    • • ΔE = qp – PΔV, or
    • • qp = ΔH = ΔE + PΔV
    • • and ΔE = ΔH – PΔV
  30. ch 5
  31. 5.4 Enthalpies of Reaction
    • • For a reaction, ΔHrxn = Hproducts – Hreactants.
    • • The enthalpy change that accompanies a reaction is called the enthalpy of reaction or heat of
    • reaction (ΔHrxn).
    • • Consider the thermochemical equation for the production of water:
    • 2H2(g) + O2(g) 􀃆 2H2O(g) ΔH = –483.6 kJ
    • • The equation tells us that 483.6 kJ of energy are released to the surroundings when water is
    • formed.
    • • ΔH noted at the end of the balanced equation depends on the number of moles of reactants and
    • products associated with the ΔH value.
    • • These equations are called thermochemical equations.
    • • Enthalpy diagrams are used to represent enthalpy changes associated with a reaction.
    • • In the enthalpy diagram for the combustion of H2(g), the reactants, 2H2(g) + O2(g), have a higherenthalpy than the products 2H2O(g); this reaction is exothermic
  32. ch 5
  33. • Enthalpy is an extensive property.
    • • Therefore, the magnitude of enthalpy is directly proportional to the amount of reactant consumed.
    • • Example: If one mol of CH4 is burned in oxygen to produce CO2 and water, 890 kJ of heat is
    • released to the surroundings. If two mol of CH4 is burned, then 1780 kJ of heat is released.
    • • The sign of ΔH depends on the direction of the reaction.
    • • The enthalpy change for a reaction is equal in magnitude but opposite in sign to ΔH for the
    • reverse reaction.
    • • Example: CH4(g) + 2O2(g) 􀃆 CO2(g) + 2H2O(l) ΔH = –890 kJ,
    • • But CO2(g) + 2H2O(l) 􀃆 CH4(g) + 2O2(g) ΔH = +890 kJ.
    • • Enthalpy change depends on state of the products and reactants.
    • • 2H2O(g) 􀃆 2H2O(l) ΔH = –88 kJ
  34. ch 5
    Calorimetry is a measurement of heat flow.• A calorimeter is an apparatus that measures heat flow

    • Heat Capacity and Specific Heat
    • • Heat capacity is the amount of energy required to raise the temperature of an object by 1 °C.
    • • Molar heat capacity is the heat capacity of 1 mol of a substance.
    • • Specific heat, or specific heat capacity, is the heat capacity of 1 g of a substance.
    • • Heat, q = (specific heat) × (grams of substance) × ΔT.
    • • Be careful with the sign of q.
  35. ch 5
  36. Constant-Pressure Calorimetry
    • • The most common technique is to use atmospheric pressure as the constant pressure.
    • • Recall ΔH = qp.
    • • The easiest method is to use a coffee cup calorimeter.
    • qsoln = (specific heat of solution) × (grams of solution) × ΔT = –qrxn
    • • For dilute aqueous solutions, the specific heat of the solution will be close to that of pure water.
  37. ch 5
  38. Bomb Calorimetry (Constant-Volume Calorimetry)
    • • Reactions can be carried out under conditions of constant volume instead of constant pressure.
    • • Constant volume calorimetry is carried out in a bomb calorimeter.
    • • The most common type of reaction studied under these conditions is combustion.
    • • If we know the heat capacity of the calorimeter, Ccal, then the heat of reaction,
    • qrxn = –Ccal × ΔT.• Since the reaction is carried out under constant volume, q relates to ΔE
  39. ch 5
  40. 5.6 Hess’s Law
    • Hess’s Law: If a reaction is carried out in a series of steps, ΔH for the reaction is the sum of ΔH for each of the steps.
  41. ch 5
  42. 5.7 Enthalpies of Formation
    • • If a compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, ΔHf.
    • • Standard state (standard conditions) refer to the substance at:
    • • 1 atm and 25 °C (298 K).
    • • Standard enthalpy, ΔH°, is the enthalpy measured when everything is in its standard state.• Standard enthalpy of formation of a compound, ΔH°f, is the enthalpy change for the formation of 1mol of compound with all substances in their standard states
  43. ch 5
  44. ΔH°f for selected substances are tabulated in Appendix C.
