Thermodynamics. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

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1 Thermodynamics Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2 Thermodynamics is the scientific study of the interconversion of heat and other kinds of energy. The system is the specific part of the universe that is of interest in the study. open closed isolated Exchange: mass & energy energy nothing 2

3 Energy is the capacity to do work. Radiant energy comes from the sun and is earth s primary energy source Thermal energy is the energy associated with the random motion of atoms and molecules Chemical energy is the energy stored within the bonds of chemical substances Nuclear energy is the energy stored within the collection of neutrons and protons in the atom Potential energy is the energy available by virtue of an object s position 3

4 Energy Changes in Chemical Reactions Heat is the transfer of thermal energy between two bodies that are at different temperatures. Temperature is a measure of the thermal energy. Temperature = Thermal Energy 4

5 Exothermic process is any process that gives off heat transfers thermal energy from the system to the surroundings. 2H 2 (g) + O 2 (g) 2H 2 O(l) + energy H 2 O(g) H 2 O(l) + energy Endothermic process is any process in which heat has to be supplied to the system from the surroundings. energy + 2HgO(s) energy + H 2 O(s) 2Hg(l) + O 2 (g) H 2 O(l) 5

6 Thermochemistry is the study of heat change in chemical reactions. State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. energy, pressure, volume, temperature ΔU = U final - U initial ΔP = P final - P initial ΔV = V final - V initial ΔT = T final - T initial Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. 6

7 First law of thermodynamics energy can be converted from one form to another, but cannot be created or destroyed. ΔU system + ΔU surroundings = 0 or ΔU system = -ΔU surroundings C 3 H 8 + 5O 2 3CO 2 + 4H 2 O Exothermic chemical reaction! Chemical energy lost by combustion = Energy gained by the surroundings 7 system surroundings

8 Another form of the first law for ΔU system ΔU = q + w ΔU is the change in internal energy of a system q is the heat exchange between the system and the surroundings w is the work done on (or by) the system w = -PΔV when a gas expands against a constant external pressure 8

9 w = F x d Work Done On the System w = -P ΔV P x V = F d 2 x d 3 = F x d = w ΔV > 0 -PΔV < 0 w sys < 0 Work is not a state function. Δw = w final - w initial initial final 9

10 A sample of nitrogen gas expands in volume from 1.6 L to 5.4 L at constant temperature. What is the work done in joules if the gas expands (a) against a vacuum and (b) against a constant pressure of 3.7 atm? w = -PΔV (a) ΔV = 5.4 L 1.6 L = 3.8 L P = 0 atm W = -0 atm x 3.8 L = 0 L atm = 0 joules (b) ΔV = 5.4 L 1.6 L = 3.8 L P = 3.7 atm w = -3.7 atm x 3.8 L = L atm w = L atm x J 1L atm = J 10

11 Enthalpy and the First Law of Thermodynamics ΔU = q + w At constant pressure: q = ΔH and w = -PΔV ΔE = ΔH - PΔV ΔH = ΔU + PΔV 11

12 Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. ΔH = H (products) H (reactants) ΔH = heat given off or absorbed during a reaction at constant pressure H products > H reactants ΔH > 0 H products < H reactants ΔH < 0 12

13 Thermochemical Equations Is ΔH negative or positive? System absorbs heat Endothermic ΔH > kj are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm. H 2 O(s) H 2 O(l) ΔH = 6.01 kj/mol 13

14 Thermochemical Equations Is ΔH negative or positive? System gives off heat Exothermic ΔH < kj are released for every 1 mole of methane that is combusted at 25 0 C and 1 atm. CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) ΔH = kj/mol 14

15 Thermochemical Equations The stoichiometric coefficients always refer to the number of moles of a substance H 2 O(s) H 2 O(l) ΔH = 6.01 kj/mol If you reverse a reaction, the sign of ΔH changes H 2 O(l) H 2 O(s) ΔH = kj/mol If you multiply both sides of the equation by a factor n, then ΔH must change by the same factor n. 2H 2 O(s) 2H 2 O(l) ΔH = 2 x 6.01 = 12.0 kj 15

16 Thermochemical Equations The physical states of all reactants and products must be specified in thermochemical equations. H 2 O(s) H 2 O(l) ΔH = 6.01 kj/mol H 2 O(l) H 2 O(g) ΔH = 44.0 kj/mol How much heat is evolved when 266 g of white phosphorus (P 4 ) burn in air? P 4 (s) + 5O 2 (g) P 4 O 10 (s) ΔH = kj/mol 266 g P 4 x 1 mol P g P 4 x 3013 kj 1 mol P 4 = 6470 kj 16

