Ashford University Week 1 to 4 I Basic Chemistry Lab

Microscopy: Theory and Practice
After completing this exercise, you should be able to
1. Identify the parts of the compound microscope and state the function of each
2. Describe and demonstrate the correct way to :
a. Carry a microscope
b. Clean the lenses of a microscope, and prepared slides
c. Find and focus on the image of a specimen using low and high power
objectives on the compound microscope
d. Estimate the size of a microscopic specimen
3. Briefly explain the concepts of inversion, and depth of field.
4. State the contributions of the following people to the development of cell
theory: Galileo, Kepler, Leeuwenhoek, Hooke, Schleiden and Schwann.
During most of man’s history, it was not possible to “see” any of the small units
that make up living things. The first compound microscope (a magnifying tube
with two lenses) was invented by a father and son Dutch team, Hans and Zacharias
Janssen between about 1590 and 1600. By 1609 the Italian Astronomer and
physicist, Galileo (1564-1642) had heard of the invention and had constructed his
own magnifying tube having a power of 32X. Galileo sent one of his magnifying
tubes to the German astronomer, Johann Kepler (1571-1630) who used it as a
telescope but who also was stimulated to describe a compound microscope (one
with more than one lens). Microscopy became all the rage in scientific circles and
Marcello Malphigi, an Italian physiologist, used it to study frog lungs (opening the
way for an understanding of respiration) and bat wings, resulting in his describing
tiny blood vessels later called “capillaries”, from the root meaning “hair like”.
Anton van Leeuwenhoek (1632-1723), a Dutch microscopist, was a contemporary
of Malphigi who is perhaps the most interesting observer of microbiological life in
the early days of microscopy. He discovered protozoa, bacteria and was the first to
describe sperm. His findings greatly impressed and influenced the Royal Society of
England. Robert Hooke (1635-1703), secretary of the Royal Society and himself an
accomplished microscopist, built microscopes according to Leeuwenhoek’s design
and was impressed with the Dutchman’s discoveries. Hooke is the individual who
coined the term “cell” to describe the small chambers he observed in the cork of
tree bark utilizing his own compound microscope.
Many other investigators working with microscopes studied all kinds of living
things. They too identified the cells that Hooke had seen. In 1838, Matthias
Schleiden (1804-1881), a German botanist developed a theory that all plants were
composed of living units called cells. A year later, a German physiologist, Theodor
Schwann (1810-1882), developed what he called the “cell theory” stating that
animals as well as plants had cells as the basis for their living structure.
Microscopes are precision instruments and essential tools in the study of cells.
Please handle microscopes carefully: carry them with two hands, one hand under
the base and the other grasping the arm. Microscopes have one or more lens
systems, light sources, and mechanisms for adjusting the distance between the lens
systems and an object. This exercise provides an opportunity for you to develop
expertise in your knowledge and use of microscopes.
Section 1: The Compound Microscope
Activity A: Identifying the Parts of the Compound Microscope
Activity B: Viewing a slide: Letter e
Activity C: The Size of the Field: Estimating the Size of
Microscopic Organisms
Activity D: Depth of Field
Section 1. The Compound Light Microscope
The name compound light microscope indicates that it uses two sets of lenses and
light to view an object. The two sets of lenses are the ocular lenses located near the
eyes and the objective lenses located near the object. Illumination is from below,
and the light passes through clear portions but does not pass through opaque
portions of a specimen. This microscope is used to examine small or thinly sliced
sections of objects.
Activity A: Identifying the Parts of the Compound Microscope
Obtain a compound light microscope from the storage area, and place it securely
on the table. Plug the microscope’s wire into the nearest electrical outlet. Using
the following descriptions and instructions, identify the following parts on your
microscope, and label them in Figure 4.1.
Figure 4.1 Compound light microscope.
Compound light microscope with binocular head and mechanical stage. Label the
drawing of this microscope with the help of the text material.
1. Ocular lenses (eyepieces): You should adjust the eyepieces by carefully
moving them apart or together, to match the separation of your eyes. Each
eyepiece is marked with a number followed by an X.. That number is the
magnifying power of the eyepiece lenses. What is the magnifying power of the
ocular lenses on your microscope?___________
2. Body tube: Holds the revolving nosepiece at one end and eyepiece at the other
end; conducts light rays.
3. Arm: Supports upper parts and provides carrying handle.
4. Nosepiece or turret: Revolving device that holds objectives.
5. Objective lenses (objectives):
a. Scanning objective: This is the shortest of the objective lenses and
is used to scan the whole slide. The magnification is stamped on the
housing of the lens. What is the magnifying power of the scanning
objective lens on your microscope?
b. Low-power objective: This lens is longer than the scanning objective
lens and is used to view objects in greater detail. What is the magnifying
power of the low-power objective lens on your microscope?
c. High-power objective: This lens is used to view an object in even
greater detail. What is the magnifying power of the high-power objective
lens on your microscope?
d. Oil immersion objective: Holds a 100 X lens and is used in conjunction with immersion oil to view objects with the greatest magnification.
Your instructor will discuss its use when the lens is needed.
As the magnifying power increases, what happens to the length of the objective?
To find the total magnifying power from the specimen to your eyes, you must
multiply the magnification of the ocular lens by the magnification of the objective
lens. Thus when using the 4 X lens, the total magnifying power is 4X x 10X (for
typical oculars) = 40X. What is the total magnifying power if you are using the
10X lens?_____________
6. Coarse-adjustment knob: Knob used to bring an object into approximate
focus. With no slide in place, turn the knob. Notice that it lifts and lowers the
objective lens A LOT. Use this knob only with the low-power objective.
7. Fine-adjustment knob: Knob used to bring object into final focus.
8. Stage: Holds and supports microscope slides.
9. Mechanical stage: A movable stage that aids in the accurate positioning of the
10.Mechanical stage control knobs (not shown): Two knobs usually located
below the stage. One knob controls forward/reverse movement, and the other
controls right/left movement.Turn the knobs to decide which one moves the slide
in which direction.
11.Condenser: Lens system found below the stage that “gathers” light from the
light source and directs that light through the hole in the stage. There is a knob
under one side of the stage that controls the height of the condenser: the condenser
works best if it is in position at the top (nearest the stage) at all times.
12.Light source: An attached lamp that directs a beam of light up through the
object. Turn it on using the switch on the front of the bse of the microscope.
13.Variable light control: wheel-type control at one side of the base. Rotating
the wheel from lower to higher numbers increases the brightness of the light.
