Unit 2 K sp and Equilibrium

1. Unit 2Ksp and Equilibrium • By: Michael Nolan • Kristi Rice • Erika Baucom

2. Important Equations

3. Bonds • Van der Waals • London Dispersion • Dipole-Dipole • Hydrogen Bonds (N,O,F)

4. Solids • Crystalline vs. amorphous • Crystalline- ionic and network (diamond) • Amorphous- glass and rubber • Unit Cell • Simple Cubic- just corners (1/8), one atom • Body-centered- corners and middle (1), two atoms • Face-centered- corners and faces (1/2), four atoms

5. Crystalline Structures Simple cubic Body-centered Face-centered (Prentice Hall Student Component CD)

6. Example • Sodium Chloride has a face-centered crystal structure. How many Cl– ions • and how many Na+ ions? Also, • what is the empirical formula? • MgO has the same crystal • structure and a density of • 3.58 g/cm3. What is the length of one edge?

7. Solubility • The amount of the substance that can be dissolved in a given quantity of solvent at a given temperature • Soluble solution: clear solution, light passes straight through b/c particles are about same size as light waves • Insoluble solution: cloudy solution, light path is scattered b/c particles larger than light waves

8. Solubility (continued) • Saturated: solution is in equilibrium with undissolved solute so additional solute won’t dissolve • Solubility vs. temp. graphs for gases and solids on the board • Unsaturated: can dissolve more solute • Supersaturated: contains more solute than needed to form a saturated solution; unstable; agitation can spur precipitates

9. Equilibrium Equations • Equilibrium Constant Equation: • Kp = Kc(RT)Δn Δn = product coeff. – reactant coeff.

10. Ksp and Solubility • Ksp is equilibrium constant for dissolved ionic substances • Solubility is affected by: • Henry’s law • Sg = kHPg • Solubility of a gas is directly proportional to partial pressure • Temperature • Inversely proportional with gas solubility • Directly proportional with solid solubility • Surface Area • Polarity of solvent vs. solute; like dissolves like • pH - Example: basic anions are more easily dissolved in acidic solutions

11. Reaction Equation Types • Molecular Equation: • Pb(NO3)2 (aq) + 2KI (aq) → PbI2 (s) + 2KNO3 (aq) • Complete Ionic Equation: • Pb2+ (aq) + 2NO3– (aq) + 2K+ (aq) + 2I– → PbI2 (s) + 2K+ (aq) + 2NO3– (aq) • Net Ionic Equation: • Pb2+ (aq) + 2I– (aq) → PbI2 (s)

12. Basic Solubility Terms • Dilute - to lower a concentration by adding water • Miscibility - the amount of liquid that can be dissolved in another liquid; “dissolvability” of liquids • Saturated - solution is in equilibrium with undissolved solute so added solute will not dissolve • Equilibrium - when opposing reactions proceed at equal rates and have constant concentrations (“Dynamic” Equilibrium)

13. Reaction Quotient • Reaction Quotient: • if Q<K, shift right • if Q>K, shift left • if Q=K, equilibrium is established

14. Example • Will AgIO3 (Ksp = 3.1 x 10–8) precipitate when 20 mL of 0.10 M AgNO3 is mixed with 10 mL of 0.015 M NaIO3? • Suppose a solution contains 0.10 M Ca(NO3)2 and 0.10 M Ba(NO3)2.The cations are to be separated by adding NaF to form CaF2 (Ksp=3.9 x 10–11) and BaF2 (Ksp=1.7 x 10–6). Which precipitate will form first? What would be the F– concentration at that point?

15. Le Châtlier’s Principle • If a system at equilibrium is disturbed by a change, the system will shift its equilibrium position to counteract the effect of the disturbance • Examples: • Increase Temperature: shift to reaction with +ΔH • Decrease Temperature: shift to reaction with –ΔH • Increase Pressure/Decrease Volume: shift to side with fewer moles of gas to decrease pressure • Add Catalyst: No Shift, catalyst increases both forward and reverse reaction rates

16. Solubility and Equilibrium Calculations • Molarity = moles = M L • Molality = moles solute = m kg solvent • Mass percentage = mass component x 100% total mass • ppm: ______ x 10⁶ • Volume fraction = V solute V total • Mole fraction = X = moles component moles total

17. Example • Wine contains 12.5% ethanol by volume. It’s density is 0.789 g/cm3. Calculate the mass percent and molal concentration of alcohol.

18. Common Ion Effect • Common ions in a solution suppress the reactant’s solubility and push the equilibrium to the left.

19. Complex Ion Formation • A complex ion is the assembly of a metal ion and the lewis bases bonded to it. • Kf is the equilibrium constant of its formation with the hydrated metal ion. • Example: • Fe2+ + 6CN– → [Fe(CN)6]4– • Kf = 7.7 x 1036 • Knet = Ksp x Kf

20. Example • FeCO3 has Ksp = 3.5 x 10–11. [Fe(CN)6]4– has a • Kf = 7.7 x 1036. • a.) Combine 2 reaction equations to form a net reaction describing a system in which concentrated KCN is added to an excess of solid FeCO3 in H2O. • b.) Calculate Knet and write a general expression for it. • c.) Calculate the theoretical solubility of FeCO3 in 2.50 M KCN.

21. Phase Changes Vocabulary • Vapor Pressure - pressure of a vapor in contact with its liquid or solid form; boiling point is when Pvap=Patm • Volatile - evaporates readily • Phase Diagrams - pressure vs. temperature graphs that show where the states of matter exist at equilibrium • Colligative Properties - depend on the quantity or concentration, but not the identity, of solute particles • Boiling Point Elevation: ΔTb = ikbcmolal • Freezing Point Depression: ΔTf = ikfcmolal • Osmotic Pressure - pressure required to prevent osmosis by pure solvent toward the solution with the higher solute concentration • ∏ = cmolarRT • i - Van’t Hoff Factor; number of products solute dissociates into (“ideally”; “real” Van’t Hoff factor is lower)

22. B.P. Elevation and F.P. Depression • Boiling Point Elevation - the ions dissolved in a solvent hinder the solvent molecules from vaporizing; lowers the solvent’s vapor pressure, so more heat is necessary for Pvap to equal Patm and boil • Freezing Point Depression - the ions dissolved in a solvent physically hinder the solvent’s molecules from creating an orderly crystalline arrangement (StickerGiant.com) (http://www.avromysegalonline.com/wp-content/uploads/2010/08/sad-smiley.jpg)

23. Example • A polyprotic acid w/ molar mass of 103.2 g/mol is dissolved in H2O. Suppose 74.3 g of the acid is dissolved in 250. g H2O, and the resulting solution is heated to its boiling point. H2O’s b.p. elevation constant, kb, is 0.512°C-kg/mol. If the first proton dissociates 100% and the remaining protons don’t dissociate, what is the boiling point of the solution?

24. Phase Diagrams • Triple Point - at which all three phases are in equilibrium • Critical Point - Critical Temperature and Critical Pressure • Critical Temperature - highest temperature at which there is a distinct liquid phase • Critical Temperature - pressure required to condense substance at critical temp. • Supercritical Fluid - indistinguishable liquid/gas phase supercritical fluid critical point (Science Fair Projects Encyclopedia; http://www.all-science-fair-projects.com/science_fair_projects_encyclopedia/Phase_(matter) )