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Spontaneity of Reaction

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Spontaneity of Reaction

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  1. Spontaneity of Reaction Chapter 16

  2. Key Concepts of Chapter 16 • Identifying Spontaneous Processes. • Identifying reversible and irreversible processes. • Entropy and its relation to randomness. • Second Law of Thermodynamics. • Predicting Entropy Changes of a Process. • Third Law of Thermodynamics. • Relate temperature change to entropy change. • Calculating change in standard entropy.

  3. • Free energy in terms of enthalpy and entropy. • Relating free energy change to spontaneity. • Calculating standard free energy change. • Relationship between free energy and work. • Calculating free energy Δ under nonstandard conditions.

  4. From Chapter 8 Chemical thermodynamicsis the study of energy relationships in chemistry. The First law of Thermodynamics - energy cannot be created or destroyed only converted from one form to another.

  5. Enthalpy • heat transfer between the system and its surroundings under constant pressure. • Enthalpy is a guide to whether a reaction is likely to proceed. • It is not the only factor that determines whether a reaction proceeds.

  6. Spontaneous Processes • Occur without outside intervention • Have a definite direction. • The reverse process is not spontaneous. • Temperature has an impact on spontaneity. • Ex: Ice melting or forming • Ex: Hot metal cooling at room temp.

  7. KI (aq) + Pb(NO3)2 (aq)  PbI2(s) + KNO3 (aq) When mixed  Precipitate forms spontaneously. *It does not reverse itself and become two clear solutions.

  8. Reversible & Irreversible • Reversible: System changes state and can be restored by reversing original process. Ex: Water (s) Water (l) • Irreversible: System changes state and must take a different path to restore to original state. Ex: CH4 + O2 CO2 + H2O

  9. Whenever a system is in equilibrium, the reaction can go reversiblyto reactants or products (water  water vapor at 100 º C). • In a Spontaneous process, the path between reactants and products isirreversible. (Reverse of spontaneous process is not spontaneous). *Scrambled eggs don’t unscramble*

  10. The Second Law of Thermodynamics - The entropy of the universe always increases in a spontaneous process and remains unchanged in an equilibrium process.

  11. “But ma, it’s not my fault… the universe wants my room like this!”

  12. Entropy (S) • A measure of randomness or disorder • S = entropy in J/K·mole • Increasing disorder or increasing randomness is increasing entropy. • Three types of movement can lead to an increase in randomness.

  13. Entropy is a state function • Change in entropy of a system • S= Sfinal- Sinitial • Depends only on initial and final states, and not the pathway. -S indicates a more ordered state(think: < disorder or - disorder) Positive (+) S = less ordered state (think: > disorder or + disorder)

  14. Entropy, S - a measure of disorder Ssolid  Sliquid  Sgas

  15. Increasing Entropy

  16. Increasing Entropy

  17. Increasing Entropy

  18. If entropy always increases, how can we account for the fact that water spontaneously freezes when placed in the freezer? • Movement of compressor + • Evaporation and condensation of refrigerant + • Warming of air around container Net increase in the entropy of the universe

  19. On the AP exam, you will likely be asked to: • predict whether a process leads to an increase in entropy or a decrease in entropy. • Determine if ΔS is + or – • Determine substances or reactions that have the highest entropy.

  20. Processes that lead to an Increase in Entropy • When a solid melts. • When a solid dissolves in solution. • When a solid or liquid becomes a gas. • When the temperature of a substance increases. • When a gaseous reaction produces more molecules. • If no net change in # of gas molecules, can be + or -, but small.

  21. Predict whether the entropy change is greater than or less than zero for each of the following processes: • Freezing liquid bromine • Evaporating a beaker of ethanol at room temperature • Dissolving sucrose in water • Cooling N2 from 80ºC to 20ºC S<0 S>0 S>0 S<0

  22. Predict whether the entropy change of the system in each of the following reactions is positive or negative: 1)S – 2)S+ 3)S? 1.) Ag+(aq)+ Cl-(aq)AgCl(s) 2.) NH4Cl(s) NH3(g)+ HCl(g) 3.) H2(g) + Br2(g)2HBr(g)

  23. According to the 2nd law of thermodynamics; the entropy of the universe always increases. ? What if the entire senior class assembles in the auditorium? Aren’t we decreasing disorder, and therefore decreasing entropy? If so, how can the second law of thermodynamics be true?

  24. If we consider the senior class as the system, the • ΔS of the system would indeed decrease. • ΔS of the system is – • In order for the students to gather, they would: • Metabolize food (entropy increase of surroundings) • Generate heat (entropy increase of surroundings) • The magnitude of the entropy increase of the surroundings will • always be greater than the entropy decrease of the system.