14.Diaphragm or iris diaphragm control lever (not shown: near 11): Moving
the lever regulates the amount of illumination passing upwards through the
With your eye to the side, look at the hole in the stage. Rotate the variable light
control wheel through the numbers. What do you notice about the light coming
through the hole in the stage with regard to the numbers?____________________
Set the wheel at number 5 or 6. Again with your eye at the side, move the iris
diaphragm lever back and forth. What do you notice about the light coming
through the hole in the stage?
Notice that you have two different ways to control the amount of light that travels
through the specimen.
Activity B. Viewing a slide.
1. Turn the nosepiece so that the lowest power objective on your microscope is in
straight alignment over the stage.
2. Always begin focusing with the lowest power objective on your microscope
(4X [scanning]).
3. With the coarse- adjustment knob, lower the stage until it stops.
4. Go to the front table and take a slide labeled “letter e” from the tray.
5. Hold the slide up to the overhead lights and notice the position of the letter e.
Draw that position in the left circle below.
6. Pull the vertical cylinder to open the spring arm on the stage, and insert the
slide. Use the two control knobs located below the stage to center the e over the
hole in the stage.
7. Again, be sure that the lowest-power objective is in place. Then, as you look
through the eyepieces, decrease the distance between the stage and the tip of the
objective lens until the object-in this case, the letter e-comes into view, or focus.
8. Once the object is seen, you may need to adjust the amount of light. To increase
or decrease the contrast, rotate the iris diaphragm lever slightly, or use the
numbered wheel on the base and go to a higher or lower number.
9. Now carefully use the fine-adjustment knob to sharpen the focus if necessary.
10.Practice having both eyes open when looking through the eyepiece, as this
greatly reduces eyestrain.
11.Move the slide around until the entire letter e is within the field of view. In the
circle at the right, make a rough drawing of what you see.
Orientation with naked eye
How it looks through the microscope
Observation: Inversion
12.What differences do you notice between the e on the left (naked eye) and that
on the right (microscope)?
13.Move the slide to the right. Which way does the image appear to move?
Summarize what the microscope does to specimens.
Higher Power Objectives
Compound light microscopes are parfocal; that is, once the object is in focus at a
low power, it should also be (almost) in focus with the next highest power.
1. Make sure that the letter e is centered in the field of the lowest objective.
2. Move to the next higher objective (low power [10X]) by turning the nosepiece
until you hear it click into place. Do not change the focus (move the stage);
parfocal microscope objectives will not hit normal slides when changing the focus
if the lowest objective is initially in focus.
3. Now pick a portion of the e you would like to study more carefully. Move the
slide around until that portion is in the center of the field of view. With your eye at
the side of the microscope, slowly turn the high power [40X] objective into place.
Again, do not move the stage.
4. Look through the ocular at the letter e. If any adjustment is needed, use only
the fine-adjustment knob. (Note: Always use only the fine-adjustment knob with
high power.) On your drawing of the letter e (above), draw a circle around the
portion of the letter that you are now seeing with high-power magnification.
5. When you have finished your observations of this slide (or any slide), rotate the
nosepiece until the lowest-power objective clicks into place, and then remove the
slide. If it is a purchased slide, return it to the slide tray in which it belongs. If you
have made the slide, put it into the “used slide” containers near the sink.
Activity C: The Size of the Field:
Estimating the Size of Microscopic Organisms
When using each objective, you must know the diameter of the field to estimate the
size of observed objects. Estimate the diameter of field for each magnification of
your microscope as described in the section that follows. Record your results
1.Turn the nosepiece until the scanning (4 X) objective is in position. Place a clear
plastic ruler under the objective so that the millimeter lines go horizontally across
the field of view. The edge of the ruler should extend across the diameter of the
field. Focus on the metric scale with the 4X objective, and move the ruler so one of
the millimeter marks is at the left edge of the field. Estimate and record the
diameter of the field at a total magnification of 40X (remember, 10 X ocular by 4
X objective lens) by counting the spaces and portions thereof visible in the field.
______________mm diameter at 40 X.
2. Use these equations to calculate the diameter of the field at 100X and 400X.
A. Field diameter at 100 X ( in mm)
40X x diameter (mm) at 40 X
B. Field diameter at 400 X ( in mm)
40X x diameter (mm) at 40 X
a. Indicate the estimated diameter of field at each total magnification for your
Diameter of Field
40 X
100 X
400 X
b. Using the diameters determined above, decide on the estimated lengths of
objects that extend across:
1. Two-thirds of the field at 40 X
2. 25% of the field at 100X
3. One-half of the field at 40OX
4. 80% of the field at 40X
(1 mm = 1,000 µm)
Activity D: Depth of Field
When a specimen is prepared for microscopic examination it is very thinly sliced
to allow light to pass through it. Even though the specimen is thinly sliced, it still
has some thickness (think of a piece of cellophane — very thin but it still has some
thickness). The amount of material that is in focus when a specimen is viewed
under the compound microscope is called the depth of field. The higher the
magnification the thinner the depth of focus. This concept is important when
viewing objects under high power because the depth of field is thinner than the
thickness of the specimen. Therefore the lens must be raised and lowered using the
fine adjustment knob to see all levels in the specimen.
Figure 4.2
In Figure 4.2 above, under the 10X lens the letters x, y and z would be in focus
(and magnified 100X); under the 40X lens (400X magnification) only letter y
would be in focus — in order to see x and z clearly the objective lens must be raised
and lowered (respectively) using the fine adjustment knob.
From the demo table obtain a slide labeled silk threads. On the slide are three
differently colored, cross mounted silk threads. The threads are made of many
smaller and finer fibers. With the 10X lens in place, find an area where you can see
all three colored threads at once. Focus very slowly up and down with the fine
adjustment — as each thread comes into focus, you will see the fibers of which it is
made more clearly than those of the other threads. See if you can determine which
thread is on the bottom, which is in the middle and which is on top. Repeat the
process using the high power (40X) lens. You may check your answer with the key
available from your instructor.
Basic Chemistry: Atoms, Molecules, Electrolytes.
Acids, Bases, pH, Buffers. Organic Molecules.
Adapted from Peter Lanzetta, Ph.D.
1. Describe the organization of an atom
2. Define: ionic bond, covalent bond, ion, molecule, electrolyte, non-electrolyte,
acid, base, pH, neutralization, salt, buffer.
3. Given a diagram of the pH scale identify the following: acid region; neutral
point; basic (alkaline) region.
4. State two properties of carbon that enable it to serve as a “backbone” for the
formation of an almost limitless variety or organic molecules.
5. Given the molecular structures, identify the following organic molecules:
saturated and unsaturated hydrocarbons, alcohols, organic acids, amino acids, and
simple sugars.
* Before coming to lab, you should read through all of Lab Topic 3 and the
Basic Chemistry chapter in your textbook.
