Introduction to the Physiology Laboratory
Physiology is commonly described as the study of biological function, while
anatomy is described as the study of biological structure. In fact, each discipline must
consider BOTH structure and function in order for either discipline to be meaningful. A
very practical difference between anatomy and physiology, however, is whether the
subject is dead or alive. Anatomical structures can be studied using preserved specimens.
Physiology, on the other hand, must be studied in a live subject (or in a close
approximation).
Studying a live subject poses certain problems for both the researcher and the
subject in a course like Human Physiology. The risk of transmission of infectious agents
is too high to monitor bodily fluids, typically saliva, urine, or blood, in a classroom
setting. We have also made a conscious decision to minimize the use of live animals,
although we have not altogether eliminated them. To circumvent these problems, we are
using a variety of alternatives including the non-invasive monitoring of bodily functions
(pulse, ECG, etc.), computer simulations, in vitro (in glass) demonstrations, and the
analysis of case studies.
Scattered throughout and at the end each exercise are study questions that are
meant to help guide you to the relevant points in the lab. Weekly quizzes will be based on
these questions and on the assigned background reading from the textbook, Human
Physiology: An Integrated Approach, 5th ed., by Dee Silverthorn. The pages for the
background reading can be found under the title of each exercise.
These lab exercises are meant to teach you something about experimental design
and methods, as well as help you understand the physiological concepts we discuss in
lecture. In choosing and writing up these exercises, we attempted to always ask “What is
the point of doing this?” and “What do we want the student to learn?”. We want these
exercises to be fun, but also relevant and useful. We hope that you will find your
experience with us this semester all of that.
A NOTE CONCERNING "AFFECT" AND "EFFECT": Students very commonly
get the terms "affect" and "effect" confused. "Affect" is a verb that means to have an
influence. "Effect" is a noun meaning the result or outcome. A change or stimulus will
affect the body's physiological state (example: Arteriosclerosis affects blood pressure).
The resulting effect of that change or stimulus is the response (example: High blood
pressure is one effect of arteriosclerosis). In other words, "af-" is the imput and "ef-" is
the output.
Exercise 1: Acids, Bases, and Buffers
Readings: Silverthorn 5th ed, pg. 36 – 39; 6th ed. pg. 47-49, 681-684
The acidic or basic properties of solutions are due to their relative concentrations
of hydrogen ion (H+) and hydroxyl (OH-) ion. Acidic solutions have high concentrations
of hydrogen ion and low concentrations of hydroxyl ion. Basic (also called alkaline)
solutions have low hydrogen ion concentrations and high hydroxyl ion concentrations.
Acids will bring about an increase in the hydrogen ion concentration of a solution, while
bases will bring about an increase in the hydroxyl ion concentration.
Why are acids and bases important to biology? Because life is a series of chemical
interactions that occur in water, and the acidity or alkalinity of a water solution greatly
influences the chemistry that is possible. Acids and bases most strongly affect the
enzymes that serve as catalysts for life's chemistry. If the proper enzymes are not
functioning because of acid or alkaline conditions, life processes will cease.
Enzymes are composed of proteins whose biological activity is dependent upon
an exact, three-dimensional shape. Substances that alter the shape or charge on proteins
are going to have significant effects on life processes. For example, the reason high
temperatures kill living things is because the high temperatures change protein structures.
Like high temperatures, the relative concentration of hydrogen ion and hydroxyl
ion exert effects on the shape of proteins, primarily by disrupting hydrogen bonding. This
results in a loss of enzyme activity. Most multicellular organisms have the ability to
regulate the acid-base conditions in their bodies and can live in environments that are
quite acidic or basic. Individual cells on the other hand, can generally tolerate only a very
narrow acid-base range.
The human body must maintain a very narrow range of pH between 7.38 and 7.42
(note that this pH is not precisely neutral). The human body uses three main mechanisms
for maintaining this narrow range of acid/base balance. The first and quickest acting
mechanism is the chemical buffering systems, the most important of which is the
carbonic acid/bicarbonate ion system. When chemical buffers are overwhelmed, usually
by increasing concentrations of H+, the second mechanism, the respiratory system, can
regulate acid/base balance by eliminating CO2 (see your textbook for an explanation of
this mechanism). These two mechanisms can only neutralize pH by converting H+ or
OH- into H2O. The only way to actually remove excess H+ is by the third mechanism, the
renal system. This mechanism actively secretes excess H+ into the urine during acidosis,
or HCO3– during alkalosis.
Today’s Objectives
1.
Review the basic principles of acid/base chemistry.
