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Chem. 1B – 10/11 Lecture

Explore solubility, pH effects, thermodynamics, and complex ions in aqueous chemistry. Discover how acids influence solubility, along with precipitation and ion selective removal methods. Dive into complex ion formations and their impact on solubility shifts. Learn about practical applications in analytical chemistry and water titrations. Encounter "U"-shaped solubility curves and the intriguing effects of metal hydroxides on solubility in various pH environments.

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Chem. 1B – 10/11 Lecture

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  1. Chem. 1B – 10/11 Lecture

  2. Announcements • Mastering (Last Ch. 16 assignment – sparingly soluble solids and complex ions – due Saturday) • Today’s Lecture • Solubility (pH effects on solubility, precipitation, and selective precipitation for separations) • Complex Ions • Thermodynamics (Chapter 17 – starting with review from Chem 1A)

  3. Chem 1B – Aqueous ChemistrySolubility (Chapter 16) • Effect of pH on Solubility • Which of the following salts have solubility increase by addition of acids? • AgCl - Mg3(PO4)2 • Mg(OH)2 - BaSO4 • Hg2Br2 RULE: • If conjugate base is a strong or weak base, acid addition increases solubility • If conjugate base is neutral (conjugate to strong acid), no effect on solubility

  4. Chem 1B – Aqueous ChemistrySolubility (Chapter 16) • Another Example • “Ocean Acidification” Effect of Carbon Dioxide • Increase of CO2 in the atmosphere leads to an increase in H2CO3 (an acid) in oceans (even if effect is to make ocean less basic) • This leads to both increasing (due to acid addition from H2CO3) and decreasing (due to common ion effect of added CO32-) solubility of CaCO3 • This matters because ocean creatures have CaCO3 containing structures (e.g. shells) • Which effect matters more? • Ka1 = 4.3 x 10-7 Ka2 = 5.6 x 10-11 and Ksp = 4.96 x 10-9 • Compare: CaCO3(s) + H2CO3(aq) ↔ Ca2+(aq) + 2HCO3- and Ca2+(aq) + H2CO3(aq)↔CaCO3(s) + 2H+(aq)

  5. Chem 1B – Aqueous ChemistrySolubility (Chapter 16) • Precipitation • If we mix an ion pair from a sparingly soluble salt together, how do we know if precipitation will occur? • Example: 0.040 M Ca2+ + 0.0010 M F- – do we get CaF2 (s)? • Using Ksp reaction (backwards of what is occurring), we can compare Q with Ksp • If Q < Ksp, no precipitation occurs (solution stays clear) • If Q > Ksp, either precipitation occurs or supersaturated solution [Example problem]

  6. Chem 1B – Aqueous ChemistrySolubility (Chapter 16) Precipitation for selective ion removal Example: An old battery plant had a leak of lead and sulfuric acid that was neutralized by addition of CaCO3. The collected liquid has [Ca2+] = 0.20 M and [Pb2+] = 0.10 M. A chemist wants to save the lead (in form Pb2+) but not the Ca2+. Can she selectively precipitate out >98% of the Pb2+ without precipitating Ca2+ by addition of SO42-? Ksp(CaSO4) = 7.10 x 10-5 and Ksp (PbSO4) = 1.82 x 10-8

  7. Chem 1B – Aqueous ChemistryComplex 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(NH3)2+≠ Ag+ Complexation also affects reactivity and uptake For example: spinach is high in Fe, but also in oxalate (C2O42-) which complexes Fe (Kf = 2 x 1020 for Fe3+) making uptake more difficult. Additionally, oxalate potentially can bind Fe from the body

  8. Chem 1B – Aqueous ChemistryComplex Ion Formation (Chapter 16) Complex Ions – Why do we study? – cont. Uses of complex ions in analytical chemistry (besides for increasing solubility) Water hardness titration – done in experiment 6 Separations – complexed metals are more organic soluble Reduction of oxidation (for example Fe3+ + H2O2 produces strong oxidants, unless Fe3+ is bound)

  9. Chem 1B – Aqueous ChemistryComplex Ion Formation (Chapter 16) Complex Ions – Effects on Solubility Examples: 2) Ca2+ + C2O42- (oxalate anion) – anion can lead to both precipitation and complex formation solubility rxn: CaC2O4(s) ↔ Ca2+(aq) + C2O42-(aq) 1.3 x 10-8 complex rxn: Ca2+(aq) + 2C2O42-(aq) ↔ Ca(C2O4)22-(aq) Kf = 2.3 x 104 At low [C2O42-] (e.g. 1.0 x 10-7 M at equilibrium), Ca2+ is fairly soluble (0.13 M) and almost no complex forms ([Ca(C2O4)22-] = 3 x 10-13 M and doesn’t contribute to solubility) At moderate [C2O42-] (e.g. 1.0 x 10-3 M), Ca2+ is less soluble (1.3 x 10-5 M), and complex still hasn’t formed much (3 x 10-7) to affect solubility At high [C2O42-] (e.g. 0.5 M), very little Ca2+ is present (2.6 x 10-8 M), but complex is increasing net solubility (1.5 x 10-4 M)

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

  11. Chem 1B – Aqueous ChemistryComplex 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)42-) Calculate the solubility of Zn2+ in buffers at pH = 7, 10 and 13. Ksp(Zn(OH)2) = 3 x 10-17 and Kf (Zn(OH)42-) = 2 x 1015

  12. Chem 1B – ThermodynamicsChapter 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

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

  14. Chem 1B – ThermodynamicsChapter 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

  15. Chem 1B – ThermodynamicsChapter 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 NH4NO3 in 1.0 L of water (solubility is much higher) oxidation of Fe(s) in air Non-spontaneous process: one that won’t occur without intervention

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