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pH, Buffers and Biological Molecules
The physical and chemical environment of biological molecules may affect the
properties of these molecules. The term "homeostasis" (homeo = the same, stasis = state) refers
to maintaining a constant internal environment within cells. This is important if molecules are
to exhibit the desired properties. Factors that may be kept constant include temperature, salt
concentration, osmolarity, and pH. Human and most other cells maintain an internal pH of
approximately 7.2. Any pH greater than 7 is considered basic, 7 is neutral and any pH less than
7 is considered acidic.
pH is actually a measure of the hydrogen ion concentration [H+] in an aqueous solution
(don't panic here, but in chemistry you will learn a more rigorous definition, but for most of
biology this definition will do the job). The "p" in pH stands for the negative logarithm to the
base ten (-Log10) and the "H" stands for the hydrogen ion concentration [H+]. In pure water, the
pH is exactly 7 (neutral). This means that the negative log of the hydrogen ion concentration in
water is 7. This number is determined based on the dissociation constant for water. Water may
ionize to form H+ and OH- ions:
+
H 2 O_ H + OH
The equilibrium constant (K'eq) for water can be found by dividing the product of the
molar concentrations of the ions produced when water ionizes by the total concentration of
water:
As the concentration of pure water is 55.5 M:
K'eq has been very carefully measured for pure water and found to be 1.8 x 10-16 M. Therefore,
(55.5M)(1. 8x 10-16 ) = [ H + ][ OH - ]
solving for [H+][OH-]:
1.0x 10-14 = [ H + ][ OH - ]
or:
Therefore, the product of [H+] and [OH-] is always 1.0 x 10-14. At neutral pH, [H+] = [OH-] so
the concentration of both must be 1.0 x 10-7 M (remember that we are dealing with exponents
here). The negative log (p) of 1.0 x 10-7 is 7. This is how the value of 7 is arrived at for neutral
pH.
As the value of the pH goes up, the [H+] goes down. This may seem counter intuitive,
but if you think it through, you will see why. The larger the absolute value of a negative
exponent, the smaller the number. For example, 2-1 = 0.5, but 2-2 = 0.25. The absolute value of
-2
-2 is grater than the absolute value of -1 and 0.25 is smaller than 0.5. A solution with a pH of 7
has a [H+] of 1.0 x 10-7 M. A solution with a larger pH, lets say 10, has a [H+] of 1.0 x 10-10 M
or 0.001 the concentration of a neutral solution! For every change of 1 pH unit, there is a 10 x
change in [H+]. Table 1 gives the relative concentrations of H+ and OH- with different pHs.
TABLE 1
The pH Scale With Examples
[H+]
M
1.0
0.1
0.01
0.001
0.0001 4
0.00001
10-6
10-7
10-8
10-9
10-10
10-11
10-12
10-13
10-14
pH
[OH-]
M
pOH
0
1
2
3
10-14
10-13
10-12
10-11
14
13
12
11
10-10
5
6
7
8
9
10
11
12
13
14
10
10-9
10-8
10-7
10-6
0.00001
0.0001 4
0.001
0.01
0.1
1.0
9
8
7
6
5
3
2
1
0
Everyday
Examples
1 M HCl
Gastric juice
Lemon juice
Cola, vinegar
Tomato juice
Black coffee
Blood, tears, saliva
Seawater, egg white
Baking soda (NaHCO3)
Household ammonia
Household bleach
1 M NaOH
Because maintaining homeostasis within an organism is so important, buffer systems
exist to prevent rapid massive internal pH changes. Buffers are solutions containing a weak acid
and its salt. If the pH goes up, the salt tends to dissociate and hydrogen ions associate with the
acid thus removing H+ from the solution so that pH remains constant. Addition of a base tends
to remove H+ from the solution, but when this happens, the weak acid tends to loose H+ to the
solution so that the pH is maintained close to the original pH of the solution. An example of a
buffer in human blood is the bicarbonate buffer system.
Methods of Determining pH
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Two methods are commonly used to determine the pH of solutions; 1) colorimetric
methods and 2) pH meters. Each method offers advantages and disadvantages and the method
chosen to determine the pH of a particular solution must be carefully evaluated before use to
ensure usable data are produced.
Colorimetric methods
Colorimetric methods rely on changes in color of a chemical indicator in response to
changes in pH of a solution. You may have used litmus paper in the past to determine whether a
solution was acidic or basic. Litmus is a chemical that turns red in acidic and blue in basic
solutions. Thymol blue is another commonly used chemical that changes color over a wide
range of pHs. It goes through a whole rainbow of colors depending on the pH of the solution it
is in. By examining the color of thymol blue in a solution of known pH, it may be possible to
determine the pH of a solution whose pH is unknown. Special paper called pH paper is also
available. This paper has been treated with chemicals that may change color over a very wide or
a small pH range and is valuable in making more accurate pH determinations than are possible
with chemicals that change color over a wide pH range.
Advantages of colorimetric methods include:
1)
2)
3)
4)
Ease of use.
Calibration may not be needed.
Materials are easily portable.
No power supply or expensive equipment is needed.
Disadvantages include:
1)
2)
3)
Colors produced by an unknown may not exactly match those on a color chart or
standard solution.
Opaque solutions, or solutions that already have a color may not produce
readable results.
Chemicals other than hydrogen ions may react with the indicator causing
problems in interpretation.
Exercise 1
SAFETY NOTE - SOME OF THE SOLUTIONS USED IN THIS LAB CAN CAUSE
SEVERE BURNS IF THEY GET ON SKIN OR CLOTHING, OR IN YOUR EYES. USE
EXTREME CARE WHEN HANDLING ALL CHEMICALS. IF SOLUTIONS SPLASH
ON ANY PART OF YOUR BODY FLUSH WITH WATER IMMEDIATELY THEN
INFORM THE LAB ASSISTANT. IMMEDIATELY REMOVE ALL CLOTHING
CONTAMINATED WITH CHEMICALS.