2.
Model a carbonic acid/bicarbonate buffering system using molecular models.
3.
Test the pH of a variety of common liquids.
4.
Observe pH fluctuations in an unbuffered solution.
5.
Create a chemical buffering system and examine its effect on pH fluctuations.
The Chemistry of Acids and Bases
Logically, you might only expect to find only water molecules (H2O) in a beaker
of pure water. However, if you perform an exact analysis, you find very small but equal
concentrations of hydrogen ion, (H+) and hydroxyl ion (OH-) in addition to water
molecules. The reason is a small fraction of the water molecules have dissociated to give
equal concentrations of hydrogen ion and hydroxyl ion:
H2O → H+ + OHThe solution is neither acid nor base, and is said to be neutral.
Further observations on pure water would demonstrate that the number of water
molecules dissociating is always the same, namely 1X10-7 M (0.0000001 moles per liter).
This doesn't sound like much but actually accounts for more than 1X1016 (more than a
thousand trillion) dissociated molecules in a liter of water. Since each dissociated water
molecule creates one H+ and one OH- both of these ions are at concentrations of 1X10-7M.
When an acid such as hydrochloric (HCl) is added to the water solution, the acid
dissociates:
HCl → H+ + ClThe acid adds to the [H+] ([H+] is read as "Hydrogen ion concentration") and it increases
[H+] above 1X10-7M. An interesting thing happens here. As the [H+] increases, there is a
proportional decrease in [OH-]. In fact, you would find that if the [H+] is multiplied by
the [OH-], the result will always be 1X10-14. This number is known as the water
dissociation constant.
Acids can be defined as H+ (or proton) donors, while bases can be defined as H+
acceptors. In other words, acids increase the [H+], while bases decrease it. Examine the
following list of common acids and bases. Notice how the acids all dissociate to produce
H+, thus increasing the [H+]. Some bases, like NaOH, can dissociate to produce OH- as
the H+ acceptor. Compounds like ammonia (NH3) and bicarbonate ion (HCO3–) are also
H+ acceptors and are considered bases.
Some Common Acids
Hydrochloric Acid
HCl → H+ + Cl-
Nitric Acid
HNO3 → H+ + NO3–
Acetic Acid
CH3COOH → H+ + CH3COO–
Some Common Bases
Sodium Hydroxide
NaOH → Na+ + OH–
Potassium Hydroxide
KOH → K+ + OH–
Ammonia
H+ + NH3 → NH4
It is awkward and time consuming to always write out the [H+] in Molar notation.
Instead, there is a better way of denoting the [H+]. The term for this notation is pH, and is
defined as the negative logarithm of the hydrogen ion concentration, or abbreviated:
pH = –log[H+]
Even though no units are used with pH notation, remember that pH notation
represents a concentration in moles per liter (M). Let's say you want to know the pH
of a neutral solution. You know that at neutrality:
[H+] = [OH–] and that [H+] X [OH–] = 1X10–14 M
thus:
[H+] = 1X10–7M
and:
pH = -log [H+] = -log(1X10–7M) = -(-7) = 7
Notice that the log is simply the exponent of 10 of the number in question. Therefore, the
pH of a solution at neutrality is 7.
Now, suppose we add some HCl to increase the [H+] to 1X10-4M. What is the pH at this
[H+]?
pH = -log [H+] = -log(1X10–4M) = -(-4) = 4
What happens if a base, such as NaOH is added to a neutral solution? The [H+] decreases
because [H+] X [OH–] must equal the dissociation constant. Let's say the [H+] is
decreased to 1X10–9M. What is the pH?
pH = -log [H+] = -log(1X10–9M) = -(-9) = 9
You should now see a relationship between [H+] and pH developing. At pH 7, a solution
is neutral; as the pH decreases, the [H+] is increasing and the solution is becoming more
acidic; as the pH increases above 7, the [H+] is decreasing and the solution is becoming
more basic.
Modeling the Carbonic Acid/Bicarbonate Buffer
In order to maintain a narrow pH range, the blood and other body fluids must be buffered.
Buffers are composed of conjugate acid-base pairs. Conjugate acid-base pairs are
compounds that differ by the presence of one proton, or H+. There are several buffering
systems in body fluids, but the most common is the carbonic acid/bicarbonate ion system:
H2O + CO2 → H2CO3 → H + HCO3–
+
Water and carbon dioxide combine to form a weak acid, carbonic acid (H2CO3) some of
which dissociates to form H+ and bicarbonate ions (HCO3-). Note that the reactions are
reversible and thus may go in either direction, depending on relative concentrations.
