Enthalpy (H) • The heat transferred sys ↔ surr during a chemical rxn @ constant P • Can’t measure H, only ΔH • At constant P, ΔH = q = mCΔT, etc. • Literally, ΔH = Hproducts - Hreactants • ΔH = + (endothermic) • Heat goes from surr into sys • ΔH = - (exothermic) • Heat leaves sys and goes into surr In this example, the energy of the system (reactants and products) ↑, while the energy of the surroundings ↓ Notice that the total energy does not change Reactant + Energy Product Endothermic Reaction Surroundings Energy Surroundings System System Myers, Oldham, Tocci, Chemistry, 2004, page 41 Before reaction After reaction In this example, the energy of the system (reactants and products) ↓, while the energy of the surroundings ↑ Notice again that the total energy does not change Reactant Product + Energy Energy Surroundings Myers, Oldham, Tocci, Chemistry, 2004, page 41 System Before reaction Exothermic Reaction Surroundings System After reaction Reaction Coordinate Diagrams: Endothermic Reaction Activation Energy Ea D(PE) PE Products Reactants Progress of the Reaction ΔHrxn= + Reaction Coordinate Diagrams: Exothermic Reaction Activation Energy Ea ΔHrxn= - PE Reactants D(PE) Products Progress of the Reaction Reaction Coordinate Diagrams Draw the reaction coordinate diagram for the following rxn: C(s) + O2(g) CO2 + 458.1kJ Activation Energy EXOTHERMIC Ea ΔHrxn= -458.1 kJ PE C + O2 D(PE) CO2 Progress of the Reaction Enthalpies of Reaction • All reactions have some ΔH associated with it H2(g) + ½ O2(g) → H2O(l) ΔH = - 483.6 kJ • How can we interpret this ΔH? • Amount of energy released or absorbed per specific reaction species • Use balanced equation to find several definitions - 483.6 kJ 1 mol H2 or - 483.6 kJ ½ mol O2 or - 483.6 kJ 1 mol H2O Able to use like conversion factors in stoichiometry Enthalpies of Reaction • Formation of water H2(g) + ½ O2(g) → H2O(l) ΔH = - 483.6 kJ • ΔH is proportional to amount used and will change as amount changes 2H2(g) + O2(g) → 2H2O(l) ΔH = - 967.2 kJ • For reverse reactions, sign of ΔH changes 2H2O(l) → 2H2(g) + O2(g) ΔH = + 967.2 kJ • Treat ΔH like reactant or product H2(g) + ½ O2(g) → H2O(l) ΔH = - 483.6 kJ H2(g) + ½ O2(g) → H2O(l) + 483.6 kJ (exo) Enthalpies of Reaction Practice Consider the following rxn: C(s) + 1/2O2(g) CO + 458.1kJ Is the ΔH for this reaction positive or negative? NEGATIVE (E released as a product) What is the ΔH for 2.00 moles of carbon, if all the carbon is used? 2.00 mol C - 458.1 kJ 1 mol C = - 916 kJ What is the ΔH if 50.0g of oxygen is used? 50.0 g O2 1 mol O2 32.0 g O2 - 458.1 kJ = -1430 kJ 0.5 mol O2 1 mol CO 28.0 g CO 458.1 kJ = 818 kJ 1 mol CO What is the ΔH if 50.0 g of carbon monoxide decompose, in the reverse reaction? 