Introduction, Rationale and Review
As you already know, all things in the world are made of atoms. Atoms in turn are
composed of smaller building blocks called – believe it or not – leptons and quarks
(and even quarks with “flavors”). Our objective today is to gain some working
knowledge and familiarity with a few of the very fundamental chemical concepts.
Today’s exercises are designed to complement in a “hands-on” way your text
readings and lectures. We will not delve deeply into chemical theory, but we will
try to gain some information about atoms and molecules and how they behave
under certain circumstances. Since living things may in part be viewed as chemical
machines, some knowledge of basic chemistry is essential to understanding most
biological phenomena.
There are 92 naturally occurring elements which make up the earth and perhaps the
entire universe. There are more than ten additional elements which have been
created in special atomic reactors. An element is matter made up of identical atoms.
We give each element a different symbol. For instance, oxygen is given the capital
letter O, hydrogen capital H, carbon capital C. It’s easy to see that with 92 elements
we quickly run out of letters, so some elements have symbols which are
abbreviations of their names: magnesium is Mg; chlorine is Cl. Some elements are
given a symbol which is the abbreviation for their Latin or Greek name -sodium in
Latin is natrium, so the symbol for sodium is Na. Please note that the first letter is a
capital and the second is lower case and the two letters together represent one
The structure of the atom is sometimes compared with the layout of the solar
system – at the center of the atom, instead of the sun there is a nucleus. (This is not
to be confused with the nucleus of a cell.) Within the nucleus there are two parts:
protons which are particles with one positive charge and neutrons which, as their
name implies, are neutral. Orbiting the nucleus are electrons — these are much
smaller particles with a negative charge. In an atom, there are always the same
number of electrons as there are protons — therefore atoms are electrically neutral
[that is, they have the same number of positive (+) and negative (-) charges]. An
element is defined by the number of protons in the nuclei of its atoms. The number
of protons is called the atomic number. For example, the element carbon (C) has
an atomic number of 6 because it has 6 protons in its nucleus (and therefore 6
orbiting electrons). Almost all the carbon atoms on this earth also have 6 neutrons
in their nuclei. However some carbon atoms have 4, 5, 7, 8 or 9 neutrons — they are
known as isotopes of carbon. These atoms with the varying number of neutrons are
all carbon regardless of the number of neutrons because they all have 6 protons.
Just about every element exists in nature in various isotopic forms.
Fig. 3.1
As you know from lecture, electrons orbit the nucleus in distinct “shells”. Figure
3.1 shows the electron shells, their names (which are letters starting with K) and
the maximum number of electrons (e) each shell can hold.
Figure 3.2 gives examples of atoms of several different elements and shows the
number and positions of the protons, neutrons and electrons.
Fig. 3.2
Electrolytes and Non-Electrolytes
Atoms interact with each other to form molecules using three kinds of interactions
called chemical bonds. The first of these, hydrogen bonds, do not involve
exchanges of parts of the atoms. However, in both ionic and covalent bonds
electrons are exchanged or shared. We will now examine a relatively simple
chemical interaction that involves exchanges of electrons in the outermost shell the nucleus and inner electrons do not become involved.
Because of their physical properties, the atoms of some elements have the ability to
take electrons from atoms of other elements which tend to give up electrons.
Sodium (Na) is an atom that tends to give up electrons while chlorine (Cl) is an
atom that tends to take electrons. Sodium normally has 11 positively charged
protons and 11 negatively charged electrons and therefore has no charge. Chlorine
has 17(+) and 17(-) and therefore it is also neutral. If chlorine takes one electron
from sodium, it now has 18 negative charges whereas it still has only 17(+) charges
in the nucleus. It therefore has one extra negative charge and is now chlorine with
a negative charge. When an atom has a charge other than neutral it is called an ion.
Chlorine with a negative charge is called chloride ion and the symbol is Cl-.
Similarly, sodium, after giving up an electron has 10 negative charges and 11
positive charges left in its nucleus and it therefore has one extra positive charge -it is now sodium ion, written Na+. The sodium ion and the chloride ion, because
they are oppositely charged, attract one another and form a molecule of sodium
chloride — common table salt. A molecule is two or more atoms or ions bonded
together. The chloride and sodium ions are said to be held together by an ionic
sodium atom (Na):
11 protons (+), 11 electrons
chlorine atom (Cl):
17 protons (+), 17 electrons
one electron has moved from Na to Cl
sodium ion(Na+): 11+,10chloride ion (Cl-): 17+,18Sodium chloride (NaCl) – held together by an ionic bond
Figure 3.3
It is interesting to note that sodium and chlorine atoms exchanged one tiny
negatively charged electron and that exchange turned sodium and chlorine atoms,
which are poisonous, to sodium and chloride ions, which are necessary for life. We
will learn later in the semester some of the ways these two ions function in living
When ionically bonded molecules are placed in water, the ions separate; and the
molecule is said to have dissociated or ionized. When sodium chloride is dissolved
in water, sodium ion separates from chloride ion. Water acts somewhat like a
wedge separating the ions and then surrounding each ion to help keep them
separated. You have undoubtedly discussed this ability of water to act as a solvent
for many compounds.
The ions distributed in the water confer upon water the ability to conduct an
electric current. Pure water cannot conduct electric currents. Therefore all
molecules that ionize in water are called electrolytes, because the free ions allow
the water to conduct electric currents.
molecule is dissolved in water, ions are “pulled” apart
Na+ + Cl-
molecule is dissolved in water, ions are “pulled” apart
Mg++ + 2 (Cl-)
Some atoms do not completely exchange electrons when they react as did sodium
and chlorine. These atoms instead can share electrons in a molecule. Shared
electrons take a position between the two atoms and are held onto by both atoms.
Atoms that form a molecule by sharing electrons are held together by a covalent
bond. Covalently bonded molecules do not ionize in water. The water containing
these molecules does not have charged ions and therefore cannot carry an electric
current. Covalent molecules are called non-electrolytes. In Figure 3.4 we see a
central carbon atom surrounded by four hydrogen atoms. As you will recall, carbon
has an atomic number of 6 — it has 6 protons and 6 electrons. We are concerned
with the outer 4 electrons in the outermost (L) shell, because these do the bonding
with other atoms. Hydrogen has an atomic number of 1 — it has 1 proton and 1
electron. Each covalent bond uses one electron from H, and one from C.
Fig. 3.4 CH4 (methane)
All life characteristics ultimately depend on chemical interactions. In the following
exercises various chemical properties and reactions will be demonstrated.
Exercise 3.1
Work in groups of four. Please work carefully. Record your answers on the
Answer Sheet for Exercise 3.1-3.3.
Fig. 3.5 Light apparatus.
