CHEM102 Lecture Notes - Lecture 17: Solubility Equilibrium, Ionic Compound

17 views2 pages

skycaribou653

17 Apr 2018

School

University of Alberta

Department

Chemistry

Course

CHEM102

Professor

Camille Bowen-Forbes

For unlimited access to Class Notes, a Class+ subscription is required.

kathanjani01 and 38922 others unlocked

CHEM102 Full Course Notes

Verified Note

32 documents

Document Summary

Precipitation equilibria: any insoluble ionic compound is actually slightly soluble in aqueous solution, equilibrium exists between solid solute and aqueous ions. [pbf2] is constant as it is a solid, so qc[pbf2] can be considered a constant value. When solution is saturated the system is in equilibrium and qsp= ksp. For metal sulfides, ksp = [mn2+][hs-][oh-] because s2- is a strong base and reacts with water: qsp < ksp solubility. Qsp > or = ksp precipitation: soluble compounds. Cl, br, i salt except agx, hg2x2, pbx2. So4 salts except baso4, srso4 and pbso4. Oh salts except group 1 and nh4. Co3, po4 and s salts except group 1 and nh4: all above rules in pure water. Solubility changes with ph: effect of common ion on solubility: Common ion causes equilibrium to shift to the left. Solubility of pbi2 in water is more than its solubility in nai. 1700 times smaller than in pure water: effect of ph on solubility:

Related Questions

Solution Investigation Question: What are some of the key characteristics of unknown sample of an aqueous solution containing an ionic compound?

1. Given that the cation is from Group 1A, what is the identity of the group 1A metal ion in your solution?

2. Given that the ionic compound contains either chloride or hydroxide, how can you determine which? (Two pieces of evidence are required and one must be related to solubility rules)

3. What is the concentration of your solution? (Your unknown sample is known to be about 1 M, but you need to determine the concentration to 3 significant figures using titration. You should plan ahead so that you know approximately what volume of your sample you will need in order to use approximately 25 mL of 0.100 M sulfuric acid for the titration).

Show all your steps for calculation questions in order to get full points.

1. (4 pts) The solubility product (K_sp) of the insoluble salt AgBr is 5.0 x 10^-13. Calculate the equilibrium constant for the following reaction: Ag⁺(aq) + Br^-(aq) AgBr(s).

2. (6 pts) A solution is prepared by adding 50.0 mL of 0.050 M HCl to 150.0 mL of 0.10 M HNO₃.

a. Calculate the pH of the solution.

b. Calculate the concentrations of all species in this solution.

3. (4 pts) Calculate the volume of concentrated (12 M) HCl needed to prepare 1-L of a pH = 1.50 solution?

4. (5 pts) Arrange the following 0.10 M solutions in order of most acidic to most basic.

KOH, KCl, KCN, NH₄Cl, HCl

5. (6 pts) Calculate the pH of an aqueous solution containing 1.0 x 10^-2 M HCl and 1.0 x 10^-2 M HCN.

6. (5pts) Carbonate buffers are important in regulating the pH of blood at 7.40. What is the concentration ratio of CO₂ (usually written H₂CO₃) to HCO₃^- in blood at pH=7.40?

H₂CO₃(aq) HCO₃^-(aq) + H⁺(aq) K_a = 4.3 x 10^-7

7. (10 pts) HCN gas, a powerful respiratory inhibitor, is highly toxic. It is a very weak acid (K_a = 6.2 x 10^-10) when dissolved in water. If a 50.00-mL sample of 0.100 M HCN is titrated with 0.100 M NaOH, calculate the pH of the solution:

a. before the titration starts.

b. after 8.00 mL of 0.100 M NaOH has been added.

c. at the half-way point of the titration.

d. at the equivalence point of the titration.

Show all your steps for calculation questions in order to get full points.

1. (4 pts) The solubility product (K_sp) of the insoluble salt AgBr is 5.0 x 10^-13. Calculate the equilibrium constant for the following reaction: Ag⁺(aq) + Br^-(aq) AgBr(s).

2. (6 pts) A solution is prepared by adding 50.0 mL of 0.050 M HCl to 150.0 mL of 0.10 M HNO₃.

a. Calculate the pH of the solution.

b. Calculate the concentrations of all species in this solution.

