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Now we’re into Chapter 8

Now we’re into Chapter 8. All have a pretty basic concept of what energy is… Chapter 6 is basically “Energy in Chemical Rxns.” Thermochemistry… Instinctively, we know chem. rxn’s can produce heat, after all, natural gas heats our homes. CH 4 + 2 O 2 → CO 2 + 2 H 2 O

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Now we’re into Chapter 8

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  1. Now we’re into Chapter 8 • All have a pretty basic concept of what energy is… • Chapter 6 is basically “Energy in Chemical Rxns.” • Thermochemistry… • Instinctively, we know chem. rxn’s can produce heat, after all, natural gas heats our homes. • CH4 + 2 O2→ CO2 + 2 H2O • Reaction produces heat…a form of energy • Take a step back and define energy…

  2. Flow chart of ideas

  3. Energy…defined • Best defined as “the capacity to do work” or capacity to move something. • Several types of energy…we’ve already covered… • Kinetic energy (energy of something in motion) • Potential energy (objects that possess the ability to do work—through position or chemical makeup • Also something called chemical energy…that’s more like potential energy… • Unit of energy is called the joule (J)…and a reason why scientists are often viewed as socially inept

  4. Characterizing potential v kinetic

  5. What’s wrong with this? • While on his honeymoon in Switzerland, James Joule measured the temperature of the water at the top of a waterfall and again at the bottom. What observation do you think he made? • Did you get past the first four words? • Observations? • Wife was ticked? • Didn’t understand the concept of honeymoon? • Sigh…perhaps best to move along…defining terms

  6. Thermochemistry: Basic Terms • Thermochemistry is the study of energy changes that occur during chemical reactions. • System: the part of the universe being studied. • Surroundings: the rest of the universe. • Open: energy and matter can be exchanged with the surroundings. • Closed: energy can be exchanged with the surroundings, matter cannot. • Isolated: neither energy nor matter can be exchanged with the surroundings.

  7. Heat (q) • Technically speaking, heat is not“energy.” • Heat is energy transfer between a system and its surroundings, caused by a temperature difference. More energetic molecules … … transfer energy to less energetic molecules. • Thermal equilibrium occurs when the system and surroundings reach the same temperature and heat transfer stops.

  8. Work (w) • Like heat, work is an energy transfer between a system and its surroundings. • Unlike heat, work is caused by a force moving through a distance (heat is caused by a temperature difference). • A negative quantity of work signifies that the system loses energy. • A positive quantity of work signifies that the system gains energy. • There is no such thing as “negative energy” nor “positive energy”; the sign of work (or heat) signifies the direction of energy flow.

  9. First Law of Thermodynamics • “Energy cannot be created or destroyed.” • Inference: the internal energy change of a system is simply the difference between its final and initial states: DU = Ufinal – Uinitial • Additional inference: if energy change occurs only as heat (q) and/or work (w), then: DU = q + w

  10. From last time… • Terms defined—work, heat etc • Heat—energy transferred between two objects • Work—a force acting over a distance • “Here’s Your Sign”—very important • Energy entering system—sign is positive • Energy leaving system—sign is negative • In a chemical reaction (define system) • Energy given to surroundings “exothermic” (-) • Energy absorbed by system “endothermic” (+)

  11. A bit more on that “exothermic” thing • Exothermic/endothermic rxns have specific meanings for us • Exothermic rxn means chemical energy becomes thermal energy. • If system is isolated (matter/energy contained) T . • If system is not isolated—thermal energy given to surroundings • Endothermic reactions are just the opposite • Isolated system implies thermal energy converted to chemical energy, and the temperature drops • Non-isolated system results in heat absorbed from the surroundings

  12. Properties of Enthalpy • Enthalpy is an extensive property. • It depends on how much of the substance is present. Two logs on a fire give off twice as much heat as does one log. • Enthalpy changeshave unique values. DH = qP • We’ll make use of these soon • Occasionally use “enthalpy diagrams” to represent changes

  13. Enthalpy Diagrams • Values of DH are measured experimentally. • Negative values indicate exothermic reactions. • Positive values indicate endothermic reactions. An increase in enthalpy during the reaction; DH is positive. A decrease in enthalpy during the reaction; DH is negative.

  14. Same magnitude; different signs. Reversing a Reaction • DH changes sign when a process is reversed. • Therefore, a cyclic process has the value DH = 0.

  15. Given the equation (a) H2(g) + I2(s)  2 HI(g) DH = +52.96 kJ calculate DH for the reaction (b) HI(g)  ½ H2(g) + ½ I2(s). Note here that the reaction on the bottom (b) is the reverse of (a) ALSO a fractional coefficient Extensive property (the amount matters)—reversed reaction (the sign changes) SPECIFIC to each reaction—we can incorporate this into our stoichiometric scheme.

