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Chemistry

Chemistry. Thermodynamics. I. Enthalpy: Endothermic & Exothermic Reactions. A. NRG can be thought of as a reactant or product: NRG + Reactant → Product (Endothermic) Reactant → Product + NRG (Exothermic). B. Terminology

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Chemistry

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  1. Chemistry Thermodynamics

  2. I. Enthalpy: Endothermic & Exothermic Reactions A. NRG can be thought of as a reactant or product: NRG + Reactant → Product (Endothermic) Reactant → Product + NRG (Exothermic)

  3. B. Terminology 1. Enthalpy (H): Potential Energy associated with a substance (Wile) - Kind of right – we’ll go with it because it’s simple 2. Enthalpy of reaction (∆H): The difference between the enthalpies of the products minus the enthalpies of the reactants ∆H = H (products) – H (reactants)

  4. 3. Units of ∆H: NRG units (calories, Joules) 1.000 cal = 4.184 Joule 4. Signs: ∆H is + for endothermic ∆H is - for exothermic

  5. 5. Example: When CH4 is burned, the enthalpy of the reaction is -802.2 kJ/mol. Write a balanced chemical equation representing this process. Include NRG in your equation.

  6. OYO 13.1 Write balanced chemical equations for each of the following processes. Include NRG as either a reactant or product. A) Formation of H2SO4; ∆H = -814 kJ B) Decomposition of K3PO4; ∆H = 2,030 kJ C) Reaction of H2CO3 & KOH; releases 21 kcals

  7. II. Determining ∆H for a Chemical Reaction by Calorimetry Remember q = mc∆T ? q is negative when NRG is released q is positive when NRG is absorbed ∆H is equal (or very close) to qrxn

  8. III. Determining ∆H using Standard Enthalpy of Formations A. The definition Standard Enthalpy of Formation (∆Hfo): the ∆H of a formation reaction at standard conditions “o” means Standard Conditions (25oC, 1.00 atm) -- different from STP (0.0 oC, 1.00 atm) “f” means formation

  9. We can use ∆Hfo for reactants and products instead of H (which we can’t measure) So, we said ∆H= H (products) – H (reactants) can’t measure can’t measure We’ll use: B. The equation ∆Hrxno = ∑ ∆Hfo (prod) - ∑ ∆Hfo (reac.) (∑ means “sum of”)

  10. l C. Some Standard Enthalpies of Formation -By convention, the standard enthalpy of formation for any element = 0 - Note that the phase is important! - Note that units are kJ/mole

  11. D. Exothermic or Endothermic?Using ∆Hrxno = ∑ ∆Hfo (prod) - ∑ ∆Hfo (reac.) 1. What is∆Hrxnoof the following reaction? 2NO(g) + O2(g) → 2NO2(g) 2. Some gasoline companies put a small amount of liquid ethyl alcohol (C2H6O) into their gasoline. What is the ∆Hrxno?

  12. OYO’s 13.2 When you use a pocket cigarette lighter, you are burning gaseous butane (C4H10). Assuming it is complete combustion, what is ∆Ho for this reaction? 13.3 In order to make the shells that house and protect them, many shellfish take solid lime (CaO) from the ocean floor and react it with gaseous carbon dioxide that has been exhaled by marine organisms. The resulting calcium carbonate is the major component of their shell. What is the ∆Ho for this reaction?

  13. IV. Applying Enthalpy to Stoichiometry Example 1: The ∆Horxn for the combustion of C2H6O is -1236 kJ. How much energy can be produced by burning 150.0 g of ethyl alcohol in excess oxygen?

  14. Example 2: One way of producing butane (C4H10) for lighters is to force ethane (C2H6) to react with itself: 2C2H6(l) → C4H10(g) + H2(g) If you have 1.23 kg of ethane and would like to turn it all into butane, how many kJ of energy would be necessary?

  15. OYO’s 13.4 Ethane gas (C2H6) can be used for auto fuel because it releases a lot of NRG when burned. How many kJ of NRG are released when 250.0 g of ethane are burned? 13.5 When mined, coal has a small amount of sulfur in it. When the coal is burned in a power plant, SO3(g) is produced. When the SO3(g) reacts with water in the air, the following reaction occurs: H2O(g) + SO3(g) → H2SO4(l) The sulfuric acid falls to earth as acid rain. How much NRG is produced when 50.0 g of H2SO4 are produced in this way?

