hydrolysis of salts, buffers, common ion effect

1. 21.2
Salts in Solution

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. 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. 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. 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. 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. 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. Buffers
ml HCl added
pH
buffered
unbuffered

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. 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. 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. 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)

14. 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-
]

15. 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

16. • 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

17. • 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

18. 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.

19. 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.

20. • If the ions concentrations are
multiplied together, [ion][ion],
and is greater than the Ksp,
then a precipitate will form.
Common Ion Effect

21. • 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

22. • 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