Materials
Bond models:
• Single bonds: fat gray pegs
• Multiple bonds: skinny gray flexible pegs. These are used to form two bonds
between the same atoms.
Ball models:
• Carbon (C): black balls with 4 holes (each hole represents one potential bond)
• Hydrogen (H): white balls with 1 hole
• Oxygen (O): red balls with 2 holes
• Nitrogen (N): blue balls with 4 holes (we won't be using these)
Procedure:
In the following diagrams, single lines connecting atoms indicate single bonds. Two
lines indicate a double bond. Double bonds are two bonds that form between the same
two atoms. The two atoms are sharing a total of 4 electrons.
Use the long flexible connectors to form double bonds. Each connector represents one
pair of electrons, so you need two connectors to make a double bond.
1.
Start out by making a water molecule and a carbon dioxide molecule:
Water: H2O
O
H
Carbon dioxide: CO2
2.
H
O=C=O
Combine these two molecules to form carbonic acid: H2CO3
O
||
C
/ \
HO OH
3.
Remove one of the hydrogen atoms to make bicarbonate ion: HCO3 –
4.
Combine your carbonic acid or bicarbonate ion and hydrogen ion with everyone
else's to make a buffering system. The instructor will demonstrate how this conjugate
acid and base pair makes a buffer and prevents fluctuations in H– concentration.
The pH of Some Common Solutions
You will now use a hand held Checkmite pH meter to measure the pH of some common
solutions. Make sure they are all at room temperature before making any measurements.
You will need to calibrate the meter before you begin your measurements.
Calibrating the pH meter
1.
The pH meters are stored with buffer solution in the cap to prevent the sensor from
drying out. Carefully remove the cap on the sensor tip of meter. Put the cap some
place safe so it doesn’t get lost. Rinse sensor tip with tap water and blot to remove
excess water. Press ON/OFF button to turn the meter on.
2.
Immerse about 1/2" to 1" of the sensor tip in pH 7.00 buffer.
3.
Press CAL button to enter Calibrate (CA) mode. 'CA' flashes on the display. Then, a
pH value close to 7 will flash repeatedly.
4.
After at least 30 seconds (about 30 flashes) press the HOLD/CON button to confirm
calibration. The display will show 'CO' and then switch to the buffer value reading.
5.
Rinse the electrode in tap water before pH testing.
Measuring the pH of a Sample
1.
Select about six different samples to test.
You do not have to test all available
samples. Space is available for samples
not listed in Table 1.
2.
Pour about 1" of the sample into the small
plastic cup provided. Immerse the sensor
tip in your sample solution. Let the reading
stabilize.
3.
Note the pH or press HOLD/CON button
to freeze the reading. Press HOLD/CON
button again to release the reading.
4.
Rinse and blot the sensor tip with tap
water. Always remember to rinse the tip
between each sample.
Table 1. The pH of some common
solutions
Solution
Deionized water
Apple juice (canned)
Vinegar
Coffee (instant, decaf)
Tea (black from bag)
Lemon juice (bottled)
Coke (Classic)
Milk (nonfat)
Mylanta
pH
5.
If you do not press a button on the pH meter for 8.5 minutes, the meter will
automatically shut off.
Unbuffered and Buffered Solution
In this exercise, you will compare the effect of adding a strong acid and base to first an
unbuffered solution, then a buffered solution.
Measuring pH Fluctuations in an Unbuffered Solution (Negative Control)
1.
Start with clean 1000 ml beaker filled with about 300 ml DI H2O.
2.
Place the beaker on the magnetic stirrer and drop a clean stirring bar into it.
3.
Turn on the stirrer and adjust the speed so that a small whirlpool is just barely
established.
4.
Add 2 ml of Universal Indicator, a pH indicator that will change color as the pH
changes.
5.
Record the pH and the color of the water in Table 2.
6.
Add 1 drop of the 1.0N HCl to the beaker. When the pH has stabilized, record the pH
of this solution.
7.
In the same way, add two more drops of the 1.0N HCl to the beaker, recording the
pH after each drop is added.
8.
This time, add one drop of 1.0N NaOH. Be sure to continue to record the pH after
each drop of acid or base is added and the pH has stabilized.
9.
In the same way, continue to drops of NaOH and record the pH until a total of 6
drops has been added, or until the pH of the solution has increased to around pH 9 to
10.