1. Pour enough distilled water into a beaker or plastic cup to half fill it. Dip the
electrodes of the light apparatus into the water in the beaker. Turn the instrument
on by plugging it into the nearest outlet. Does the light go on? Why?
Turn the apparatus off by pulling out the plug. Remove the beaker.
2. Dissolve one half teaspoon of table sugar into a second beaker half full of
distilled water. Dip the electrodes of the light apparatus into the beaker containing
the sugar water. Turn the apparatus on. Does the light go on? Why?
Turn the apparatus off. Remove the beaker with sugar water and place an empty
waste container under the electrodes. Rinse the electrodes off by squirting them
with distilled water from a squeeze bottle (allow the water to collect in the empty
waste container). After rinsing, make sure the apparatus is off, then wipe the
electrodes with a paper towel.
3. Dissolve one half teaspoon of table salt into another beaker with distilled water.
Dip the electrodes into the salt water. Turn the apparatus on. Does the light go on?
Explain your results.
Turn the apparatus off. Again rinse and clean as above.
Acids and Bases
With the apparatus you just used you could test just about anything to see if it is an
electrolyte or not. Among the things that are electrolytes are two special groups of
molecules called acids and bases. Acids are electrolytes. When acids are dissolved
in water, one of the ions formed is hydrogen ion, H+. Bases are also electrolytes.
When bases are dissolved in water, one of the ions formed is hydroxyl ion (OH)-.
Bases are also called alkalies.
Acids are electrolytes which produce hydrogen ion
(H+) in water. Following are some examples of acids:
(hydrochloric acid) HCl ——– > H+ + Cl(hydrofluoric acid) HF ——– > H+ + F(sulphuric acid) H2SO4 ——– > 2 (H+ ) + S04-2
Bases are electrolytes that upon ionization in water give a hydroxyl ion (OH-). The
hydroxyl ion is a complex ion made of one oxygen, one hydrogen atom and one
extra electron. Following are some typical bases:
(sodium hydroxide – lye) NaOH ——– > Na+ + OH(potassium hydroxide) KOH ——– > K+
+ OH+2
+ 2 (OH-)
(magnesium hydroxide) Mg(OH)2 —– > Mg
Exercise 3.2
1. Add 20 drops of 0.1 N HCl to 50 cc of distilled water in a beaker or cup and dip
the electrodes of the light apparatus into the beaker. Turn the apparatus on. What
Turn the apparatus off and rinse and wipe the electrodes as before.
2. Repeat the electrolyte test done on HCl only this time use 0.1 N NaOH. What
happens when the apparatus is turned on?
Turn off the apparatus and rinse and wipe the electrodes.
Hydrogen ions confer acid properties upon water whereas hydroxyl ions confer
basic properties. When equal amounts of both hydrogen and hydroxyl ion are in the
same water the water is neither acidic or basic. Hydrogen ions and hydroxyl ions
react with one another to form water; we call this reaction neutralization. The
remaining ions — those other than the hydrogen and hydroxyl — form a salt. For
HCl ——– >
NaOH ——– >
NaCl (salt)
H2O (water) +
An acid and a base always neutralize one another to water and a salt. What would
be the salt formed if we mixed HF and KOH?
If we were to take an acid and add some to an amount of water there would be a
certain number of hydrogen ions (H+) in that water. If we kept adding acid the
hydrogen ion (H+) concentration would go higher. The same can naturally be said
of the hydroxyl ion (OH-) concentration if we keep adding base to an amount of
water. There is a convenient way of measuring the amount of acidity or basicity
(also known as alkalinity) of a solution; this acid/base measurement scale is called
the pH scale. Look at the diagram below. At pH 1, a solution has a great quantity
of hydrogen ions and no hydroxyl ions. As you go up the scale there are fewer
hydrogen ions and more hydroxyl ions.
ion concentration
ion concentration
pH of Solution
Figure 3.6 pH chart
When you reach seven there is an equal amount of both ions. Seven represents a
neutral solution. If we continue to go up the pH scale the hydrogen ion
concentration continues to fall while the hydroxyl ion concentration continues to
rise until we reach pH 14 where there are virtually all hydroxyl ions and no
hydrogen ions.
The numbers on the pH scale are the negative logarithms of the number of (H+)
ions in a solution. The table below reminds us of how logarithms are related to
numbers. The amount of (H+) ions is very small, so the pH scale starts at 0.1 (H+)
ions. That is pH 1. The fewer the (H+) ions, the larger the pH becomes.
Number Representation and Logarithms
Exponent Notation
Log of the Number
There are several ways to determine the pH of a solution – that is, how acid or
basic a solution is. One method is to use an instrument called a pH meter; another
method is to use specially treated paper strips which turn different colors in
solutions of different pH; a third is to use a computer. The computer records the
pH and can also graph pH changes over time.
Exercise 3.3
1. Pour 25 ml of distilled water into a beaker or cup. Using a glass stirring rod test
the pH of the water by dipping the glass rod into the water and then touching the
rod to the pH test paper. Compare the color of the paper to the chart on the test
paper container and record the pH.
Rinse the glass rod with distilled water and wipe.
2. Using a glass dropper (Pasteur pipet) or plastic dropper, add 10 drops (count
carefully) of 0.1 N HCl to the water, stir and test the pH with the glass rod and test
paper. Record the result.
Rinse and wipe the glass rod.
3. Add 10 drops (again, count carefully) of 0.1 N NaOH to a second beaker
containing 25 ml of distilled water. Be sure the dropper for the NaOH is the same
size as the dropper for the HCl you used in part 2. Test the pH again and record you
Rinse and wipe the glass rod.
4. Now pour the contents of the beaker with the NaOH into the beaker with the
HCl. Stir and test the pH. Record your result.
You can see by the results of your acid and base experiment that water is neutral
but can be made acidic or basic. The acidic and basic solutions can be mixed and
the resulting solution will then be (close to) neutralized. Please realize that to
come to perfect neutrality you must add exactly equal quantities of exactly equal
concentrations of acid and base. Your result in (j) should be somewhere between
your answers to (h) and (i).
Your instructor will tell you whether to continue with the exercises on organic
chemistry (Exercise 3.5), or do the computer exercise on buffers (Exercise 3.4).
