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BUFFER SOLUTION
A buffer is a solution that resists the change in pH when limited amounts of acid or base
are added to it.
One buffer system consists of a mixture of acetic acid and sodium acetate. The mixture
would contain the following particles: Na
, C
2
H
3
O
2
, HC
2
H
3
O
2
The reactions of added H
or of added OH
with particles in this buffer solution will be:
H
added
+ C
2
H
3
O
2
HC
2
H
3
O
2
OH
added
+ HC
2
H
3
O
2
H
2
O + C
2
H
3
O
2
The equations above show that a buffer system will maintain a relatively fixed pH even
when considerable acid or base is added. Until enough acid has been added to almost
exhaust the supply of C 2
H
3
O
2
ions, the pH of the solution will remain in the order of
magnitude of its original value. In the same way, the pH of the mixture will remain
almost constant until about all HC 2
H
3
O
2
molecules are exhausted by the added OH
ions.
Notice, as demonstrated by the equilibrium expression below, that a mixture of equal
volumes of equimolar HC 2
H
3
O
2
and C 2
H
3
O
2
(for example 1M solution of each) will
have a [H 3
O
] ion concentration nearly equal to its ionization constant.
HC
2
H
3
O
2
+ H
2
O H
3
O
+ C
2
H
3
O
2
K
a
= [H
3
O
] [C
2
H
3
O
2
]
[HC
2
H
3
O
2
]
TO MAKE A BUFFER SOLUTION:
One way to make a buffer solution is to choose a weak acid (with a K a
value of
approximately the H 3
O
concentration you want.) or a weak base (with a K b
value of
approximately the OH
concentration you want). Then calculate the conjugate-base/acid
or conjugate-acid/base ratio needed for your exact pH. (See the calculation immediately
below.) For example, if you want to prepare a buffer solution of pH=4.2, that is, a
[H
3
O
] = 6.3x 10
M, you need to choose a weak acid that has a K a
in the neighborhood
of 6.3x 10
H
3
O
2
is a good choice because it has a K a
of 1.8 x
You will see from the following calculation that by rearranging the equation for the
equilibrium expression, you may obtain the required ratio of volumes of equimolar
solutions of acid and conjugate base that you should mix.
HC
2
H
3
O
2
+ H
2
O H
3
O
+ C
2
H
3
O
2
K
a
= [H
3
O
] [C
2
H
3
O
2
]
[HC
2
H
3
O
2
]
K
a
= [C
2
H
3
O
2
]_ = moles of conjugate base/V final
[H
3
O
] [HC
2
H
3
O
2
] moles of acid/ V final
Since both compounds will now be in the same total volume of solution, V final
, this same
ratio will be the ratio of moles needed.
K
a
__ = moles of conjugate base = M conj. base
. V
conj. base
[H
3
O
] moles of acid M acid
. V
acid solution
If you use equimolar solutions of the two solutes (solutions of the same molarity before
mixing), then the calculated ratio will also be the ratio of volumes of the two solutions to
use.
K
a
__ = V
conj. base
[H
3
O
] V
acid solution
See how we specifically arrange to have:
molarity ratio= mole ratio=volume ratio
This is only true for equimolar starting solution.
Again fill the 10 ml graduate cylinder with the 1 M NaOH solution. Determine the
volume of the NaOH needed to obtain the same color change in the ammonium
acetate solution. Answer the questions on the report sheet.
SAVE THE REMAINNG 1 M NaOH FOR USE IN PART 4
SODIUM BICARBONATE AS A BUFFER:
- Put 5.0 ml portions of 0.10 M NaHCO 3
into each of two clean test tubes.
To one add 2 drops of methyl orange, followed by 1 M HCl until the indicator color
change is reached, about pH 3. Record the volume of acid used. Complete part 4 of the
report.
To the other, add 2 drops of Alizarin yellow R, followed by 1 M NaOH solution until the
indicator color change is reached, about pH 11. Record the volume of NaOH used.
Complete part 4 of the report.