  16. V. NRG Diagrams A. Exothermic Reaction C2H6O(l) + H2(g) → C2H6(g) + H2O(g)

  17. Potential NRG = ∆H Reaction Coordinate = Rxn progress Intermediate Stage = “Activated Complex” B. Activation NRG: The NRG necessary to start a chemical reaction All reactions require activation NRG

  18. C. NRG Diagram for an Endothermic Rxn

  19. C. Using NRG Diagrams Which is most exothermic? ∆H? Which is easiest to get started? Which is endothermic? ∆H?

  20. OYO 13.6 You perform 3 experiments: Experiment A: hard to start, gets very hot Experiment B: easy to start, gets only slightly warm Draw NRG diagrams for both. (Label the y and x axis – but you will have no numbers)

  21. VI. Entropy A. The definition: Entropy (S): a measure of how spread out the NRG is in any system 1. Units: J/mol K 2. Can measure absolute entropy (3rd Law of Thermodynamics: a perfect crystal has 0 S at 0.00 K)

  22. 3. Factors that affect entropy a. Phases of matter solid – least liquid – next gas – most b. Entropy increases with temperature c. Entropy increases with increasing matter

  23. B. Determining whether entropy increases or decreases Is ∆S positive or negative? Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g) N2(g) + 3H2(g) → 2NH3(g) 2NH4NO3(s) + Ba(OH)2*8H2O(s) → Ba(NO3)2(aq) + 2NH4OH(aq) + 8H2O(l)

  24. OYO’s 13.7 How does the entropy of a system change for each of the following processes? a. A solid melts b. A liquid freezes c. A liquid boils d. A vapor is converted to a solid e. Urea dissolves in water

  25. 13.8 Determine the sign of ∆S for the following: a. C2H4O2(aq) + NaHCO3(s) → NaC2H3O2(aq) + H2O(l) + CO2(g) b. 2H2O2(g) → 2H2O(g) + O2(g) c. 2AgNO3(aq) + Mg(OH)2(aq) → 2AgOH(s) + Mg(NO3)2(aq)

  26. C. Determining ∆So for a Chemical Rxn All substances have a positive entropy (except for a perfect crystalline solid at 0 K) S has the units J/mol K (notice the units are J instead of kJ – entropy NRG is very small comparatively) ∆Srxno = ∑ ∆Sfo (prod) - ∑ ∆Sfo (reac.)

  27. OYO 13.9 Calculate ∆Sofor the following: CaO(s) + CO2(g) → CaCO3(s)

  28. VII. 2nd Law of Thermodynamics A. Definition: 1. The entropy of the universe increases or remains the same for spontaneous processes 2. Energy of all types spontaneously flows from being more localized or concentrated to becoming more dispersed or spread out, if not hindered. ∆Suniverse≥ 0 B. Spontaneous: a reaction that occurs naturally under certain circumstances One of the main reasons chemists study thermodynamics is so they are able to predict whether a reaction will be spontaneous!

  29. Spontaneous Processes - Waterfall runs downhill, not up - Sugar lump dissolves in water, sugar lump doesn’t reappear - Water freezes below 0oC; ice melts above 0oC - Heat flows from hot object to a colder one; the reverse never happens spontaneously - Gas expands into an evacuated bulb spontaneously; the reverse process (gas molecules gathering into one bulb) is not spontaneous

  30. OYO’s 13.10 Which of the following processes are spontaneous and which are nonspontaneous? a. Dissolving NaCl in soup b. Climbing Mt. Everest c. Spreading fragrance in a room by removing the bottle cap d. Separating He and Ne from a mixture of gases

  31. VIII. Gibbs Free NRG (G – energy available to do work): How to tell if a reaction is spontaneous or not A. Two driving forces determine whether a reaction is spontaneous or not 1. Whether the product(s) have more entropy than the reactant(s) 2. Whether the reaction is endothermic or exothermic a. Exothermic increases the entropy of the surroundings b. Endothermic decreases the entropy of the surroundings

  32. B. The Gibbs Free Energy equation puts them together 1. The equation: ∆G = ∆H - T∆S 2. When ∆G< 0, reaction is spontaneous 3. When ∆G > 0, reaction is not spontaneous

  33. 4. Can also use: ∆Grxno = ∑ ∆Gfo (prod.) - ∑ ∆Gfo (reac.)

  34. C. Factors Affecting the sign in ∆G

  35. D. Using ∆G = ∆H - T∆S 1. 2NO(g) + O2(g) → 2NO2(g) ∆H = -114.2 kJ ∆S = -147 J/mol K *What is ∆G at 532oC (805 K)? *Is it spontaneous at that temperature? *Is there a temperature at which this rxn is spontaneous?

  36. 2. What is the ∆G for the following rxn at 298 K? Is it spontaneous? Ca(s) + 2HCl(aq) → CaCl2(s) + H2(g)

  37. OYO’s 13.11 The ∆Ho of a certain chemical reaction is 10.0 kJ whereas ∆So is 123 J/K. At what temperatures is this reaction spontaneous? 13.12 Is the following synthesis of ethyl alcohol a spontaneous reaction at 298 K? 2CH4(g) + O2(g) → C2H6O(l) + H2O(l)

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