50.0 g CO Hess’s Law Reactants Products The change in enthalpy is the same whether the reaction takes place in one step or a series of steps Victor Hess Why? Because enthalpy is a state function To review: 1. If a reaction is reversed, ΔH is also reversed 2 CH4 + O2 2 CH3OH ΔHrxn = -328 kJ 2 CH3OH 2 CH4 + O2 ΔHrxn = +328 kJ 2. If the coefficients of a reaction are multiplied by an integer, ΔH is multiplied by that same integer CH4 + 2 O2 CO2 + 2 H2O 2(CH4 + 2 O2 CO2 + 2 H2O) ΔHrxn = -802.5 kJ ΔHrxn = -1605 kJ Example: Methanol-Powered Cars 2 CH3OH(l) + 3 O2(g) 2 CO2(g) + 4 H2O(g) ΔHrxn = ? - ( ΔHrxn = -328 kJ) - (2 CH4(g) + O2(g) 2 CH3OH(l) ) 2 (CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g)) 2 (ΔHrxn = -802.5 kJ ) 2 CH3OH(l) 2 CH4(g) + O2(g) ΔHrxn = +328 kJ 3 2 CH4(g) + 4 O2(g) 2 CO2(g) + 4 H2O(g) ΔHrxn = -1605 kJ 2 CH3OH + 3 O2 2 CO2 + 4 H2O ΔHrxn = -1277 kJ Tips for applying Hess’s Law… Look at the final equation that you are trying to create first… • Find a molecule from that eq. that is only in one of the given equations (i.e. CH3OH, CO2) • Make whatever alterations are necessary to those • Once you alter a given equation, you will not alter it again • Continue to do this until there are no other options • Next, alter remaining equations to get things to cancel that do not appear in the final equation 1. Given the following data: S(s) + 3/2O2(g) → SO3(g) ΔH = -395.2 kJ - ½ ( 2SO2(g) + O2(g) → 2SO3(g) ) - ½ (ΔH = -198.2 kJ) . Calculate ΔH for the following reaction: S(s) + O2(g) → SO2(g) * * S(s) + 3/2O2(g) → SO3(g) 2SO3(g) → O2(g) + 2SO2(g) SO3(g) → ½ O2(g) + SO2(g) ΔH = -395.2 kJ ΔH = +198.2 kJ ΔH = +99.1 kJ S(s) + O2(g) → SO2(g) ΔH = -296.1 kJ 2. Given the following data: -(C2H2(g) + 5/2O2(g) → 2CO2(g) + H2O(l) ΔH = -1300. kJ) 2( C(s) + O2(g) → CO2(g) ) 2(ΔH = -394 kJ) H2(g) + 1/2O2(g) → H2O(l) ΔH = -286 kJ Calculate ΔH for the following reaction: 2C(s) + H2(g) → C2H2(g) 2C(s) + 2O2(g) → 2CO2(g) ΔH = -788 kJ 2CO2(g) + H2O(l) → C2H2(g) + 5/2O2(g) ΔH = +1300 kJ H2(g) + 1/2O2(g) → H2O(l) ΔH = -286 kJ * * * 2C(s) + H2(g) → C2H2(g) ΔH = +226 kJ 3. Given the following data: -½( 2O3(g) → 3O2(g)) -½(O2(g) → 2O(g) ) NO(g) + O3(g) → NO2(g) + O2(g) -½( ΔH = - 427 kJ ) -½(ΔH = + 495 kJ) ΔH = - 199 kJ Calculate ΔH for the following reaction: NO(g) + O(g) → NO2(g) * * * NO(g) + O3(g) → NO2(g) + O2(g) O(g) → ½ O2(g) 3/2 O2(g) → O3(g) NO(g) + O(g) → NO2(g) ΔH = - 199 kJ ΔH = - 247.5 kJ ΔH = + 213.5 kJ ΔH = - 233 kJ 4. Given the following data: Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g) ΔH = -23 kJ 3Fe2O3(s) + CO(g) → 2Fe3O4(s) + CO2(g) ΔH = -39 kJ Fe3O4(s) + CO(g) → 3FeO(s) + CO2(g) ΔH = 18 kJ Calculate ΔH for the following reaction: FeO(s) + CO(g) → Fe(s) + CO2(g) Hess’s Law HW Questions 1. A B C B ∆H = + 30 kJ ∆H = - 60 kJ Calculate ΔH for the following reaction: A C 2. Suppose you are given the following reactions: 4X 2Y DH = - 40 kJ X ½Z DH = - 95 kJ Calculate ΔH for the following reaction: Y Z 3. From the following heats of reaction: 2 H2 (g) + O2 (g) 2 H2O (g) 3 O2 (g) 2 O3 (g) DH = -483.6 kJ DH = +284.6 kJ Calculate the heat of the reaction (∆H): 3 H2 (g) + O3 (g) 3 H2O (g) 4. From the following enthalpies of reaction: H2 (g) + F2 (g) 2 HF (g) DH = - 537kJ C (s) + 2 F2 (g) CF4 (g) DH = - 680 kJ 2 C (s) + 2 H2 (g) C2H4 (g) DH = + 52.3 kJ Calculate the DH for the reaction of ethylene with F2. C2H4 (g) + 6F2 (g) 2 CF4 (g) + 4 HF(g) 5. Given the following data: N2 (g) + O2 (g) 2 NO (g) 2 NO (g) + O2 (g) 2 NO2 (g) 2 N2O (g) 2 N2 (g) + O2 (g) Calculate DH for the reaction below: N2O (g) + NO2 (g) 3 NO (g) DH = + 180.7 kJ DH = - 113.1 kJ DH = - 162.3 kJ DH = ? Heats of Formation, ΔH°f The enthalpy change when one mole of a compound is formed from the elements in their standard states ° = standard conditions • Gases at 1 atm pressure • All solutes at 1 M concentration (remember M = mol/L) • Pure solids and pure liquids f = a formation reaction • 1 mole of product formed • From the elements in their standard states (1 atm, 25°C) For all elements in their standard states, ΔH°f = 0 What’s the formation reaction for adrenaline, C9H12NO3(s)? 