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1.
Obtain a spot plate and ensure that it is clean of any contamination.
2.
Number 12 wells with a wax pencil to ensure there is no confusion about what is in each
well. Into well 1 place 5 drops of pH 1 buffer. Into well 2 place 5 drops of pH 2 buffer
and so on until each well has 5 drops of a buffer whose pH corresponds to the well
number. It is essential that only buffer solutions are used in this part of the lab.
3.
Place one drop of thymol blue solution in each well containing buffer solution and
gently swirl the plate to mix thymol blue and buffer.
4.
Carefully record the colors in each well. Do not throw out this set of standards, you will
need it for later use.
5.
Determine the pH of the unknown solutions by using the above procedure and
comparing the color produced with the standards you just made.
pH Meters
pH meters are devices that measure changes in electrical potential (Voltage) between the
inside and the outside of a special electrode. At the tip of the electrode is special pH sensitive
glass. Inside the electrode is a special buffer solution. Differences in potential develop across
the tip of the electrode due to differences in hydrogen ion activity between the solution within
the electrode and the solution being tested.
Advantages of pH meters include:
1)
2)
3)
Their readings may be more precise than colorimetric methods allow.
Readings are not influenced by the natural color or opacity of a solution.
Results may be easier to read as they are displayed numerically and do not
require interpretation from color standards.
Disadvantages include:
1)
2)
3)
4)
Exercise 2
1.
Care must be taken to avoid damaging the electrode.
The equipment used may be expensive and complex.
Careful calibration of the instrument is essential.
Some chemicals other than H+ may cause spurious readings.
Obtain a pH meter and calibrate it by doing the following:
a.
b.
c.
Switch on the pH meter.
Take the temperature of the buffer solution you intend to calibrate with.
Adjust the knob marked "Temperature" on the pH meter to correspond with the
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d.
e.
f.
g.
h.
temperature of the buffer solution.
Carefully remove the electrode from its storage buffer. NEVER LET THE
ELECTRODE DRY OUT. DO NOT LEAVE THE ELECTRODE IN
HIGH SALT OR LOW pH SOLUTIONS FOR LONG PERIODS OF
TIME.
Rinse the electrode with distilled water.
Place the tip of the electrode in the calibration solution of know pH.
Adjust the calibration knob until the display indicates the known pH of the
calibration solution.
Rinse the electrode again with distilled water and either place it back in the
storage buffer or into a solution that you want to measure the pH of.
2.
Measure the pH of all the buffer solutions and the 1 M NaOH and 1 M HCl using the pH
meter. When doing this it is essential that the temperature of all solutions be the same as
the temperature of the calibration solution. In addition IT IS ESSENTIAL THAT
THE ELECTRODE BE RINSED WITH DISTILLED WATER BETWEEN
EACH MEASUREMENT.
3.
Measure the pH of the unknown solutions. Are the results obtained using this method
the same as those obtained using thymol blue? Why?
Exercise 3
1.
Make 10 x serial dilutions of the following solutions: 1 M NaOH, 1 M HCl, the highest
and the lowest pH buffers. This is done by taking the original solution and adding one
part of the original to nine parts distilled water, mixing then taking one part of the new
solution and adding nine parts distilled water. Do this six times so that you have 1.0,
0.1, 0.01, 0.001, 0.0001, 0.00001, and 0.000001 dilutions.
2.
Measure the pH of each dilution. Did you get results you expected? How do you
explain discrepancies between what you expected to happen and your actual data? How
do you explain differences in results obtained using buffered and unbuffered solutions?
Exercise 4
BE CAREFUL NOT TO BURN YOURSELF OR YOUR COLLEAGUES. IF YOU ARE
BURNED, IMMEDIATELY PLACE THE BURNT AREA UNDER COLD RUNNING
WATER AND INFORM YOUR LAB INSTRUCTOR.
1.
Make a boiling water bath by bringing 150 ml of water to a boil in a 500 ml beaker.
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2.
Place purple kale or cabbage into boiling water for 5 min.
3.
Drain off and save the now purple water. Discard the boiled kale.
4.
Allow the water to cool to room temperature.
5.
Repeat Exercise 1 using the purple water instead of thymol blue. Did the biological
molecule you extracted using boiling water change its properties when the pH was
changed? Does this indicate that other biological molecules may also change their
properties when the pH changes?
Exercise 5 For A students only
Design a simple experiment of your own. For example, measure the pH of your body
fluids; saliva, urine, sweat, blood etc.. Examine the effect of changing pH on solutions like
milk, egg white, orange juice or grape juice. Check other biological pigments ie chlorophyll or
hemoglobin to see if they may also work as indicators. Be sure to indicate what method you
used to measure pH if you chose to do that and why you chose that method.
Materials:
Equipment
Droppers
pH meters
Spot plates
Beakers, 500 ml
Hot plates
10 ml Test tubes
Test tube racks
Chemicals
Buffer solutions ranging from pH 1 to 12
Distilled water
HCl 1 M solution
NaOH 1 M solution
pH paper
Thymol Blue solution
3 Unknown pH solutions
Consumables Purple kale
Various solutions of biological origin ie. milk, grape juice, orange juice etc.
For more information on pH - Read chapter 3 in Campbell (1993). Most chemistry texts also
have extensive discussions of pH. Biochemistry texts, like Principles of Biochemistry by
Lehninger, discuss the biological significance of pH.