EXERCISE 3.1 – 3.3
Exercise 3.1
a) _____________________________________________________________
b) _____________________________________________________________
c) _____________________________________________________________
Exercise 3.2
d) _____________________________________________________________
e) _____________________________________________________________
f) _____________________________________________________________
Exercise 3.3
g) ____________________________________________________________
h) _____________________________________________________________
i) _____________________________________________________________
j) _____________________________________________________________
k) _____________________________________________________________
Exercise 3.5
We’ve just learned that atoms involved in a bond in which electrons were exchanged
are said to be ionically bonded whereas those that share electrons are said to be
covalently bonded. Carbon atoms form a great number of covalently bonded
molecules, and form long chains and rings of carbon atoms to which many other
atoms may be bonded. The area of chemistry that deals with carbon compounds is
called organic chemistry. It is called organic chemistry because it was first
believed that organic molecules could come only from living things, which of
course we know now is not true. Organic chemistry is an immense area of study
because of the almost endless variety of different compounds that can be formed by
carbon and other elements sharing electrons with carbon. We will attempt to learn
only some very basic concepts of organic chemistry, but be sure, organic chemistry
is very relevant to the study of biology. In this section, presentation of a concept
will be followed with a demonstration and/or exercise.
Figure 3.7
Hydrocarbons: Carbon atoms have four electrons in the outer shell that are
available for sharing in covalent bonds. They can share these electrons with many
other kinds of atoms.
Figure 3.8
The left carbon atom is sharing three electrons with three separate hydrogen atoms
and the fourth electron with another carbon atom. This second carbon atom shares
its three remaining electrons with three more hydrogen atoms. For convenience
sake, we will henceforth show a shared pair of electrons as a simple straight line,
and the above molecule will look like the following:
Figure 3.9
We can build very simple to very immense molecules just with carbons and
hydrogens; these molecules are known as hydrocarbons. One of these is a
component of swamp gas: the other, part of a liquid known as gasoline.
H–C–C–C–C–C–C–C –C–H
Figure 3.10
As we add carbon and hydrogen atoms and change configurations, the physical and
chemical properties of the molecule change. The smallest hydrocarbon molecules
exist as gases, while those of intermediate size are liquid and the very large
molecules are solids (for example, margarine or the fat on steaks). All hydrocarbons
may be used as fuels – they burn in the presence of oxygen to form carbon dioxide
(CO2) and water (H2O). The paraffin (“wax”) in candles is a large, solid
Some organic molecules may form branches and rings as below:
Figure 3.11
Alchohols. Molecules similar to hydrocarbons but with some other atoms involved
in the sharing of electrons have major biological significance. For instance, if we
add an oxygen atom into any of the hydrocarbons we have discussed on previous
pages we form a group of molecules known as alcohols.
methane becomes
methyl alcohol or methanol
ethane becomes
Figure 3.12
ethyl alcohol or ethanol
Unsaturated Hydrocarbons. Carbon has another important characteristic. Under
certain circumstances, carbon atoms may share more than one electron with a
neighboring carbon atom (it may even share as many as three electrons). When
carbon shares more than one electron with a neighboring carbon it has that many
less electrons to share with other atoms. Let us look at the organic molecule ethane
in Figure 3.12, above. Each carbon is sharing its four electrons with four different
atoms (three hydrogens and one carbon). The molecule is said to be saturated. If
the two carbon atoms were to share more than one pair of electrons then ethane
would look like the molecule below. Each carbon is sharing electrons with three
other atoms (two hydrogens and one carbon). The molecule is said to be
Figure 3.13
The molecule in Figure 3.13 above is called ethylene. The sharing of two pairs of
electrons by two carbon atoms is called a double bond. The occurrence of
unsaturated (double) bonds in organic molecules has an interesting biological
significance. Large organic molecules like long hydrocarbons and long fatty acids
(about which you will shortly learn) are usually solids, but if any two carbons
among the many in the chain have a double bond, the molecule exists in the liquid
state. See Figure 3.14.
Figure 3.14
saturated tail:solid
monounsaturated:fluid, oily liquid
Some molecules may have several carbon atoms with double bonds and they, too,
are liquids. These molecules with many double bonds are said to be
Organic Acids. The organic acids are among the few organic molecules that will
ionize (partially) in water. Like all acids, they free a hydrogen ion (H+) when they
ionize. Organic acids are hydrocarbons, but have an end carbon with a new electron
sharing pattern.
Figure 3.15
The carbon with the double bonded oxygen and the OH group is the acid (proton
(H+) donor) part of the molecule. We have seen two long organic acids (long
hydrocarbons with the last carbon an acid group), called fatty acids, in Figure 3.14.
They are solids (when they are saturated), and if they have double bonds in the
hydrocarbon portion are called unsaturated fatty acids and are liquids As you have
probably assumed from the name, fatty acids are important constituents in fats.
Amino Acids. The -next group of organic molecules we will study is amino acids.
Amino acids have two important groups sharing electrons with the carbon atoms.
First, there is the acid group we just learned about; secondly, the next carbon shares
electrons with a nitrogen atom which in turn shares electrons with two hydrogen
atoms – this nitrogen with two hydrogens is called an amino group. Hence the
name amino acid (Figure 3.16).
Figure 3.16
NOTE: The parts of the molecule contained in the box – the amino group and the
acid group – are common to all the amino acids. The “R” above refers to any number
of possible combinations of carbon, hydrogen and atoms of other elements.
You can imagine an almost infinite number of different amino acids according to
what atoms and group of atoms (the “R” group) are bonded to the amino acid part.
However, there are only twenty or so biologically important amino acids. These
amino acids are subunits of proteins, important molecules that give structure and
activity (as enzymes) to organisms. Each amino acid has the same basic structure: an
acid group on the first carbon, an amino group on the second, and then a variety of
twenty different molecular arrangements where the diagram shows R. Below are
several examples:
Figure 3.17
Exercise 3.7
Go to the demonstration table and look at the examples and models of various
organic chemical structures. There are extra model kits available; take one to your
desk (one for every four students) and at the end of the period you will be asked to
construct a model of several organic molecules.
Simple Sugars.The last individual group we will study are the simple sugars.
Simple sugars can contain three to seven carbons; the sugar we will talk most about
during the course is the six carbon simple sugar glucose. Don’t let the word simple
fool you, simple sugars are quite complex in structure and cannot be derived from
the chains of hydrocarbons as the other organic molecules were. You will find two
forms representing the glucose molecule in the figure below. You will not be
required to reproduce from memory any of these organic structures – however, given
a diagram of the structures, you will be asked to identity the type of organic
molecule – i.e. whether it is a simple sugar, amino acid, fatty acid, alcohol and so on.
a) straight chain form
b) ring and space-filling ring form
Figure 3.18
Even with all the “OH” groups the sugar molecule is not an alcohol. The top carbon
in the straight chain form is not an acid group: it is known as an aldehyde. The
combination of aldehyde and alcohol groups in the same molecule makes this a new
chemical entity called a simple sugar or monosaccharide.
Exercise 3.8
Work in groups of four.