9 Cgr + 6 H2(g) + 1/2 N2(g) + 3/2 O2(g) C9H12NO3(s) Thermite Reaction Fe2O3(s) + 2 Al(s) Al2O3(s) + 2 Fe(l) ΔHrxn = ? Welding railroad tracks Thermite Reaction Fe2O3(s) + 2 Al(s) Al2O3(s) + 2 Fe(l) Reactants Elements Products (standard states) Fe2O3(s) 2 Al(s) 2 Fe(s) 3/ 2 O2(g) 2 Fe(l) Al2O3(s) 2 Al(s) ΔHrxn = 2ΔH°f(Fe(l)) + ΔH°f(Al2O3(s)) - ΔH°f(Fe2O3(s)) - 2ΔH°f(Al(s)) ΔHrxn = 2(15 kJ) + (-1676 kJ) - (-822 kJ) – 2(0) ΔHrxn = - 824 kJ ΔHrxn = nΔH°f(products) - nΔH°f(reactants) ΔH°f Example Problems ∆Hrxn = Σ n∆Hof Products - Σn∆Hof Reactants 1. CH4(g) + 2 Cl2(g) CCl4(g) + 2 H2(g) ΔHrxn = ? (- 106.7) 2 (0) (- 74.8) 2 (0) ∆H = [(-106.7) + 0] – [(-74.8)+0] = -106.7 + 74.8 = - 31.9 kJ 2. 2 KCl(s) + 3 O2(g) 2KClO3(s) 2 (- 435.9) 3 (0) 2 (- 391.2) ∆H = [(2)(- 391.2)] – [(2)(- 435.9) + (3)(0)] = - 782.4 + 871.8 = 89.4 kJ ΔHrxn = ? ΔH°f Example Problems ∆Hrxn = Σ n∆Hof Products - Σn∆Hof Reactants 3. AgNO3(s) + NaCl (aq) AgCl(s) + NaNO3(aq) ΔHrxn = ? (-124.4) (-407.1) (-127.0) (-446.2) ∆H = [(-127.0) + (-446.2)] – [(-124.4) + (-407.1)] = -573.2 + 531.5 = - 41.7 kJ 4. C2H5OH(l) + 7/2 O2(g) 2CO2(g) + 3H2O(g) ΔHrxn = ? (-277.7) (7/2)(0) (2)(-393.5) (3)(-241.8) ∆H = [(2)(-393.5) + (3)(-241.8)] – [(-277.7) + (7/2)(0)] = -1512.4 + 277.7 = -1234.7 kJ Enthalpy Review #2. Calculate DH for the following reaction: N2H4 (g) + O2 (g) → N2 (g) + 2H2O (g) Given: DH (kJ/mol) 2 NH3 (g) + 3 N2O (g) → 4 N2 (g) + 3 H2O (g) -1010 N2O (g) + 3 H2 (g) → N2H4 (g) + H2O (g) -317 2 NH3 (g) + ½ O2 (g) → N2H4 (g) + H2O (g) -143 H2 (g) + ½ O2 (g) → H2O (g) -286 Bond Energies • Chemical reaction ⇔ Bond breakage & bond formation • Bond energy = energy required to break a bond • Bond breaking is endothermic (raises potential energy) • Bond formation is exothermic (lowers PE) • Average energy for one type of bond in different molecules • Common Bond Energies C-H : 413 kJ/mol C=O : 799 kJ/mol O=O : 495 kJ/mol O-H : 467 kJ/mol • ΔHrxn = (bonds broken) – (bonds formed) Energy required Energy released Bond Energies • ΔHrxn = (bonds broken) – (bonds formed) ex. CH4 + H H C H + H 2 O2 CO2 + 2 H2O O O O O C-H : 413 kJ/mol O=O : 495 kJ/mol O C + O H H O O H H C=O : 799 kJ/mol O-H : 467 kJ/mol ΔHrxn = [4(C-H) + 2(O=O)] – [2(C=O) + 4(O-H)] ΔHrxn = [4(413 kJ) + 2(495 kJ)] – [2(799 kJ) + 4(467 kJ)] ΔHrxn = -824 kJ Compare to ΔHrxn = -802.5 kJ Enthalpy Summary • Enthalpy (ΔH) rxn = heat = q • All rxns have some ΔH • ΔH = +…endo • ΔH = - …exo • If given ΔH, can use stoich to quantify • Three ways to estimate ΔH (if not given) • Hess’s Law • known eq’s manipulated into desired eq. ΔHrxn • Heats of Formation (ΔHf): values from appendix • ΔHrxn = nΔH°f(products) - nΔH°f(reactants) • Bond energies • ΔHrxn = bonds broken – bonds formed