E09. Exp 09 - Solubility. Solubility. Using Q. Solubility Equilibrium. This Weeks Experiment. Factors Effecting Solubility.

E09 Exp 09 - Solubility Solubility Solvation The reaction coefficient Precipitating Insoluble Substances Comparing Q to Ksp Solubility Equilibrium Solubility Product, Ksp Relating Molar Solubility Factors Effecting Solubility Common Ion Effect Solution ph 1 Using Q Predicting Solubility This Weeks Experiment The reactions Titration Next Meeting

Producing Solutions Solutions are produced by solute dissolving into a solvent. All ionic compounds dissolve in water to some degree. A solution is a homogenous mixture. A solvent is the largest component of the mixture. A solute is a smaller components of the mixture. 2

How Ionic Solids Dissolve in Water Salt dissolves in water. The the ions separate. Water molecules get in between the ions. The result is a mixture of ions and water. Mostly water. H 2 O Ions separating in solution is a process called dissociation. NaCl(s) Na + (aq) + Dissocia'on of sodium chloride in water Cl - (aq)

Dissociation is an Equilibrium Process H 2 O NaCl(aq) Na + (aq) + Dissocia'on of sodium chloride in water Cl - (aq) Dissolved ions in solution can find other dissolved ions. If the attraction between those ions is strong, they can re-associate. These dissolved ions form ion pairs. The ion pair is not a solid, it s still dissolved in solution. Ions that dissociate and re-associate in solution are a kind of reversible reaction.

Producing Solutions Solutions are produced by solute dissolving into a solvent. All ionic compounds dissolve in water to some degree. However, many compounds have such low solubility in water that we classify them as insoluble. MgCl2(aq) + Pb(OAc)2(aq) Mg(OAc)2(aq) + PbCl2(s) In the first semester of General Chemistry we introduce you to the solubility rules, for identifying substances classified as insoluble. 5

The Solubility Rules Check each step, in order. Solubility Rules you are responsible for. Soluble no precipitate Insoluble forms precipitate 2 + Step 1 Step 2 Step 3 Step 4 has exceptions ANIONS CA+IONS ANIONS ANIONS Acetates (OAc 1- or CH3COO 1- ) Nitrates (NO3 1- ) Ammonium (NH4 1+ ) Alkali metal (Na 1+, Li 1+, K 1+...) Acids (the ones we learned) Carbonates (CO3 2- ) Phosphates (PO4 3- ) Halogens (Cl 1-, Br 1-, I 1-, F 1- ) Sulfates (SO4 2- ) Sulfides (S 2- ) Hydroxy Salts (OH 1- ) Always Always Never Usually Usually Except: Sr 2+, Ba 2+, Ca 2+ Never Never Always Except: Ag+, Hg2 2+ or Pb 2+ Hg2 2+ or Pb 2+ Sr 2+, Ba 2+ Usually Hg2 2+ mercury (I) ion 2 + Hg 2+ mercury (II) ion If you remember 1-3 you ll be good 85% of the time If you remember 1-3 and 4 you ll be good 95% Remembering the exceptions isn t that hard there s only six ions that cause exceptions and lead, mercury, and silver are the most commonly encountered ones.

Producing Solutions Solutions are produced by solute dissolving into a solvent. All ionic compounds dissolve in water to some degree. However, many compounds have such low solubility in water that we classify them as insoluble. MgCl2(aq) + Pb(OAc)2(aq) Mg(OAc)2(aq) + PbCl2(s) Solubility is an equilibrium process. We can apply the concepts of equilibrium to salts dissolving, and use the equilibrium constant to measure even slight solubilities in water. 7

E09 Exp 09 - Solubility Solubility Solvation Precipitating Insoluble Substances Using Q The reaction coefficient Comparing Q to Ksp Predicting Solubility Solubility Equilibrium Solubility Product, Ksp Relating Molar Solubility Factors Effecting Solubility Common Ion Effect Solution ph This Weeks Experiment The reactions Titration Next Meeting 8

Solubility Product The equilibrium constant for the dissociation of a solid salt into its aqueous ions is called the solubility product, Ksp Even ionic substances described as insoluble have some solubility. For an ionic solid MnXm, the dissociation reaction is MnXm(s) nm m+ (aq) + mx n (aq) The solubility product would be Ksp = [M m+ ] n [X n ] m For example, the dissociation reaction for PbCl2 is PbCl2(s) Pb 2+ (aq) + 2 Cl (aq) PbCl2 And its equilibrium constant is Ksp = [Pb 2+ ][Cl 1 ] 2 10

