1. Chem. 1B – 10/15 Lecture

2. Announcements I Exam 2: Two Weeks from Today (10/29) Lab: Experiment 4 Report – due Mon./Tues. Today’s Lecture –Complex Ion Formation Why study? Effects on Solubility –Thermodynamics Reviewing Ch. 6 Entropy and Change in entropy

3. Chem 1B – Aqueous Chemistry Complex Ion Formation (Chapter 16) Complex Ions – Why do we study? –Metals bound as complexes do not react the same as “free” metals –In determining solubility, for example, only the “free” metal is in the equilibrium equation Example: AgCl(s) ↔ Ag + (aq) + Cl - (aq) or K = [Ag + ][Cl - ] – Ag complexed as Ag(NH 3 ) 2 + ≠ Ag + –Complexation also affects reactivity and uptake For example: spinach is high in Fe, but also in oxalate (C 2 O 4 2- ) which complexes Fe (K f = 2 x 10 20 for Fe 3+ ) making uptake more difficult. Additionally, oxalate potentially can bind Fe from the body.

4. Chem 1B – Aqueous Chemistry Complex Ion Formation (Chapter 16) Complex Ions – Why do we study? –Some uses of complex ions (besides for increasing solubility) Water hardness titration – done in experiment 6 Separations – complexed metals are more organic soluble and it is possible to move metals to a separate liquid phase such as an ether phase Avoiding oxidation (for example Fe 3+ + H 2 O 2 produces strong oxidants, unless Fe 3+ is bound)

5. Chem 1B – Aqueous Chemistry Complex Ion Formation (Chapter 16) Complex Ions – Effects on Solubility –Examples: 2) Ca 2+ + C 2 O 4 2- (oxalate anion) – anion can lead to both precipitation and complex formation solubility rxn: CaC 2 O 4 (s) ↔ Ca 2+ (aq) + C 2 O 4 2- (aq) 1.3 x 10 -8 complex rxn: Ca 2+ (aq) + 2C 2 O 4 2- (aq) ↔ Ca(C 2 O 4 ) 2 2- (aq) K f = 2.3 x 10 4 - Because 2 oxalates are used for complexation (vs. 1 released by dissolution), complexation is more important at higher concentrations At low [C 2 O 4 2- ] (e.g. 1 x 10 -7 M at equilibrium), Ca 2+ is fairly soluble (0.13 M) with almost no complex forms ([Ca(C 2 O 4 ) 2 2- ] = 3 x 10 -13 M) At moderate [C 2 O 4 2- ] (e.g. 1.0 x 10 -3 M), Ca 2+ is less soluble (1.3 x 10 -5 M), and complex at similar (3 x 10 -7 M) concentration At high [C 2 O 4 2- ] (e.g. 0.5 M), very little Ca 2+ is present (2.6 x 10 -8 M), but complex starts to increase net solubility (1.5 x 10 -4 M)

6. Complex Ions – “ U ” Shaped Solubility Curves Solubility in water Common ion effect Complex ion effect Note: looks “U” shaped if not on log scale (otherwise “V” shaped)

7. Chem 1B – Aqueous Chemistry Complex Ion Formation (Chapter 16) Complex Ions – Effects on Solubility –The second example also applies to metal hydroxides (e.g. Zn(OH) 2 = sparingly soluble salt, but solubility increases at high pH due to formation of Zn(OH) 4 2- ) –Calculate the solubility of Zn 2+ in buffers at pH = 7, 10 and 13. K sp (Zn(OH) 2 ) = 3 x 10 -17 and K f (Zn(OH) 4 2- ) = 2 x 10 15

8. Chem 1B – Thermodynamics Chapter 17 Chapter 6 – Review –Types of Energy: kinetic energy (associated with motion) potential energy (stored energy – e.g. ball at the top of a hill) Chemical energy (a type of stored energy) Heat (a molecular scale type of kinetic energy) –Conservation of Energy Energy can change forms – but can not be created or destroyed

