Chapter 6. Heat Flow
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1 Chapter 6 Thermochemistry Heat Flow Heat (q): energy transferred from body at high T to body at low T Two definitions: System: part of universe we are interested in Surrounding: the rest of the universe Direction of heat flow defined in terms of system 1
2 The System For chemists, the system usually is the set of chemicals involved in a reaction Everything else (solvent, test tube, ourselves, etc.) is part of the surroundings Direction of heat flow defined in terms of system System loses energy to surroundings: negative heat System gains energy from surroundings: positive heat Exothermic & Endothermic q <0: System loses energy exothermic q > 0: System gains energy endothermic Remember: We are part of the surroundings! What does an exothermic process feel like to us? What about endothermic? 2
3 Heats of Reaction 3 C 11 H 22 O 11 (s) + 22 KClO 3 (l) ----> 33 CO 2 (g) + 33 H 2 O (l) + 22 KCl (s) System = gummy bear and potassium chlorate What do you observe? Exothermic or endothermic reaction? Heats of Reaction Ba(OH) 2 10H 2 O (s) + 2 NH 4 SCN (s) > Ba(SCN) 2 (aq) + 2 NH 3 (aq) + 12 H 2 O (l) System = The two chemicals Exothermic or Endothermic? 3
4 Heat Capacity Adding heat to a substance raises its temperature High heat capacity A lot of heat only raises temperature a little High ability to absorb energy without making molecules move Low heat capacity Doesn t take much heat to raise temperature Small amount of absorbed energy really gets molecules moving Heat Capacity Mathematical expression for heat capacity In terms of Calculus C = dq/dt 4
5 Heat Capacity and Specific Heat Specific heat (s) = heat capacity of one gram C = (mass)(specific heat) Question: For water, s = J/(g o C). How much heat required to raise temp. of 50.0 g water by 10 o C? Thermodynamics Scientific study of interconversion of heat and other types of energy Deep and subtle topic I am handling topic differently from text Based on three laws 5
6 First Law of Thermodynamics Conservation of Energy Energy cannot be created or destroyed It can be converted from one form to another Potential energy (ball at top of hill) -> kinetic energy (moving ball) Chemical energy (gasoline) -> thermal energy (moving molecules) Radiant energy (sunlight) -> chemical energy (carbohydrates) First Law of Thermodynamics ΔE = q + w Internal energy change = heat + work Chemists usually interested in the heat, q I am not covering calculations of work 6
7 Enthalpy Usual lab conditions are constant pressure Heat is same as another thermodynamic function, the enthalpy, at constant pressure ΔH = q at constant pressure Enthalpy (H) 2 CH 3 OH (l) + 3 O 2 (g) ---> 4 H 2 O (l) + 2 CO 2 (g) ΔH = kj/mole ΔH units are per mole of reaction Conversion factors: kj/2 moles CH 3 OH kj/3 moles O kj/4 moles H 2 O kj/2 moles CO 2 7
8 Enthalpy 2 CH 3 OH (l) + 3 O 2 (g) ---> 4 H 2 O (l) + 2 CO 2 (g) ΔH = kj/mole Exo- or endothermic? Feel hot or cold to us? What is ΔH in kj if 5.00 g of methanol are burned? Constant Pressure Calorimetry Measure heat of reaction when P = constant When a reaction occurs inside this calorimeter heat lost or gained by rxn = heat gained or lost by cup and water q rxn = - (C cup + C water ) T Note negative sign C water = (mass water)(4.18 J/(g o C) For an insulated calorimeter C cup 0 Won t do constant V calorimetry 8
9 Constant Pressure Calorimetry Example: ml of 1.00 M HCl are mixed with ml of 1.00M NaOH. The temperature rises by 6.72 o C. Assume C cup = 0 q rxn = - C water T = -m water s water T What is the net ionic reaction? What is the heat of the reaction in kj? Standard Enthalpies of Formation Want to tabulate enthalpies Problems Only enthalpy changes can be measured (ΔH = q p ) Need to specify pressure Need to specify concentrations of solutions Solutions Set some reference enthalpies = 0 Define other enthalpies relative to reference enthalpies Choose 1 atm as standard pressure Choose 1 M as standard concentration 9
10 Standard Enthalpy of Formation Tabulated values (Appendix 3 in text) A few on page 253 Standard Conditions H f o = 0 for elements in standard state (most stable form) P = 1 atm Concentrations = 1 M H o (reaction) = H f o (products) - H f o (reactants) Standard Enthalpy of Formation C (s, graphite) + O 2 (g) -----> CO 2 (g) Produce 500 ml CO 2 at 1.00 atm and 25 o C. Heat of reaction is J What is ΔH f o of CO 2 in kj/mole? 10
11 Standard Enthalpy of Formation Question: What is ΔH o for combustion of g of methane? CH 4 (g) + 2 O 2 (g) -----> CO 2 (g) + 2 H 2 O (l) Exo- or endothermic? Tables on last slide Combining Reactions 1. Magnitude of H is directly proportional to the amount of reactant or product. 2. Reversing reaction changes sign of H. 3. Hess s Law: If two or more reactions can be added to give a total reaction, then H for the total reaction is the sum of H s for individual reactions. 11
12 Combining Reactions Given these reactions: C 2 H 5 OH (l) + 3 O 2 (g) -----> 2 CO 2 (g) + 3 H 2 O (l) H 1 = kj CO (g) + ½ O 2 (g) -----> CO 2 (g) H 2 = -566 kj What is H for reaction below? C 2 H 5 OH (l) + 2 O 2 (g) -----> 2 CO (g) + 3 H 2 O (l) H 3 =? We know that Combining Reactions 3 Fe (s) + 2 O 2 (g) -----> Fe 3 O 4 (s) H 1 = kj 4 Fe (s) + 3 O 2 (g) -----> 2 Fe 2 O 3 (s) H 2 = kj What is H for the reaction below 4 Fe 3 O 4 (s) + O 2 (g) -----> 6 Fe 2 O 3 (s) H =? Exothermic or endothermic? What is H per mole of Fe 2 O 3 formed? 12
13 Thermodynamic Tables Pressure = 1 atm Concentrations = 1 M 25 o C Substance H o f (kj/mole) G o f (kj/mole) S o (J/mole K) CH 4 (g) CO (g) CO 2 (g) Cl 2 (g) Fe (s) Fe 2 O 3 (s) H 2 O(l) H 2 O(g) I 2 (g) ICl(g) O 2 (g)
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