Butane is a hydrocarbon with a structure like octane, but it has only 4 carbon
molecules and 10 hydrogen molecules. Your instructor will explain the proper use
of the model making kits: please check with him/her after each assembly. With the
molecular model kit construct the following four models. Make models 1-3 in order,
then do model 4.
1 – butane (saturated);
2 – butanol;
3 – butanoic acid (unsaturated carbons at the opposite end from the acid group).
4 – one amino acid of your choice. You may use the ones in the lab book, or get
one from your instructor.
Hint: first draw the molecule in the space below and then construct the model.
Your instructor may modify this list.
NOTE: Please make every effort to utilize lab time to do Exercise 3.8 — it is an
important learning experience.
Exercise 3.9
If you have time, go to the demonstration table labeled Exercise 3.9, and with pH
paper, test the pH of the following items. Record the pH of each.
distilled H20
cleaner with ammonia
We can see that some common household items show a wide range of pH !!!.
Exercise 3.10
Again, if you have time, go to the demonstration table and look at the various
hydrocarbons. Record your observations. With how many are you familiar?
EXERCISE 3.5 – 3.10
Exercise 3.6
l) _____________________________________________________________
m) _____________________________________________________________
n) _____________________________________________________________
o) _____________________________________________________________
Exercise 3.9
p) _____________________________________________________________
q) _____________________________________________________________
r) ____________________________________________________________
Exercise 3.10
s) _____________________________________________________________
Part 2
Living organisms are composed of molecules that come in diverse shapes and
sizes and serve a variety of purposes. Some molecules form the structure of an
organism’s body — for example, the cellulose that makes up the cell walls in
plants, the proteins and phospholipids that comprise cell membranes, and the
fibers that make up animal muscles.
There is also a wide array of molecules that perform all the functions of life. For
example, enzymes catalyze the chemical reactions necessary for biological
processes, neurotransmitters convey information from one brain cell to another,
and visual pigments absorb light so that you can read the words on this page,
In this laboratory you will do further study on three classes of the largest
biological molecules, called macromolecules: carbohydrates, lipids, and proteins.
Your objectives are listed with each set of exercises. A fourth class of
macromolecules, the nucleic acids, will be studied near the end of the semester.
Exercise 4.1: Carbohydrates
Activity A: Monosaccharides and Disaccharides
Activity B: Starch
Exercise 4.2: Lipids
Exercise 4.3: Proteins
Exercise 4.4: Macromolecules in Food
Note: As you do this investigation, please write your answers to the questions at
the end of Lab Topic 4 Part 2.
After completing this exercise, you should be able to
1. Define monosaccharide, disaccharide, and polysaccharide and give examples
of each.
2. Name the monosaccharide components of sucrose and starch.
3. Describe the test that indicates the presence of most small sugars.
4. Describe the test that indicates the presence of starch.
5. Define hydrolysis and give an example of the hydrolysis of carbohydrates.
Most carbohydrates contain only carbon (C), oxygen (O), and hydrogen (H). The
simplest form of carbohydrate molecules are the monosaccharides (“single
sugars”). One of the most important monosaccharides is glucose (C6H1206), the end
product of photosynthesis in plants. It is also the molecule that is metabolized to
produce another molecule, ATP, whose energy can be used for cellular work.
There are many other common monosaccharides, including fructose, galactose,
and ribose.
Some disaccharides (“double sugars”) are also common. A disaccharide is simply
two monosaccharides linked together. For example, maltose consists of two
glucose molecules, lactose (milk sugar) consists of glucose and galactose, and
sucrose (table sugar) consists of glucose and fructose. Can you discern a rule used
in naming sugars?
4.1 a)
Carbohydrates are also found in the form of polysaccharides (“many sugars”),
which are long chains of monosaccharide subunits linked together.
Starch, a polysaccharide composed of only glucose subunits, is an especially
abundant component of plants. Most of the carbohydrates we eat are derived from
plants. What was the last starch you ate?
4.1 b)
Starch is the plant’s way of storing the glucose it makes during photosynthesis.
When you eat starch, you are consuming food reserves that the plant has stored for
its own use. The starch of potatoes and root vegetables, for example, would be
used the next spring for the plant’s renewed growth after the winter die-back. All
perennial plants (those that come up year after year, such as tulips) have some kind
of food storage for overwintering. Beans, on the other hand, contain starch in the
seeds. Beans are annual plants; they will die at the end of the growing season. So
the seeds are stocked with starch to use when they have a chance to germinate the
next spring.
Animals store glucose in glycogen, which is another form of polysaccharide.
Although starch and glycogen are both composed of glucose subunits, the glucose
molecules are bonded together in different ways, so these polysaccharides are not
identical. Glucose subunits are bonded together a third way in the polysaccharide
cellulose. While starch and glycogen are meant to be metabolized for energy,
cellulose, which is the most abundant carbohydrate in the world, is a structural
molecule that is designed not to be metabolized. Cellulose makes up the cell walls
of plants and is a primary component of dietary fiber. For most animals it is
completely indigestible. Those that can digest it, such as termites and cows, do so
only with the assistance of organisms such as bacteria, fungi, or protistans.
Most disaccharides and polysaccharides can be broken down into their component
monosaccharides by a process called hydrolysis, which is accomplished in
organisms by digestive enzymes. This process is important in seeds. If the seed’s
food resource is starch, it must be able to convert the starch to glucose. The
glucose is then used to generate ATP which in turn is used to provide the growing
plant embryo with energy for metabolic work. Hydrolysis of starch begins when
the seed takes up water and begins to germinate.
Germination of barley seeds is part of the process of brewing beer. When the
barley is germinated, the starch-to-sugar conversion begins. In the breakdown of
starch, disaccharide maltose molecules are formed before the final product,
glucose, is obtained. At a certain point in the germination, the barley is dried so
that no further hydrolysis takes place. The maltose sugar is extracted and used in
the brewing process. That’s the “malt” listed on the beer can as an ingredient. The
process of germinating the barley is called malting.
A chemical hydrolysis can be done in the laboratory by heating the molecules with
acid in the presence of water. You will perform a chemical hydrolysis in this
Wear safety glasses throughout the lab session.
Activity A: Monosaccharides and Disaccharides
You will use Benedict’s reagent as a general test for small sugars
(monosaccharides and disaccharides). When this reagent is mixed with a solution
containing single or double sugars and then heated, a colored precipitate (solid
material) forms. The precipitate may be yellow, green, orange, or red. If no
monosaccharide or disaccharide is present, the reaction mixture remains clear.
However, Benedict’s reagent does not react with all small sugars. For example,
sucrose gives a negative Benedict’s reaction.
Glucose will be used in this laboratory to demonstrate a positive Benedict’s test
(Figure 4. 1). What should be used as a negative control for this test?