  16. ΔH in Stoichiometric Calculations • For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant. • We can generate conversion factors involving DH. • For example, the reaction: H2(g) + Cl2(g)  2 HCl(g) DH = –184.6 kJ can be used to write: –184.6 kJ ———— 1 mol Cl2 –184.6 kJ ———— 2 mol HCl –184.6 kJ ———— 1 mol H2

  17. Adding to our molar conversion theme • Heat is a quantity of a reaction just like mass, molarity, ideal gas laws…etc. Heat MOLARITY Heat MOLARITY MOLE MOLE Gram Gram PV = nRT PV = nRT

  18. Calculating heat flow (enthalpy) What is the enthalpy change associated with the combustion of 5.67 mol CH4 in the following reaction? CH4(g) + 2 O2(g)  CO2 (g) + 2 H2O (g) DH = –890.3 kJ Simple conversion really 5.67 molmethane x -890.3 kJ/molmethane = -5048 kJ Or stated—5048 kJ of energy RELEASED to the surroundings How to determine?

  19. Calorimetry • We measure heat flow using calorimetry. • A calorimeter is a device used to make this measurement. • A “coffee cup” calorimeter may be used for measuring heat involving solutions. A “bomb” calorimeter is used to find heat of combustion; the “bomb” contains oxygen and a sample of the material to be burned.

  20. Calorimetry, Heat Capacity, Specific Heat • Heat evolved in a reaction is absorbed by the calorimeter and its contents. • In a calorimeter we measure the temperature changeof water or a solution to determine the heat absorbed or evolved by a reaction. • The heat capacity (C) of a system is the quantity of heat required to change the temperature of the system by 1 °C. C = q/DT (units are J/°C) • Molar heat capacity is the heat capacity of one mole of a substance. • The specific heat (s) is the heat capacity of one gram of a pure substance (or homogeneous mixture). s = C/m = q/(mDT) q = s m DT

  21. Table of useful specific heats • Note water is ridiculously large relative to most other materials • Metals really have very low SH’s, that’s why they heat up so fast.

  22. Heat Capacity: A Thought Experiment • Place an empty iron pot weighing 5 lb on the burner of a stove. • Place an iron pot weighing 1 lb and containing 4 lb water on a second identical burner (same total mass). • Turn on both burners. Wait five minutes. • Which pot handle can you grab with your bare hand? • Iron has a lower specific heat than does water. It takes less heat to “warm up” iron than it does water.

  23. Calculate the heat capacity of an aluminum block that must absorb 629 J of heat from its surroundings in order for its temperature to rise from 22 °C to 145 °C. This is a simple P&C example from previous eq. 2nd example—how much heat does it take to raise the temp of water (814 g) from 38 °C to 50 °C? Again, know eq., the Cp, the mass and the DT…done!

  24. Some real world examples • Everyone knows by now the combustion rxn of methane with oxygen. • Methane burns, the system is the molecular level of CH4 CO2 and H2O • System is open, which implies heat is given to surroundings (our homes, for example). qrxn • ENDOTHERMIC reactions—not as common • The “ice-pack” in sports…H2O and NH4NO3 mix (endothermically). Heat is taken from surroundings (the air and the skin of injured athlete)—temp drops

  25. Properties of Enthalpy • Enthalpy is an extensive property. • It depends on how much of the substance is present. Two logs on a fire give off twice as much heat as does one log. • Enthalpy changeshave unique values. DH = qP • We’ll make use of these soon • Occasionally use “enthalpy diagrams” to represent changes

  26. ΔH in Stoichiometric Calculations • For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant. • We can generate conversion factors involving DH. • For example, the reaction: H2(g) + Cl2(g)  2 HCl(g) DH = –184.6 kJ can be used to write: –184.6 kJ ———— 1 mol H2 –184.6 kJ ———— 1 mol Cl2 –184.6 kJ ———— 2 mol HCl

  27. Adding to our molar conversion theme • Heat is a quantity of a reaction just like mass, molarity, ideal gas laws…etc. Heat MOLARITY Heat MOLARITY MOLE MOLE Gram Gram PV = nRT PV = nRT

  28. Hess’s Law • So…where do these heat values come from anyway? • Some reactions cannot be carried out “as written.” • Consider the reaction: C(graphite) + ½ O2(g)  CO(g). • If we burned 1 mol C in ½ mol O2, both CO and CO2 would probably form. Some C might be left over. However …

  29. Hess’s Law Continued • … enthalpy change is a state function. • The enthalpy change of a reaction is the same whether the reaction is carried out in one step or through a number of steps. • Hess’s Law: If an equation can be expressed as the sum of two or more other equations, the enthalpy change for the desired equation is the sum of the enthalpy changes of the other equations.