    • • A large majority of ΔH°f values tabulated in Appendix C are negative, meaning that most
    • formation reactions are exothermic.
    • • If there is more than one state for a substance under standard conditions, the more stable one is used.
    • Example: When dealing with carbon we use graphite because graphite is more stable than diamond or C60.
    • • The standard enthalpy of formation of the most stable form of an element is zero.
  45. ch 5
  46. Using Enthalpies of Formation to Calculate Enthalpies of Reaction
    • • Use Hess’s law!
    • • Example: Calculate ΔH for
    • C3H8(g) + 5O2(g) 􀃆 3CO2(g) + 4H2O(l)
    • • We start with the reactants, decompose them into elements, then rearrange the elements to form
    • products. The overall enthalpy change is the sum of the enthalpy changes for each step.
    • • Decomposing into elements (note O2 is already elemental, so we concern ourselves with C3H8):
    • C3H8(g) 􀃆 3C(s) + 4H2(g) ΔH1 = –ΔH°f [C3H8(g)]
    • • Next we form CO2 and H2O from their elements:
    • 3C(s) + 3O2(g) 􀃆 3CO2(g) ΔH2 = 3 ΔH°f [CO2(g)]
    • 4H2(g) + 2O2(g) 􀃆 4H2O(l) ΔH3 = 4 ΔH°f [H2O(l)]
    • • We look up the values and add:
    • ΔH°rxn = –1(–103.85 kJ) + 3(–393.5 kJ) + 4(–285.8 kJ) = –2220 kJ
    • • In general:
    • ΔH°rxn = n ΔH°f (products) – m ΔH°f (reactants)
    • • Where n and m are the stoichiometric coefficients.
  47. ch 5
  48. 5.8 Foods and Fuels
    • • Fuel value is the energy released when 1 g of substance is burned.
    • • The fuel value of any food or fuel is a positive value that must be measured by calorimetry.
    • Foods
    • • Fuel value is usually measured in Calories (1 nutritional Calorie, 1 Cal = 1000 cal).
    • • Most energy in our bodies comes from the oxidation of carbohydrates and fats.
    • • In the intestines carbohydrates are converted into glucose, C6H12O6, or blood sugar.
    • • In the cells glucose reacts with O2 in a series of steps, which produce CO2, H2O, and energy.
    • C6H12O6(s) + 6O2(g) 􀃆 6CO2(g) + 6H2O(l) ΔH°= –2803 kJ
    • • Fats, for example tristearin, react with O2 as follows:
    • 2C57H110O6(s) + 163O2(g) 􀀅 114CO2(g) + 110H2O(l) ΔH° = –75,250 kJ
    • • Fats contain more energy than carbohydrates. Fats are not water soluble. Therefore, fats are good for energy storage.
  49. ch 5
  50. Fuels
    • • In 2008 the United States consumed about 1.05 × 1017 kJ/year (9.4 × 105 kJ of fuel per person per day).
    • • Most of this energy comes from petroleum and natural gas.
    • • The remainder of the energy comes from coal, nuclear and hydroelectric sources.
    • • Coal, petroleum, and natural gas are fossil fuels. They are not renewable.
    • • Natural gas consists largely of carbon and hydrogen. Compounds such as CH4, C2H6, C3H8 and
    • C4H10 are typical constituents.
    • • Petroleum is a liquid consisting of hundreds of compounds. Impurities include S, N, and O
    • compounds.
    • • Coal contains high molecular weight compounds of C and H. In addition compounds
    • containing S, O, and N are present as impurities that form air pollutants when burned in air.
    • Other Energy Sources45
    • • Nuclear energy: energy released in splitting or fusion of nuclei of atoms.
    • • It is used to produce about 21% of the electric power in the US.
    • • Fossil fuels and nuclear energy are nonrenewable sources of energy.
    • • Renewable energy sources include:
  51. ch 5
  52. • Solar energy
    • • Wind energy
    • • Geothermal energy
    • • Hydroelectric energy
    • • Biomass energy
    • • These are virtually inexhaustible and will become increasingly important as fossil fuels are
    • depleted.

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