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hydrolysis of salts, buffers, common ion effect

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  1. 1. 21.2 Salts in Solution
  2. 2. Salt Hydrolysis • A salt is an ionic compound that: –comes from the anion of an acid –comes from the cation of a base –is formed from a neutralization reaction –some neutral; others acidic or basic • “Salt hydrolysis” - a salt that reacts with water to produce acid or base
  3. 3. Salt Hydrolysis • Hydrolyzing salts usually come from: 1. strong acid + weak base, or 2. weak acid + strong base • Strong refers to the degree of ionization • How do you know if it’s strong? – Refer to the handout provided (downloadable from my web site)
  4. 4. Salt Hydrolysis • To see if the resulting salt is acidic or basic, check the “parent” acid and base that formed it. Practice on these: HCl + NaOH H2SO4 + NH4OH CH3COOH + KOH
  5. 5. Strong Acid + Strong base neutral solution Strong acid + weak base acidic solution (salt hydrolyzes acidic) Weak acid + Strong base basic solution (salt hydrolyzes basic) Salt Hydrolysis
  6. 6. Buffers • Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added –made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts
  7. 7. Buffers • ml HCl pH • added unbuffered buffered • 0 7.00 7.00 • 10 1.04 6.91 • 20 0.78 6.82 • 30 0.64 6.73 • 40 0.54 6.63 • 50 0.48 6.52 • 60 0.43 6.40 • 70 0.39 6.25 • 80 0.35 6.05 • 90 0.32 5.72
  8. 8. Buffers ml HCl added pH buffered unbuffered
  9. 9. Buffers • A buffer system is better able to resist changes in pH than pure water • Since it is a pair of chemicals: –one chemical neutralizes any acid added, while the other chemical would neutralize any additional base –AND, they produce each other in the process!!!
  10. 10. Buffers • Example: Ethanoic (acetic) acid and sodium ethanoate (also called sodium acetate) • Examples on page 628 of these • The buffer capacity is the amount of acid or base that can be added before a significant change in pH
  11. 11. Buffers • The buffers that are crucial to maintain the pH of human blood are: 1. carbonic acid (H2CO3) & hydrogen carbonate (HCO3 1- ) 2. dihydrogen phosphate (H2PO4 1- ) & monohydrogen phoshate (HPO4 2- ) – Table 21.2, page 629 has some important buffer systems
  12. 12. Buffers • How does the carbonic acid- hydrogen carbonate buffer system work? IF there is base added: H2CO3 + OH- -------H2O + HCO3 - If there is acid added: HCO3 - + H+ --------H2CO3 ( which is a weak acid)
  13. 13. Solubility Products, KSP • KSP expressions are used for ionic materials that are only slightly soluble in water Look at chart on pg 631 • Their only means of dissolving is by dissociation. • AgCl(s) Ag+ (aq) + Cl- (aq) • KSP = [ Ag+ ] [ Cl- ]
  14. 14. Solubility Products, KSP • At equilibrium, our system is a saturated solution of silver and chloride ions. • The only way to know that it is saturated it to observe some AgCl at the bottom of the solution. • As such, [AgCl] is a constant and KSP expressions are really Keq x constant conc of the solid so do not include the solid form in the equilibrium expression. • Chart on pg.632 shows Ksp for common slightly soluble salts
  15. 15. • What is the concentration of silver ions and chloride ions in a saturated silver chloride solution at 250 C? (Ksp= 1.8 x 10-10 ) Ksp=[Ag+ ][Cl- ]; [Ag+ ]=[Cl- ] 1.8 x 10-10 =x2 1.3 x 10-5 =x which is the conc. of silver and chloride ions Solubility Products, Ksp Problem
  16. 16. • Calcium flouride has a Ksp of 3.9 x 10-11 at 250 C. What is the flouride ion conc. At equilibrium? • CaF2------Ca2+ + 2F- Ksp=[Ca2+ ][F- ]2 Since there are twice as many F- as Ca2+ , Ksp=[x][2x]2 =2x3 3.9 x 10-11 =4x3 3.9 x 10-11 /4=x3 9.8 x 10-12 =x3 2.1 x 10-4 =x=F- Solubility Products, KSP, Problem
  17. 17. Common ion effect. • This is an example of Le Châtelier’s principle. Common ion effectCommon ion effect • The shift in equilibrium caused by the addition of an ion formed from the solute. Common ionCommon ion • An ion that is produced by more than one solute in an equilibrium system. • Adding the salt of a weak acid to a solution of weak acid is an example of this.
  18. 18. Common Ion Effect: Example • PbCrO4(s) Pb2+ (aq) +CrO4(aq) 2- If you add lead nitrate to this solution you increase the lead(II) ion and equil shifts to the left, and some lead(II) chromate will precipite out. The Ksp will stay the same. The lead ion is called the “common ion”, the lowering of the solubility of lead(II) chromate is called the common ion effect.
  19. 19. • If the ions concentrations are multiplied together, [ion][ion], and is greater than the Ksp, then a precipitate will form. Common Ion Effect
  20. 20. • The Ksp of silver iodide is 8.3 x 10-17 . What is the iodide ion conc. Of a 1.00L saturated solution of AgI to which 0.020mol of AgNO3 is added? Ksp=[Ag+ ][I- ]=x Because Ksp is so small, x is insignificant compared to 0.20mol of Ag+ so 8.3 x 10-17 =[0.20][x] 8.3x10-17 /0.20=4.2 x 10-15 M Common Ion Effect: PROBLEM
  21. 21. • Predict whether barium sulfate will precipitate when 0.50L of 0.002M Ba(NO3)2 mixes with 0.50L of 0.008M Na2SO4 to form a 1.0L solution. Ksp of BaSO4 is 1.1 x 10-10 Find [Ba2+ ] and [SO4 2- ] if this is > than Ksp,then precipitation will occur [Ba2+ ] = .50L x .002M=.001mol/1L [SO4 2- ]= .50L x .008M=.004mol/1L .001 x .004= 4 x 10-6 which is greater than the Ksp, so yes, BaSO4 will precipitate Common Ion Effect: PROBLEM