Enthalpy changes and calorimetry. Enthalpy changes in reactions Calorimetry and heat measurement Hess s Law Heats of formation

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1 Enthalpy changes and calorimetry Enthalpy changes in reactions Calorimetry and heat measurement Hess s Law Heats of formation

2 Learning objectives Describe the standard state for thermodynamic functions Explain sign of enthalpy change for changes of state Calculate enthalpy changes for reactions Use specific heat and heat capacity in calorimetric problems Apply Hess law to calculations of enthalpy change Use standard heat of formation in calculations of enthalpy of reaction Use bond energies to estimate heats of reaction

3 The importance of state The heat change in a reaction will depend upon the states of the reactants and/or products States must be specified ΔH(1) = kj 1 2. C3H8( g) 5O2 ( g) 3CO2( g) 4H O( g) ΔH(2) = kj 2. C3H8( g) 5O2 ( g) 3CO2( g) 4H 2O( l) Additional heat is released in (2) because condensation of H 2 O releases 176 kj

4 Thermodynamic standard state Heat changes will also depend upon P and T The Standard State: Most stable form of a substance at 1 atm pressure and 25ºC(usually); 1 M concentration for solutions Conventionally, enthalpy change under standard conditions is represented by: ΔHº Any thermodynamic function referring to standard conditions is identified by º

5 Calculating ΔU from ΔHº Reaction between N 2 and H 2 at P = 40 atm and ΔV 1.12 L What is ΔU? N 2 (g) + 3H 2 (g) = 2NH 3 (g) ΔHº = kj U 92.2kJ ΔU = ΔH PΔV J 1kJ Latm 1000J 40x 1.12Latm 101 ΔU > ΔH Internal energy change is less negative by amount of work done by surroundings on compressing system 4.52kJ 87. kj U 92.2kJ 7

6 Enthalpies of change of state Conversion of a substance from one state to another involves heat input or output Heat is absorbed when breaking bonds Solid liquid gas Heat is released when making bonds Gas liquid solid Heat of fusion, vaporization etc. (formerly known as latent heats heat consumed with no change in T)

7 State functions ignore path Transition from solid to gas may lie through liquid phase... Or transition may be direct sublimation Enthalpy change is the same by either route ΔH subl = ΔH fusion + ΔH vap

8 Heat of reaction Enthalpy change for a reaction is commonly known as heat of reaction Exothermic refers to reaction where heat is given out by system (we warm our hands): ΔH < 0 Endothermic refers to reaction where heat is absorbed by system and surroundings get cool (hands get cold): ΔH > 0 NOTE SIGNS OF ΔH!

9 Calculations with ΔHº ΔHº may be quoted for a given equation or for one mole of a reactant or product Either way, it can be used to calculate the heat change for a given amount of substance

10 Example What is the heat evolved from the combustion of 15.5 g of C 3 H 8? C3H8( g) 5O2 ( g) 3CO2( g) 4H 2O( l)... H 2219kJ / mol Combustion of one mole of C 3 H 8 yields kj Molar mass of C 3 H 8 = 44.0 g/mol No moles in 15.5 g = 15.5/44.0 = mol Heat produced = x = kj

11 Worked example of heat of combustion

12 Calorimetry: the universe in the palm of your hand Heat change in process measured by temperature change of water in calorimeter (surroundings) Insulation confines heat exchange to process and calorimeter universe becomes the calorimeter

13 Heat capacity and specific heat Heat capacity, C: Heat required to raise temperature of body by 1ºC (extensive) C T Specific heat (capacity), c (σ): Heat required to raise 1 gram of the substance by 1ºC (intensive) q ct m q

14 Interactive calorimetry exercise Calculate some heat changes using virtual calorimeter Heats of solution

15 Constant pressure and constant volume Heat capacity can be defined at constant pressure and constant volume C P q T C V For solids and liquids they are similar For gases the values differ by R (the ideal gas constant): C P = C V + R q T

16 Molar heat capacity Related to specific heat (sometimes known as specific heat capacity) is the molar heat capacity, C m : the heat required to raise the temperature of 1 mole of a substance by 1ºC q Cm T moles

17 Some examples The heat capacity of water is much larger than many other substances Consequence of strong H-bonding Important consequences for climate, the quality of life, and life itself Substance S.H. (J/g o C) S,H. (J/mol o C) Air Aluminum Copper Gold Iron Mercury NaCl Water (s) Water (l)

18 Hess s law and the many Enthalpy being a state function means that enthalpy change for process is independent of pathway Hess s law states that: Overall enthalpy change for reaction equals sum of enthalpy changes for individual steps pathways

19 Application: determining ΔH for unknown process in terms of ΔH for known process Consider the reaction N 2 g) 3H 2( g) 2NH 3( g)... ( H Proceeds in two steps N N 2 g) 2H2( g) N2H4( g)... ( H 2 H 4 g) H 2( g) 2NH 3( g)... ( H ΔH 1 = ΔH 2 + ΔH 3 Reaction 1 is sum of reactions

20 Obtain ΔH 2 from ΔH 1 and ΔH 3 ΔH 2, which is hard to determine experimentally, is obtained using ΔH 1 and ΔH 3 ΔH 2 is endothermic, the others are exothermic

21 Combustion of CH 4 by alternate Route 1: ΔH 1 complete combustion to CO 2 Route 2: ΔH 2 + ΔH 3 intermediate product is CO Different pathways, same enthalpy ΔH 1 = ΔH 2 + ΔH 3 routes

22 Standard heats of formation Standard heat of formation, ΔH fº : The enthalpy change for the formation of 1 mole of a substance in its standard state from it constituent elements in their standard states ΔHº f for elements in most stable state is by definition 0

23 Reactions are often hypothetical ΔHº f (CH 4 ) is kj/mol although direct reaction between C and H 2 does not occur Important to know which is the standard state for any substance ΔHº f values are tabulated for many substances

24 Heats of reaction, ΔH fº, and Hess s law Application of ΔHº f to calculation of heats of reaction (ΔHº RXN ) For reaction: aa + bb = cc + dd Products reactants H o reaction c H ( C) d H ( D) a H ( A) b H ( B) o o o o f f f f

25 Worked example of heat of reaction H rxn H f ( CO2 ) H f ( H2O) H f ( C3H8)

26 Estimating ΔHº f using bond energies What if ΔHº f value(s) for reaction is (are) unknown? Estimate enthalpy change as difference between bonds broken (energy cost) and bonds made (energy gain) ΔHº RXN = D(reactant bonds) D(product bonds) COST (Bond energies are positive values) GAIN

27 Consider reaction: CH 4( g) 3Cl2( g) CHCl 3( g) 3HCl ( g) On reactant side: 3 C-H bonds and 3 Cl-Cl bonds are broken On product side: 3 C-Cl bonds and 3 H-Cl bonds are made 3D 3D 3D D H 3 ClCl CH Reaction is exothermic H Cl H-Cl bond is much stronger than Cl-Cl CCl 3x243 3x410 3x432 3x330kJ kj H 327

28 Worked example of heat of reaction from bond energies

29 More than enthalpy Some reactions are exothermic Some reactions are endothermic Both kinds of reaction can occur spontaneously. Or not. Why? The enthalpy is not the arbiter of spontaneity There is more to energy than heat We must also consider entropy Stay tuned for the Second Law

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