  30. Hess’s Law: An Enthalpy Diagram We can find DH(a) by subtracting DH(b) from DH(c) So let’s have a look at how this works in an example…

  31. Calculate the enthalpy change for reaction (a) given the data in equations (b), (c), and (d). (a) 2 C(graphite) + 2 H2(g)  C2H4(g) DH = ? (b) C(graphite) + O2(g)  CO2(g) DH = –393.5 kJ (c) C2H4(g) + 3 O2  2 CO2(g) + 2 H2O(l) DH = –1410.9 kJ • H2(g) + ½ O2  H2O(l) DH = –285.8 kJ So this leads us to a more general way of stating and/or calculating enthalpy changes. Question: Where do these values of DH come from?

  32. Standard Enthalpies of Formation • It would be convenient to be able to use the simple relationship • ΔH = Hproducts – Hreactants • to determine enthalpy changes. • Although we don’t know absolute values of enthalpy, we don’t need them; we can use a relative scale. • We define the standard state of a substance as the state of the pure substance at 1 atm pressure and the temperature of interest (usually 25 °C). • The standard enthalpy change (ΔH°) for a reaction is the enthalpy change in which reactants and products are in their standard states. • The standard enthalpy of formation (ΔHf°) for a reaction is the enthalpy change that occurs when 1 mol of a substance is formed from its component elements in their standard states.

  33. Standard Enthalpy of Formation When we say “The standard enthalpy of formation of CH3OH(l) is –238.7 kJ”, we are saying that the reaction: C(graphite) + 2 H2(g) + ½ O2(g)  CH3OH(l) has a value of ΔH of –238.7 kJ. (ONE MOLE methanol) We can treat ΔHf° values as though they were absolute enthalpies, to determine enthalpy changes for reactions. Question: What is ΔHf°for an element in its standard state [such as O2(g)]? Hint: since the reactants are the same as the products …

  34. Calc’s Using Standard Enthalpies of Formation DH°rxn = SnpxDHf°(products) – SnrxDHf°(reactants) • The symbol S signifies the summation of several terms. • The symbol n signifies the stoichiometric coefficient used in front of a chemical symbol or formula. • In other words … PRODUCTS minus REACTANTS • Add all of the values for DHf° of the products. • Add all of the values for DHf° of the reactants. • Subtract #2 from #1 (This is usually much easier than using Hess’s Law!)

  35. Some helpful numbers Other numbers are in the back of the book—NOTE the box for H2O. The phase MATTERS!!

  36. Example 6.15 Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds. One reaction for producing synthesis gas is 3 CH4(g) + 2 H2O(l) + CO2(g)  4 CO(g) + 8 H2(g) ΔH° = ? Use standard enthalpies of formation from Table 6.2 to calculate the standard enthalpy change for this reaction. Example 6.16 The combustion of isopropyl alcohol, common rubbing alcohol, is represented by the equation 2 (CH3)2CHOH(l) + 9 O2(g)  6 CO2(g) + 8 H2O(l) ΔH° = –4011 kJ Use this equation and data from Table 6.2 to establish the standard enthalpy of formation for isopropyl alcohol.

  37. Ionic Reactions in Solution • We can apply thermochemical concepts to reactions in ionic solution by arbitrarily assigning an enthalpy of formation of zero to H+(aq).

  38. Example 6.18 H+(aq) + OH–(aq)  H2O(l) ΔH° = –55.8 kJ Use the net ionic equation just given, together with ΔHf° = 0 for H+(aq), to obtain ΔHf° for OH–(aq). Note that we’ve already sort of calculated this from coffee cup calorimetry!!

  39. Looking Ahead • A reaction that occurs (by itself) when the reactants are brought together under the appropriate conditions is said to be spontaneous. • A discussion of entropy is needed to fully understand the concept of spontaneity, and will be discussed in Chapter 17. • A spontaneous reaction isn’t necessarily fast (rusting; diamond  graphite; etc. are slow). • The difference between the tendency of a reaction to occur and the rate at which a reaction occurs will be discussed in Chapter 13.

  40. Combustion and Respiration:Fuels and Foods • Fossil Fuels: Coal, Natural Gas, and Petroleum • A fuel is a substance that burns with the release of heat. • These fossil fuels were formed over a period of millions of years from organic matter that became buried and compressed under mud and water. • Foods: Fuels for the Body • The three principal classes of foods are carbohydrates, fats, and proteins. • 1 Food Calorie (Cal) is equal to 1000 cal (or 1 kcal).

  41. CUMULATIVE EXAMPLE The human body is about 67% water by mass. What mass of sucrose, C12H22O11(s), for which ΔHf° = –2225 kJ/mol, must be metabolized by a 55-kg person with hypothermia (low body temperature) to raise the temperature of the body water from 33.5 °C to the normal body temperature of 37.0 °C? Assume the products of metabolism are in their most stable states at 25 °C. What volume of air at 37 °C having a partial pressure of O2 of 151 Torr is required for the metabolism?

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