4.1 c)
Figure 4.1. Benedict’s test for detecting small sugars.
1. Make a boiling water bath by filling a beaker about half full of water and
heating it on a hot plate. Put six or seven boiling chips in the beaker. You will
need to use this water bath in several activities.
Set the hot plate where it will not be in your way as you work. Be careful-it
will be very hot!
2. Get two test tubes and label them 1 and 2 with a wax pencil.
Make heavy marks so that they don’t melt off in the water bath.
3. Put one dropperful (NOTE: NOT drop) of glucose into Tube 1. Tube 1 is the
positive control.
4. Tube 2 is the negative control. What substance goes in it? How much should
be used?
4.1 d)
5. Add 2 droppersful of Benedict’s reagent to each tube.
6. Place the tubes in the boiling water bath and let them heat for 5 minutes.
7. After 5 minutes, remove the tubes from the water bath.
Use a test-tube holder to retrieve test tubes from the boiling water.
8. Allow the tubes to cool at room temperature for several minutes in the test
tube rack while you go on to the next procedure.
9. Record your observations (color?).
4 e) Tube 1 (glucose).
4 f) Tube 2 (negative control)
Interpretation of Results
Describe a positive Benedict’s test.
What are the limitations of this test?
4.1 g)
4.1 h)
Activity B: Starch
Starch is tested by using iodine reagent (I2KI – iodine potassium iodide). A dark
blue color indicates the presence of starch (Figure 4.2).
You will use a solution of potato starch to demonstrate a positive test. What
negative control should be used for this test?
4.1 i)
Figure 4.2. The iodine test for detecting starch.
1. Get two test tubes and label them 1 and 2.
2. Put a dropperful of starch solution in Tube 1. This is the positive control.
3. Tube 2 is the negative control. What substance goes in it? How much should
be used? Why?
4.1 j)
4. Put 3 or 4 drops of iodine reagent into each tube.
5. Record the results.
4.1 k) Tube 1 (starch):
4.1 l) Tube 2 (negative control):
Interpretation of Results
Describe a positive test for starch.
What are the limitations of this test?
4.1 m)
4.1 n)
After completing this exercise, you should be able to
1. Define lipid and give examples.
2. Describe the test that indicates the presence of lipids.
Lipids are compounds that contain mostly carbon and hydrogen. They are grouped
together solely on the basis of their insolubility in water. The lipids we will
consider in this laboratory are fats and oils, which are generally used as storage
molecules in both plants and animals. You are no doubt already familiar with the
fact that your body converts excess food into fat. This fat is stored in your adipose
tissue until your food intake is lower than your metabolic needs, at which time the
fat can be metabolized to generate ATP, whose energy can be used for cellular
work. Plants, too, can store fats. Seeds are often provisioned with fats that can be
metabolized by the developing embryo when germination time comes. Thus we
obtain corn oil, peanut oil, sunflower oil, and others by pressing the seeds.
You will use the paper test (Figure 4.4) to indicate the presence of lipids in
various foods. Although this test is not very sophisticated, it is quick and
Rub sample on
brown paper
Figure 4.4. Brown paper test for lipids.
1. Get a small square of brown paper. Write “oil” on one half and “water” on the
2. Put a tiny drop of salad oil on the half of the paper labeled oil. Rub it gently
with your fingertip.
3. As a negative control, put a tiny drop of water on the half of the paper labeled
water. Rub it gently with a different fingertip to avoid contamination.
4. Allow the spots to dry. This may take quite a while, so go on to another
exercise while you wait.
5. When the spots are dry, hold the paper up to the light.
Interpretation of Results
Describe a positive test for lipids.
What are the limitations of this test?
4.2 a)
4.2 b)
After completing this exercise, you should be able to
1. Define protein and give examples.
2. Explain why the structure of a protein is important for its function.
3. Describe the test that indicates the presence of protein.
A protein’s structure is determined by the amino acid subunits that make up the
molecule. Although there are only 20 different naturally occurring amino acids,
each protein molecule has a unique sequence. The amino acids are linked by fairly
tight bonds, and the side groups (R groups) that are part of the amino acids also
interact with each other to help shape the molecule.
Proteins have a greater diversity of roles than either carbohydrates or lipids. The
shape of a protein is key to its purpose: Proteins work by selectively binding to
other molecules.
You will use biuret reagent as a test for proteins (Figure 4.5). This reagent, which
is blue, reacts with proteins to give a light violet or lavender color.
You will use a solution of egg albumin (a protein extracted from egg whites) to
demonstrate a positive biuret test. What negative control should be used for this
4.3 a)
Figure 4.5. Biuret test for protein.
1. Get two test tubes and label them 1and 2.
2. Put two droppersful of egg albumin into Tube 1.
3. Tube 2 is the control. What substance goes in it? How much should be used?
4.3 b)
4. Put 1 dropperful of biuret reagent into each tube and swirl gently to mix.
5. After 2 minutes, record the color in each tube.
4.3 c) Tube 1 (egg albumin):
4.3 d) Tube 2 (negative control):
Interpretation of Results
Describe a positive biuret test.
4.3 e)
What are the limitations of this test?
4.3 f)
Macromolecules in Food
After completing this exercise, you should be able to
1. Interpret the results of tests that indicate the presence of sugar, starch, lipid,
and protein in food samples.
We metabolize food in order to release energy to produce the ATP needed for
cellular work. We also break down food molecules in order to use their subunits as
raw materials for synthesizing our own macromolecules. In this exercise, you will
investigate certain foods to learn which macromolecules are present in each.
Activity: Tests with Food
Test some or all of the items in Table 4.2 for the presence of simple sugars, starch,
lipid, and protein. Your instructor may want to modify the list. The procedures for
the tests are reviewed below.
Benedict’s Test (sugar)
Put 1 pasteur pipetful of sample into a test tube. Add 2 droppersful of Benedict’s
reagent; mix. Heat in a boiling water bath for 5 minutes. Allow to cool and
observe the precipitate.
*Some samples may require extra cooling time, so don’t be too hasty in
recording results.
Iodine Test (starch)
Put a pipetful of sample into a test tube and add 4 or 5 drops of iodine reagent;
*In some foods, the starch is still contained in granules inside the cells. You may
see these dark granules suspended in the yellow solution instead of seeing the
entire solution turning blue.
Paper Test (lipid)
If the sample is whole (for example, a peanut), rub a piece of it directly on the
paper. If the sample is liquid, put a small drop on the paper.
*Remember to wait for the paper to dry before you record the results.
Biuret Test (protein)
Put 1 pipetful of sample into the test tube and add 1 dropperful of biuret reagent;
*Allow at least 2 minutes for the reaction to occur. Some samples may take 5
minutes to react.
