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MBB 222 Lecture Notes - Lecture 7: Gibbs Free Energy, Thermodynamics, Equilibrium Point

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browncricket987
20 Dec 2018
School
Simon Fraser University
Department
Molec Biol & Biochem
Course
MBB 222
Professor
Adam Barlev

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Document Summary

From last time: g= h-t s, g = g " + rt ln [products]/[reactants, if g = 0, G " = -rt ln [products]eq/[reactants]eq: g " is defined based on initial standard state. Standard temp and pressure, all reactants at 1m except h+ which starts at ph 7. A more concrete example: ice liquid transition for water, role of t. Ice requires heat to melt ( h is positive) melting is enthalpically unfavorable. Water molecules become more disordered upon melting (gain entropy, s is positive) melting is entropically favorable. Whether or not the ice will melt depends on the temperature. For melting to proceed, t must be high enough for the t s term to become larger than h, making the g term negative. Neither h nor s alone will tell us what will happen, but the combination ( g) tells us what form water will take at a given temperature; a positive g = ice, but a negative g.

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Related Questions


I need help on this chemistry work. Thanks.

DATA ANALYSIS

Using the example, calculate the Gibbs free energy for the process of dissolving ammonium nitrate and the process of dissolving sodium hydroxide (equation: ΔG = ΔH – TΔS).

Use the standard temperature 298 K.

Use the experimental data to calculate enthalpy of dissolving (ΔH).

Use the following data to calculate the entropy change (ΔS):

S (J/K mol)

NH4NO3(s)

151

NH4+(aq)

113.4

NO3-(aq)

146.4

NaOH(s)

64.46

Na+(aq)

59.0

OH-(aq)

- 10.7

After completing the calculations answer the following questions separately for each experiment:

Is the dissolving exothermic or endothermic?

Does the entropy decreases or increases during dissolving?

Is the Gibbs free energy positive or negative? Is the process of dissolving spontaneous?

Using data in the appendix in your textbook, calculate Gibbs free energy for the reaction of dissolving

lithium phosphate, Li3PO4, in water at 25 °C. Do the calculations in two ways:

1) Use standard enthalpy of formation (ΔH) and entropy (S) from the appendix in your textbook, and use the same formula as you did in the lab experiments: ΔG = ΔH – TΔS. Remember that both ΔH and ΔS has to include three lithium ions.

2) Use standard free energies of formation from Appendix 6 (ΔG) and equation

ΔG = ∑ΔG products – ∑ΔG reactants

Are the results similar in both calculations?

Q1. Which of the following processes are spontaneous and which are non spontaneous.
a) The combustion of natural gas
b) Extraction of Aluminum metal from its ore
c) A ball rolling down a hill
d) A bike going up a hill


Q2. Predict the entropy change (positive or negative) for the following processes/reactions
a) a lake freezing
b) weeding a garden
c) N2(g) + 2 H2(g) ? N2H4 (g)
d) Mg(s) + 2H2O(l) ? Mg(OH)2 (s) + H2(g)


Q3. The element Cesium (Cs) freezes at 28.4 C and its molar enthalpy of fusion ΔHfusion = 2.09 kJ/mol.
a) When molten Cs solidifies to Cs (s) at its normal melting point, is the entropy change ΔS positive or negative?
b) Calculate the entropy change S when 15.0 g of Cs(l) solidifies at 28.4 C.

Q4. Using the values for standard molar entropies ΔSf° from the appendix in your textbook, calculate ΔS° for the following reactions.
a) Be(OH)2 (s) ? BeO (s) + H2O (g)
b) Mg(OH)2 (s) + 2 HCl (g) ? MgCl2 (s) + 2H2O (l)
c) 2CH4(g) ? C2H6 (g) + H2(g)


Q5. Using the data in the appendix of your textbook, calculate the Gibbs free energy change ΔG° for the following reactions. In each case, indicate whether the reaction is spontaneous or not
a) H2 (g) + Cl2(g) ? 2HCl (g)
b) 2 POCl3 (g) ? 2 PCl3 (g) + O2 (g)


Q6. Consider the reaction 2 NO2 (g)---N2O4 (g)
a) Using the data in the appendix in your text, calculate the standard Gibbs free energy change (ΔG°) for the reaction at 298K.
b) For the above reaction, calculate the value of Keq at 298K
c) For the same reaction, calculate G at 298 K when the partial pressures for NO2 and N2O4 are 0.40 atm and 0.36 atm, respectively.


Q7. Consider the reaction I2 (g) + Cl2 (g)----2 ICl (g) Kp = 81.9 at 298 K
a) Calculate the standard free energy change (ΔG°)
b) What will be the Free energy change when the above reaction is at equilibrium
c) Calculate G for the reaction at 298 K when PICl = 2.55 atm, PI2=0.325 atm and PCl2 = 0.221 atm


Q8. Consider the decomposition of Lead carbonate PbCO3 (s) ? PbO(s) + CO2 (g)
a) Calculate the Gibbs free energy change (ΔG°) for the reaction at298 K.
b) Calculate G for the reaction at 400 K.
c) Write down the equilibrium expression and calculate the value of Keq at 298 K
d) Calculate the partial pressure of CO2 at 298 K. (Hint: Look at the equilibrium expression and the value of Keq at 298 K)


Q9. For the reaction: 2Al(s) + 3/2 O2(g) ? Al2O3(s), ΔHorxn = -1617 kJ/mol and ΔHorxn = -1577 kJ/mol. Use this data to calculate ΔSorxn reaction at 298 K.


Q10. Calculate the ΔGo for the following reaction using the data from your text:
2SO2(g) + O2(g) ? 2SO3(g)


Q11. Using the thermodynamic data available from the Appendix(?Go’s), calculate ΔGo rxn and calculate the equilibrium constant(K) for the reaction at 25 oC:
CaCO3(s) ? CaO(s) + CO2(g)


Q12. What is the minimum temperature required for the spontaneous conversion of CCl4(l) to CCl4(g) when ΔH°(vap) is 57.3 kJ/mol andΔS°(vap) is 164 J/(mol K)? reaction: CCl4(l) ? CCl4(g)


Q13. Given the following
Fe2O3(s) + 3CO(g) . ? 2Fe(s) + 3CO2(g) ΔG° =  29.4 kJ
3Fe2O3(s) + CO(g) . ? 2Fe3O4(s) + CO2(g) ΔG° =  61.6 kJ
Calculate ΔG° for Fe(s) + Fe2O3(s) + CO2(g) ? Fe3O4(s) +CO(g)


Q14. Calculate ΔG° rxn for 2CO(g) + O2(g) ? 2CO2(g) and then calculate ΔG for the reaction at 298K; if the pressure of CO =20. atm,P O2 = 10. atm and P CO2= 0.50 atm. Will increasing the pressure of the reactants make the reaction more or less spontaneous? Explain using your calculations and Le Chatelier’s principle.


Q15. Write the reaction for photosynthesis (production ofC6H12O6(s) and oxygen from carbon dioxide and water as a gas).Calculate ?H, ? S, and ?G for the reaction at standard state. Is the reaction spontaneous?