Solubility Product The equilibrium constant for the dissociation of a solid salt into its aqueous ions is called the solubility product, Ksp Even ionic substances described as insoluble have some solubility. For an ionic solid MnXm, the dissociation reaction is MnXm(s) nm m+ (aq) + mx n (aq) The solubility product would be Ksp = [M m+ ] n [X n ] m For example, the dissociation reaction for PbCl2 is PbCl2(s) Pb 2+ (aq) + 2 Cl (aq) And its equilibrium constant is Ksp = [Pb 2+ ][Cl 1 ] 2 11

Molar Solubility Solubility is the amount of solute that will dissolve in a given amount of solution at a particular temperature. The molar solubility is the number of moles of solute that will dissolve in a liter of solution before it becomes saturated. The molarity of the dissolved solute in a saturated solution. Molar solubility is related to Ksp Ksp = [M m+ ] n [X n ] m For the general reaction MnXm(s) n M m+ (aq) + m X n (aq) 13

Ksp and Relative Solubility Molar solubility is related to Ksp But you cannot always compare solubilities of compounds just by comparing their Ksp To compare Ksp the compounds must have the same dissociation stoichiometry. 14

The Effect of Common Ion on Solubility Addition of a soluble salt that contains one of the ions of the insoluble salt, decreases the solubility of the insoluble salt. For example, addition of NaCl to the solubility equilibrium of solid PbCl2 decreases the solubility of PbCl2. PbCl2(s) Pb 2+ (aq) + 2 Cl 1 (aq) Addition of Cl shifts the equilibrium to the left. 16

The Effect of Common Ion on Solubility Addition of a soluble salt that contains one of the ions of the insoluble salt, decreases the solubility of the insoluble salt. 17

The Effect of ph on Solubility For insoluble ionic hydroxides, the higher the ph, the lower the solubility of the ionic hydroxide. And the lower the ph, the higher the solubility Higher ph = increased [OH ] Mn(OH)(s) Mn + (aq) + noh (aq) For insoluble ionic compounds that contain anions of weak acids, the lower the ph, the higher the solubility. M2(CO3)(s) 2 M + (aq) + CO3 2 (aq) H3O + (aq) + CO3 2 (aq) HCO 3 (aq) + H2O(l) 18

Precipitation Precipitation will occur when the concentrations of the ions exceed the solubility of the ionic compound. If we compare the reaction quotient, Q, for the current solution concentrations to the value of Ksp, we can determine if precipitation will occur. Q = Ksp, the solution is saturated, no precipitation. Q < Ksp, the solution is unsaturated, no precipitation. Q > Ksp, the solution would be above saturation, the salt above saturation will precipitate. Some solutions with Q > Ksp will not precipitate unless disturbed; these are called supersaturated solutions. 21

Exp 09: Solubility Your job is to Determine the Ksp of Calcium Iodate Ca(IO3)2 by measuring the IO31- concentration in two different saturated solutions of Ca(IO3)2 #1 - saturated Ca(IO3)2 in pure water #2 - saturated Ca(IO3)2 in 0.0100 M KIO3 solution. You will be measuring that concentration by titrating the IO31with a known concentration of Na2S2O3 stock solution.

Finding Ksp The same Ksp can be used to determine the different molar solubility of Ca(IO3)2 in different environments. Today we will go the other way we will setup different solutions and go from the different molar solubilities we measure to find Ksp. #1 - saturated Ca(IO 3 ) 2 in pure water #2 - saturated Ca(IO 3 ) 2 in 0.0100 M KIO 3 solution. Because of the common ion effect, these solutions will have different molar solubilities different concentrations. But using an equilibrium calculation you will be able to determine the same K sp. Determining that K sp with two different systems produces a more effective demonstration of your results. 24

Finding Ksp Ca(IO3)2 (s) Ca 2+ (aq) + 2 IO3 1- (aq) Determine the solubility constant (Ksp) for the equilibration of solid Ca(IO3)2 into dissociated ions. Ksp = [Ca 2+ ] [IO3 1- ] 2 Ksp = (x) (2x) 2 = (x) (4x 2 ) = 4x 3 [IO3 1- ] = x 25