17 A Comparison of ΔH and ΔU 2Na(s) + 2H 2 O(l) ΔU = ΔH - PΔV 2NaOH(aq) + H 2 (g) ΔH = kj/mol At 25 C, 1 mole H 2 = 24.5 L at 1 atm PΔV = 1 atm x 24.5 L = 2.5 kj ΔU = kj/mol 2.5 kj/mol = kj/mol 17

18 The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius. The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius. C = m x s Heat (q) absorbed or released: q = m x s x Δt q = C x Δt Δt = t final - t initial 18

19 How much heat is given off when an 869 g iron bar cools from 94 o C to 5 o C? s of Fe = J/g C Δt = t final t initial = 5 C 94 C = -89 C q = msδt = 869 g x J/g C x 89 C = -34,000 J 19

20 Constant-Volume Calorimetry q sys = q water + q bomb + q rxn q sys = 0 q rxn = - (q water + q bomb ) q water = m x s x Δt q bomb = C bomb x Δt Reaction at Constant V ΔH = q rxn No heat enters or leaves! ΔH ~ q rxn 20

21 Constant-Pressure Calorimetry q sys = q water + q cal + q rxn q sys = 0 q rxn = - (q water + q cal ) q water = m x s x Δt q cal = C cal x Δt Reaction at Constant P ΔH = q rxn No heat enters or leaves! 21

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23 Because there is no way to measure the absolute value of the enthalpy of a substance, must I measure the enthalpy change for every reaction of interest? Establish an arbitrary scale with the standard enthalpy of formation (ΔH f ) as a reference point for all enthalpy expressions. Standard enthalpy of formation (ΔH f ) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. The standard enthalpy of formation of any element in its most stable form is zero. ΔH ΔH f (C, graphite) = 0 f (O 2 ) = 0 ΔH (O 3 ) = 142 kj/mol f ΔH (C, diamond) = 1.90 kj/mol f 23

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25 The standard enthalpy of reaction (ΔH a reaction carried out at 1 atm. rxn ) is the enthalpy of aa + bb cc + dd ΔH rxn = [ cδh f (C) + dδh (D)] - [ aδh f (A) + bδh f (B) ] ΔH rxn f = Σ nδh f (products) - Σ mδh f (reactants) Hess s Law: When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. (Enthalpy is a state function. It doesn t matter how you get there, only where you start and end.) 25

26 C (graphite) + 1/2O 2 (g) CO(g) + 1/2O 2 (g) C (graphite) + O 2 (g) CO 2 (g) CO(g) CO 2 (g) 26

27 Calculate the standard enthalpy of formation of CS 2 (l) given that: C(graphite) + O 2 (g) CO 2 (g) ΔH rxn = kj/mol S(rhombic) + O 2 (g) SO 2 (g) ΔH rxn = kj/mol CS 2 (l) + 3O 2 (g) CO 2 (g) + 2SO 2 (g) ΔH = kj/mol rxn 1. Write the enthalpy of formation reaction for CS 2 C(graphite) + 2S(rhombic) CS 2 (l) 2. Add the given rxns so that the result is the desired rxn. + C(graphite) + O 2 (g) CO 2 (g) ΔH rxn = kj/mol 2S(rhombic) + 2O 2 (g) 2SO 2 (g) ΔH rxn = kj/mol x 2 CO 2 (g) + 2SO 2 (g) CS 2 (l) + 3O 2 (g) ΔH = kj/mol rxn CS 2 (l) ΔH = (2x-296.1) = 86.3 kj/mol rxn C(graphite) + 2S(rhombic) 27

28 Benzene (C 6 H 6 ) burns in air to produce carbon dioxide and liquid water. How much heat is released per mole of benzene combusted? The standard enthalpy of formation of benzene is kj/mol. 2C 6 H 6 (l) + 15O 2 (g) 12CO 2 (g) + 6H 2 O (l) ΔH rxn = Σ nδh f (products) - Σ mδh (reactants) ΔH rxn = [ 12ΔH (CO 6ΔH 2 ) + (H 2 O) ] - [ 2ΔH (C 6 H 6 )] f ΔH rxn = [ 12x x ] [ 2x49.04 ] = kj f f f kj 2 mol = kj/mol C 6 H 6 28