PREPARATION OF A BUFFER MIXTURE:
- The instructor will assign you a pH for which you are to prepare about 20 to 30 ml of
buffer mixture. A selection of 0.10 M solutions for the possible buffer combinations
listed below are available on the Chem 111 shelves. Choose a suitable pair and
calculate what volume of each to use. Make the mixture, then with your instructor
check it with the pH pen. Complete part 5 of the report sheet.
Possible buffer pair Acid ionization constant
HSO
4
, SO
4
2 -
K
a
= 1.2 x 10
HCHO
2
, CHO
2
(formic acid, formate ion)
K
a
= 2.1 x 10
HC
2
H
3
O
2
, C
2
H
3
O
2
K
a
= 1.8 x 10
H
2
PO
4
, HPO
4
2 -
K
a
= 6.3 x 10
NH
4
, NH
3
K
a
= 5.6 x 10
HCO
3
, CO
3
2 -
K
a
= 4.7x 10
Report Sheet Name _________________
Buffer Solutions Last First
Instructor’s approval____________
pH of ammonium acetate solution:
- pH of 1.0 M NH 4
C
2
H
3
O
2 solution pH __________
Ammonium acetate as a buffer:
- Volume of 1 M HCl needed to decrease the pH of distilled water from pH 7 to 3 (in
drops or in ml)
______________drops
Volume of 1 M HCl needed to decrease the pH of 1.0 M NH 4
C
2
H
3
O
2
from pH 7 to pH 3
______________ml
Write the net-ionic equation for the reaction of HCl with NH 4
C
2
H
3
O
2
___________________________________________
- Volume of 1 M NaOH needed to increase the pH of distilled water from pH 7 to 11 (in
drops or in ml)
______________drops
Volume of 1 M NaOH needed to increase the pH of 1.0 M NH 4
C
2
H
3
O
2
from pH 7 to pH
______________ml
Write the net-ionic equation for the reaction of NaOH with NH 4
C
2
H
3
O
2
_____________________________________
NaHCO 3
as a buffer:
Volume of 1 M HCl needed to decrease the pH of 0.10 M NaHCO 3
to pH=3 ________ml
4. Write the net-ionic equation for the reaction of HCl with NaHCO
3
_____________________________________
Volume of 1 M NaOH needed to increase the pH of 0.10 M NaHCO 3
to pH=11______ml
Write the net-ionic equation for the reaction of NaOH with NaHCO 3
___________________________________________
Exercises
- Write net-ionic equations for the reaction of the buffer mixture, Na 2
HPO
4
with
NaH 2
PO
4
a. with added H
ions : _____________________________
b. with added OH
ions: ____________________________
- Write the formulas of the particles present in each of the following solutions.
Then classify these particles as Bronsted-Lowry acids or bases. If more than one
Bronsted-Lowry acid is present, list the strongest one first. Decide on which of
the solutions below would show a buffer action. (A buffer must be able to react
with both H
and OH
ions.
Solute Particles
present
Bronsted base
(capable in
reacting with
H
ions)
Bronsted acid
(capable in
reacting with
OH
ions)
Buffer (yes, or
no?)
NH
4
HSO
4
NH
3
NH
4
Cl
HCl,
NaCl
NH
4
Cl,
NaCl
H
2
CO
3
KHCO
3
HC
2
H
3
O
2
HCl
H
2
C
2
O
4
KHC
2
O
4
KHS,
Na 2
S
- Calculate the pH of an equimolar mixture of H 2
S and NaHS. Show the setup.
Ka1 for H 2
S is 1.1 x
Setup:
Answer__________
- a. Define : Buffer
b. How will the pH of a buffer solution change if we add water? (increases, decreases,
or remain the same)
c. Consider the table given below and write a balanced chemical equation for any
reaction taking place between solute particles. Then write the formulas of the
major particles present ( just as you would for a net-ionic equation) in each of the
following solutions below. Decide on which of the solutions below would show a
buffer action.
Particles present
Is it a buffer?
(Yes or No)
Na 2
CO
NaHSO 4
NaF
Equal volumes of 0.10 M HCN and 0.05 M
NaOH
Equation:
Equal volumes of 0.05 M H 2
S and 0.10 M
NaOH
Equation:
NaHC 2
O