9. Chem 1B – Thermodynamics Chapter 17 Chapter 6 – Review II –Systems and Surroundings used to define energy transfers example: system with reaction that produces heat (from chemical energy) can heat surroundings –Enthalpy (H) Energy related to heat  H = q p (heat in a constant pressure system) Endothermic reaction means  H > 0, means heat from surrounding used for reaction Exothermic reaction means  H > 0, means heat from reaction goes to surroundings

10. Chem 1B – Thermodynamics Chapter 17 Chapter 17 – Overview –Spontaneous and Non-Spontaneous Processes: –Entropy: A measure of disorder –Gibbs Free Energy –Entropy and Gibbs Free Energy Changes Associated with Reactions –Relating Gibbs Free Energy to Equilibrium Constants

11. Chem 1B – Thermodynamics Chapter 17 – Spontaneous Processes Thermodynamical Definition of Spontaneous –A spontaneous process is one that will eventually occur (actually has nothing to do with speed of occurrance) –Examples of spontaneous processes: freezing of water droplet at -5°C dissolution of 0.1 moles of NH 4 NO 3 in 1.0 L of water (solubility is much higher) oxidation of Fe(s) in air

12. Chem 1B – Thermodynamics Chapter 17 – Spontaneous Processes Thermodynamical Definition of Spontaneous –Most, but not all, spontaneous processes are exothermic (e.g. H 2 (g) + O 2 (g) ↔ H 2 O(l)) –Non-spontaneous process: one that won’t occur without intervention –Example: splitting water to H 2 (g) and O 2 (g) (can be done through electrolysis, but then needs external energy)

13. Chem 1B – Thermodynamics Chapter 17 – Entropy Entropy –A few reactions that occur spontaneously are endothermic (e.g. NH 4 NO 3 (s) ↔ NH 4 + (aq) + NO 3 - (aq)) –How can a process occur if it takes energy? –There must be some trade off that makes it likely to occur –Trade off is an increase in disorder (entropy) –For example, we can see that a desk will have a natural tendency to becoming messy and that it takes energy to clean it

14. Chem 1B – Thermodynamics Chapter 17 – Entropy Entropy –A macroscopic analogy to entropy would be to have a box of 50 ping pong balls with half white and half black –Even if placed on two separate halves of the box, if the box were shaken to mix the balls, roughly half of each color would be expected in each half v initial state v final state

15. Chem 1B – Thermodynamics Chapter 17 – Entropy Entropy in Chemical Systems –From a molecular scale view, a system that appears more randomly assembled has higher entropy (can have more possible “states”) Highly orderedHighly disordered Low EntropyHigh Entropy Crystalline solid (T = 0K) Amorphous solidliquid gas large compound various small compounds N 2 O 4 (g) vs.2NO 2 (g) vs.2O 2 (g) + N 2 (g) vs.4O (g) + 2N(g) S = 0 note: gases shown are relative (still at higher entropy vs. liquids/solids Crystalline solid (T > 0K)

16. Chem 1B – Thermodynamics Chapter 17 – Entropy Determine the sign of entropy change for the following reactions: Entropy Examples: (Is ΔS > or < 0?) H 2 O(l) ↔ H 2 O(g) H 2 O(s) ↔ H 2 O(l) NaCl(s) ↔ Na + (aq) + Cl - (aq) 2H 2 (g) + O 2 (g) ↔ 2H 2 O(g) N 2 (g) + O 2 (g) ↔ 2NO(g) ΔS > 0 ΔS < 0 ΔS > 0

17. Chem 1B – Thermodynamics Chapter 17 – Entropy Quantifying Entropy Changes ( Δ S) –From 2 nd Law of thermodynamics, we know  S univ ≥ 0 and  S univ =  S sys +  S surr –Thus for a process in which Δ S sys 0 –In our particular example, energy (or enthalpy) is evolved in the process (I 2 (g) ↔ I 2 (s)) –we can set these equal in  S surr = -q sys /T or under constant pressure,  S surr = -  H sys /T

18. Chem 1B – Thermodynamics Chapter 17 – Entropy Quantifying Entropy Changes ( Δ S) II –For a particular process, we also can look up standard entropy values (S°) for reactants and products (see Appendix II B) –These are under standard conditions (25°C/1 atm for gases/1 M for solutions)