3. (4 pts) Calculate the volume of concentrated (12 M) HCl needed to prepare 1-L of a pH = 1.50 solution?

4. (5 pts) Arrange the following 0.10 M solutions in order of most acidic to most basic.

KOH, KCl, KCN, NH₄Cl, HCl

5. (6 pts) Calculate the pH of an aqueous solution containing 1.0 x 10^-2 M HCl and 1.0 x 10^-2 M HCN.

6. (5pts) Carbonate buffers are important in regulating the pH of blood at 7.40. What is the concentration ratio of CO₂ (usually written H₂CO₃) to HCO₃^- in blood at pH=7.40?

H₂CO₃(aq) HCO₃^-(aq) + H⁺(aq) K_a = 4.3 x 10^-7

a. before the titration starts.

b. after 8.00 mL of 0.100 M NaOH has been added.

c. at the half-way point of the titration.

d. at the equivalence point of the titration.

e. after 52.00 mL of 0.100 M NaOH has been added.

8. (5 pts) Calculate the K_sp value. For bismuth sulfide (Bi₂S₃), which has a solubility of 1.0 x 10^-15 mol/L at 25â°C.

9. (5 pts) Arrange the following group of salts: in order of most soluble to most insoluble.

AgI (K_sp= 1.5 x 10^-16), CuI (K_sp= 5.0 x 10^-16) and CaSO₄ (K_sp= 6.1x 10^-5)

e. after 52.00 mL of 0.100 M NaOH has been added.

8. (5 pts) Calculate the K_sp value. For bismuth sulfide (Bi₂S₃), which has a solubility of 1.0 x 10^-15 mol/L at 25â°C.

9. (5 pts) Arrange the following group of salts: in order of most soluble to most insoluble.

AgI (K_sp= 1.5 x 10^-16), CuI (K_sp= 5.0 x 10^-16) and CaSO₄ (K_sp= 6.1x 10^-5)

tanbutterfly616

THIS IS A FOUR PART QUESTION. HELP NEEDED. THANK YOU VERY MUCH.

a) Using the solubility of Ca(IO3) 2 that you determined in pure water, calculate the value of K sp using only concentrations, that is assuming that all activity coefficients are 1.0. Use 3 sig figs. Solubility of Ca(IO3)2 in Water: 0.00828 mols/L

b) Using the solubility of Ca(IO3)2 that you determined in 0.01M KIO3 as a solvent (the answer from part a), calculate the value of Ksp using only concentrations, that is assuming that all activity coefficients are 1.0. Solubility of Ca(IO3)2 = 0.00327

c) The next few questions will incorporate activity effects into the calculation of Ksp for Ca(IO3)2 using your data from the saturated solution of Ca(IO3)2 in pure water.To do so, we must first know the ionic strength of the solutions. Use your value of the solubility of Ca(IO3)2 in pure water to calculate the ionic strength of that solution.

d) Using the value of the ionic strength of a saturated solution of Ca(IO3)2 in pure water, use TABLE I under the solubility product experiment from the laboratory manual, and enter the activity coefficient for Ca2+ ions listed there at the ionic strength closest to that of the saturated solution of Ca(IO3)2 in pure water. If your ionic strength happens to lie exactly between two values in TABLE I, then use the activity coefficient for the higher ionic strength.

If you need any other information, please let me know.

CHEM102 Lecture Notes - Lecture 17: Solubility Equilibrium, Ionic Compound

CHEM102 Full Course Notes

Document Summary

Get access

Related textbook solutions

Chemistry: Structure and Properties

Basic Chemistry

Principles of Chemistry Molecular Approach

Chemistry: Structure and Properties

Principles of Chemistry Molecular Approach

Chemistry: A Molecular Approach

Chemistry: A Molecular Approach

Principles of Chemistry: A Molecular Approach

Chemistry: The Central Science

Related Documents

CHEM 230 Lecture Notes - Lecture 6: Supersaturation, Gibbs Free Energy, Mole Fraction

01:160:162 Midterm: Chem 2 Exam 3

CHEM 101 Study Guide - Midterm Guide: Barium Hydroxide, Alkali Metal, Nitric Acid

Related Questions