Table 2. pH Fluctuations in an unbuffered water
Event
pH
Color
Deionized water at start
1st drop 1.0N HCl added
2nd drop 1.0N HCl added
3rd drop 1.0N HCl added
1st drop 1.0N NaOH added
2nd drop 1.0N NaOH added
3rd drop 1.0N NaOH added
4th drop 1.0N NaOH added
5th drop 1.0N NaOH added
6th drop 1.0N NaOH added
7th drop 1.0N NaOH added
8th drop 1.0N NaOH added
Developing a Buffer System
Earlier, you made a carbonic acid/bicarbonate ion buffer using molecular models. Now
you will make one using the actual chemicals. We will use Alka Seltzer as a source of
CO2 and baking soda as a source of HCO3–.
H2O + CO2 → H2CO3 → H + HCO3–
+
Making the buffer
1.
Start with a fresh 1000 ml beaker filled with 300 ml DI H2O. Be sure to rinse the
beaker well between this experiment and the last. Place the beaker on the magnetic
stirrer and drop a clean stirring bar into it.
2.
Turn on the stirrer and adjust the speed so that a small whirlpool is just barely
established. Add 2 ml of Universal Indicator.
3.
Record the pH and the color of the water in Table 3.
4.
Place 50 ml of DI H2O in the stoppered flask. Drop 2 Alka Seltzer tablets into the
stoppered flask and snugly replace the stopper.
5.
Place the end of the glass tubing from the flask stopper into the swirling water. You
should see CO2 bubbles swirling out of its tip.
6.
When the Alka Seltzer tablets are just about dissolved, measure the pH of the
solution in the beaker.
7.
Add two more Alka Seltzer tablets to the flask and quickly replace the stopper.
8.
Keep bubbling CO2 into the beaker until the last two Alka Seltzer tablets are
dissolved. Measure the pH of the solution in the beaker.
9.
Add about 1 to 2 teaspoon of NaHCO3 (baking soda) to the beaker. Allow all of the
NaHCO3 to dissolve and record the pH.
Testing the buffer
1.
Add 1 drop of the 1.0N HCl to the beaker. Record the pH in Table 3.
2.
After about 30 seconds, add another drop of HCl. Record the pH.
3.
After another 30 seconds, add a third drop of HCl. Record the pH.
4.
You should observe only momentary fluctuations in the pH of the test solution.
5.
Allow the pH stabilize, but now add one drop of 1.0N NaOH and record the pH.
6.
In the same way, add a total of 6 drops of the NaOH solution, recording the pH after
each drop.
7.
When you are finished with the pH meter, turn it off and recap it. There should be a
damp piece of sponge in the cap to keep the sensor moist.
Table 3. pH Fluctuations in a buffered solution
Event
Deionized water at start
After first 2 Alka Seltzer tabs
After second 2 Alka Seltzer tabs
After adding baking soda
1st drop 1.0N HCl added
2nd drop 1.0N HCl added
3rd drop 1.0N HCl added
1st drop 1.0N NaOH added
2nd drop 1.0N NaOH added
3rd drop 1.0N NaOH added
4th drop 1.0N NaOH added
5th drop 1.0N NaOH added
6th drop 1.0N NaOH added
7th drop 1.0N NaOH added
8th drop 1.0N NaOH added
pH
Color
About the Buffer
The carbonic acid/bicarbonate buffering system made today is the same buffer found in
blood. The bubbling of CO2 into water mimics the CO2 produced by cellular respiration
that diffuses out of cells and into the plasma. CO2 combines with H2O to form first
carbonic acid, then bicarbonate ion:
H2O + CO2 → H2CO3 → H + HCO3
+
–
It is the dissociated H+ that caused the initial drop in pH. To make our buffer, we added
–
an excess of HCO3 by adding the sodium bicarbonate, which dissociated to form Na+
–
and HCO3 :
NaHCO3 → Na
+
+
HCO3
–
–
The Na+ ions have no effect on the pH and can be ignored. The HCO3 behaved as a
weak base and binds free H+, raising the pH and forming carbonic acid. This reaction is
written below from right to left to emphasize that adding baking soda increased the
amount of product, bicarbonate ion, to our initial reaction and shifted the equilibrium of
the reaction to the left.
H2CO3 <–– H + HCO3
It is the combination of both the carbonic acid and the excess bicarbonate ion that forms
the buffer. If acid is added, the bicarbonate ions will bind up the added H+ ions and
prevent them from decreasing the pH. If base is added, the carbonic acid will release H+
to combine with the added OH– (or other proton acceptor) to form water and prevent
them from increasing the pH. Any chemical buffering system requires both a conjugate
acid and a conjugate base to donate or accept compensatory H+.
+
−