Finding IO3 1- We cannot do an acid base titration to determine the IO3 1- concentration (it s not a strong enough base). Instead, we will convert the IO3 1- to I2 which reacts reversibly with starch to produce a dark blue solution. IO3 (aq) + 5 I 1 (aq) + 6 H3O + (aq) 3 I2 (aq) + 9 H2O(l) RXN #1 I2 (aq) + starch dark blue complex RXN #3 (indicator) We will titrate the reaction with Na2S2O3 to reduce the I2 to I 1- until all the blue color is gone. I2 (aq) + 2 S2O3 2 (aq) 2 I 1 (aq) + S4O6 2 (aq) RXN #2 When the blue is gone we can know the amount of I2 that was produced, and therefore the IO3 1- that existed, by knowing how much Na2S2O3 we added. 26

Exp 09: Solubility Get Test Solutions Saturated solutions of calcium iodate in water, and calcium iodate with added iodate ion, are available in the hoods. CAUTION: Take care to not disturb the solid calcium iodate present in the bottles when obtained your portions of solutions to use. Obtain approximately 40 ml of the two different calcium iodate solutions in small beakers. One is labeled "Calcium Iodate, Aqueous. The other is labeled "Calcium Iodate, Added Calcium Ion". You will eventually use about 150 ml of this standardized thiosulfate solution. 27

Exp 09: Solubility Setup Burets Set up a 50 ml buret and stand, fill the buret with standardized sodium thiosulfate solution. Take a reading of initial volume to two decimal places. Be sure and record the exact concentration of the thiosulfate solution that is listed on the reagent bottle. This solution has been standardized by the Preparation Lab, and its concentration should be printed on the label. 28

Exp 09: Solubility Titration #1 Prepare sample: Pipet 10.0 ml of a calcium iodate saturated solution in pure water into a clean 125 ml Erlenmeyer flask. Using a graduated cylinder, add approximately 20 ml of distilled water to this saturated iodate solution, swirl to mix. Dissolve 0.5g of solid KI into the iodate/water solution, then add 10 ml of 1 M HCl. Swirl to mix the contents, obtaining dark red homogeneous solutions. First titrate the resulting brown solution with sodium thiosulfate until the brown color (I2) is mostly gone and the solution turned pale yellow (not golden). At this point, add 5 ml off a 0.1% starch solution. The titration solutions should become a dark blue-black color. Then titrate with standardized thiosulfate solution until a colorless endpoint. 29

Exp 09: Solubility Titration #2 Prepare sample: Pipet 10.0 ml of a calcium iodate saturated solution in 0.0100 M KIO3 into a clean 125 ml Erlenmeyer flask. Using a graduated cylinder, add approximately 20 ml of distilled water to this saturated iodate solution, swirl to mix. Dissolve 0.5g of solid KI into the iodate/water solution, then add 10 ml of 1 M HCl. Swirl to mix the contents, obtaining dark red homogeneous solutions. First titrate the resulting brown solution with sodium thiosulfate until the brown color (I2) is mostly gone and the solution turned pale yellow (not golden). At this point, add 5 ml off a 0.1% starch solution. The titration solutions should become a dark blue-black color. Then titrate with standardized thiosulfate solution until a colorless endpoint. 30

Stoichiometry IO3 (aq) + 5 I 1 (aq) + 6 H3O + (aq) 3 I2 (aq) + 9 H2O(l) RXN #1 This step, which occurs after adding both solid KI, and aqueous acid, to aliquots of saturated iodate solutions, has the net effect of converting iodate ions to aqueous iodine. Thiosulfate ion then reacts with aqueous iodine according to: I2 (aq) + 2 S2O3 2 (aq) 2 I 1 (aq) + S4O6 2 (aq) RXN #2 I2 (aq) + starch dark blue complex RXN #3 (indicator) The net titration reaction can be obtained by combining the two reactions above, then balancing for mass and charge: IO3 (aq) + 5 I 1 (aq) + 6 H3O + (aq) 3 I2 (aq) + 9 H2O (l) 3 I2 (aq) + 6 S2O3 2 (aq) 6 I 1 (aq) + 3 S4O6 2 (aq) IO3 (aq) + 6 S2O3 2 (aq) + 6 H3O + (aq) I 1 (aq) + 3 S4O6 2 (aq) + 9 H2O (l) 31

Next Week Before next Meeting: Bring to class: Notebook You will not be turning in notebooks, but this permanent record of your preparations, observations and notes will be essential to your success in this class. Textbook, calculator, pencils (yes, you can use pen) Safety Glasses (you cannot participate in the next class without them) Read and bring a copy of the next experiment Free Energy of Boraz Produce and bring to class: Your pre-lab for exp 10 Your procedure summary for exp 10 Review from your lecture text: Free Energy & Entropy We will start with a quiz about the experiment and reading. 33

Questions?