  25. Theoretical values Suniverse = Ssystem+Ssurroundings Suniverse = (-10) + (+20) Suniverse = +10 + means entropy increases

  26. The same can be considered in a chemical process. When a piece of metal rusts: 4Fe(s) + O2(g)  2Fe2O3(s) The entropy of the solid slowly decreases. Although this is a slow process, it is exothermic, and heat is released into the surroundings causing an overall increase in entropy of the universe!

  27. nonspontaneous spontaneous Spontaneous Physical and Chemical Processes • A waterfall runs downhill • A lump of sugar dissolves in a cup of coffee • At 1 atm, water freezes below 0 0C and ice melts above 0 0C • Heat flows from a hotter object to a colder object • A gas expands in an evacuated bulb • Iron exposed to oxygen and water forms rust

  28. spontaneous nonspontaneous

  29. CH4(g) + 2O2(g) CO2(g) + 2H2O (l)DH0 = -890.4 kJ H+(aq) + OH-(aq) H2O (l)DH0 = -56.2 kJ H2O (s) H2O (l)DH0 = 6.01 kJ H2O NH4NO3(s) NH4+(aq) + NO3-(aq)DH0 = 25 kJ Does a decrease in enthalpy mean a reaction proceeds spontaneously? Spontaneous reactions at 25 °C

  30. TWO Trends in Nature • Order  Disorder   • High energy  Low energy 

  31. The Driving Forces • Energy Factor (ΔH) • Randomness Factor (ΔS) or more lately the “dispersal” factor.

  32. S order disorder S H2O (s) H2O (l) Entropy (S) is a measure of the randomness or disorder of a system…the tendency to spread energy out – disperse energy. DS = Sf - Si If the change from initial to final results in an increase in randomness DS > 0 Sf > Si For any substance, the solid state is more ordered than the liquid state and the liquid state is more ordered than gas state Ssolid < Sliquid << Sgas DS > 0

  33. First Law of Thermodynamics Energy can be converted from one form to another but energy cannot be created or destroyed. Second Law of Thermodynamics The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process. DSuniv = DSsys + DSsurr > 0 Spontaneous process: DSuniv = DSsys + DSsurr = 0 Equilibrium process:

  34. The Second Law of Thermodynamics - The entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process. Suniverse = Ssystem +Ssurroundings Suniverse > 0 for spontaneous rxn Suniverse = 0 at equilibrium

  35. In terms of temperature, how would you describe an object that has an entropy value of 0? 0 K Perfect solid crystal with no motion Only Theoretical It is not possible to reach absolute 0! Entropy of universe is always increasing!

  36. 3rd Law of Thermodynamics the entropy of a perfect crystalline substance is zero at absolute zero *Based on 0 entropy as a reference point, and calculations involving calculus beyond the scope of this course, data has been tabulated for Standard Molar Entropies ΔSº Pure substances, 1 atm pressure, 298 K

  37. Standard Molar Entropies • ΔSº • Standard molar entropies of elements are not 0 (unlike ΔHºf). • (0 entropy is only theoretical; not really possible) • S.M.E of gases > S.M.E of liquids and solids. • (gases move faster than liquids) • 3) S.M.E. increase with increasing molar mass. • (more potential vibrational freedom with more mass) • 4) S.M.E. increase as the number of atoms in a formula increase. • (same as above)

  38. Calculating the Entropy Change Sorxn =n So(products) - m So(reactants) Units for S S=J/mol•K Since we are considering ΔS° J/K are often used because moles are assumed and cancel in the calculations when considering standard states.

  39. Calculate the standard entropy change (Sº) for the following reaction at 298K Al2O3(s) + 3H2(g)2Al(s) + 3H2O(g)

  40. Sorxn =n So(products) - m So(reactants) So = [2Sº(Al) + 3Sº(H2O)] - [Sº(Al2O3) + 3Sº(H2)] Al2O3(s) + 3H2(g)2Al(s) + 3H2O(g) = 180.4 J/K

  41. Predict the sign of ΔSº of the following reaction. 2SO2(g) + O2(g) 2SO3(g) Entropy decreases, - Lets’ Calculate

  42. Calculate the standard entropy change (Sº) for the following reaction at 298K 2SO2(g) + O2(g) 2SO3(g) ΔSº = -187.8 J K-1

  43. Predicting spontaneous reactions • Spontaneous reactions result in an increase in entropy in the universe. • Rx’s that have a large and negative  tend to occur spontaneously. • Spontaneity depends on enthalpy, entropy, and temperature.

  44. Gibbs Free Energy (G) Provides a way to predict the spontaneity of a reaction using a combination of enthalpy and entropy of a reaction.

  45. If Both T and P are constant, the relationship between G and spontaneity is: • G is (-), forward rxn is spontaneous. • G is 0, rxn is at equilibrium. • G is (+) forward rxn is not spontaneous (requires work) reverse rxn is spontaneous.