Some of the foods to be tested are solids. Use a razor blade to mince
approximately 1 cm3 (about the size of a pea) of the sample. Put it in a test tube
with 10 mL distilled water. Put your thumb over the top of the test tube and shake
it vigorously for 1 minute. Perform the tests using the liquid (except the lipid test).
Record your results in Table 4.2. Be sure to rinse off the razor blade and cutting
board between samples to avoid contamination.
Table 4.2 (at end of Lab Topic 4 Part 2)
Interpretation of Results
Which results confirmed your previous knowledge about the composition of
4.4 a)
Which results were unexpected?
4.4 b)
What factors might result in a false negative test (that is, the food does contain a
molecule but the tests results are negative)?
4.4 c).
Why might a plant storage organ (such as a fruit or tuber) contain both starch and
4.4 d)
If you have tested foods in addition to the ones listed in Table 4.2, compare the
results from those tests with the results for the foods listed in Table 4.2.
At the end of the Answer Sheet, please answer the Questions for Review
This laboratory is from Jean Dickey, Laboratory Investigations for Biology 2nd ed.
San Francisco: Benjamin Cummings, 2003.
Procedures for the macromolecule tests were adapted from the following sources:
Armstrong, W D., and C. W Carr. Physiological Chemistry Laboratory Directions,
3rd ed. Minneapolis: Burgess Publishing, 1963.
Dotti, L. B., and J. M. Orten. Laboratory Instructions in Biochemistry, 8th ed.
St. Louis: C. V Mosby, 1971.
Oser, B. L., ed. Hawk’s Physiological Chemistry, 14th ed. New York:
McGraw-Hill, 1965.
Name: _______________________ Section:____________Date:__________
Answer Sheet
Exercise 4.1 Carbohydrates
4.1 A
4.1 B
4.1 C
4.1 D
4.1 E
4.1 F
4.1 G
4.1 H
Activity B: Starch
4.1 I
4.1 J
4.1 K
Name: _______________________ Section:____________Date:__________
4.1 L
4.1 M
4.1 N
Exercise 4.2 Lipids
4.2 A
4.2 B
Exercise 4.3 Proteins
4.3 A
4.3 B
4.3 C
4.3 D
4.3 E
4.3 F
Name: _______________________ Section:____________Date:__________
Exercise 4.4 Macromolecules in Food
Table 4.2
Use + for positive, – for negative, or +/- if inconclusive
4.4 A)
4.4 B)
4.4 C)
4.4 D)
Name: _______________________ Section:____________Date:__________
Questions for Review
1. What subunits make up
a. Carbohydrates?
b. Proteins?
2. Why is each test done initially using water as well as a known sample?
3. Why might a substance taste sweet, yet give a negative reaction with the
Benedict’s test?
What procedure could you use to check your answer to the previous question?
4. You have been given an unknown solution. Describe how you would test it,
and what the positive result would be, for the presence of
a. Starch:
b. Lipid:
c. Sugars:
d. Protein:
5. You have tested an unknown sample with biuret and Benedict’s reagents. The
solution mixed with biuret reagent is blue. The solution boiled with Benedict’s
reagent is also blue. What does this tell you about the sample?
6. Whole butter gives only a slightly positive test for protein (and may show no
reaction at all). When the same butter is clarified, however, the liquid lower layer
is definitely positive for protein. Explain why these different results might have
been obtained.
7. Since potatoes have starch in them, why don’t they taste sweet after they are
Lab Topic 1
The Metric System
Measurements and Laboratory Equipment
Peter Lanzetta, Ph.D., Georgia Lind, Ph.D., Mary Ortiz, Ph.D.
After completing this exercise, you will be able to:
1. Define length, volume, meniscus, mass, density;
2. Recognize and be able to properly use common laboratory equipment: graduated cylinders, beakers,
Erlenmeyer flasks, pipettes, a triple beam balance, thermometers, etc.;
3. Measure and estimate length, volume, mass and temperature in metric units;
4. Demonstrate the proper method for determining an accurate measurement;
5. Explain the concept of temperature;
6. Explain the advantages of the metric system of measurement over the English System.
One requirement of the scientific method is that results be repeatable. As numerical results are
more precise than purely written descriptions, scientific observations are usually made as measurements.
Of course, sometimes a written description without numbers is the most appropriate way to describe a
The Metric System
Logically, units in the ideal system of measurement should be easy to convert from one to
another (for example, inches to feet or centimeters to meters) and from one related measurement to
another (length to area, and area to volume). The metric system meets these requirements and is used by
the majority of citizens and countries in the world. Universally, science educators and researchers prefer
it. In most non-metric countries, governments have launched programs to hasten the conversion to
metrics. In fact, the U.S. Department of Defense adopted the metric system in 1957, and all cars made in
the United States have metric components. As expected there has been some reluctance on the part of
many Americans to change over – some reasons are economic (having to retool factories and
households); normal resistance to change and probably some feeling that the metric system is foreign in
more ways than one.
It is interesting to take some historical notice of how the metric system came into existence. Prior
to the French revolution there was little international agreement on standard weights and measures. Each
country had its own standards based on some transient things like the length of the king’s foot or middle
knuckle. Just before the French revolution the French Academy of Sciences decided to work on a system
of measures based on universal, scientific principles. They decided that the system should be based on a
unit of length that was one ten millionth of the distance from the North Pole to the equator. That unit
was called the meter. The revolutionists, when they came into power, agreed with the change because it
represented a further repudiation of the old ways and another step towards “reason”. Napoleon gave
several awards to the scientists and mathematicians who helped devise the new system (including one to
Joseph Lagrange who was lucky not to be beheaded in the revolution due to his association with the
monarchy). The advantage of metric measurements was quickly appreciated by scientists and others
(including merchants, mechanics etc.) from other countries and it spread from France and was eventually
adopted almost everywhere. In 1819 an English physicist and chemist, William Wollaston, argued
against Britain’s adopting the system and as a result Britain and the commonwealth countries declined to
convert to its use. Unfortunately, the United States followed Britain’s lead and also declined – we have
been saddled with the irrational English system of weights and measures ever since (ironically the
British and the other commonwealth countries have since converted but, alas, we are taking an
inordinate amount of time to do so).
The metric reference units are the meter for length, the liter for volume, the gram for mass, and
the degree Celsius for temperature. Regardless of the type of measurement, the same prefixes are used
to designate the relationship of a unit to the reference unit. Table 1.1 lists the prefixes we will use in this
and subsequent exercises. The metric system is a decimal system of measurement (based on ten (deci)).
Metric units are 10, 100, 1000, and sometimes 1,000,000 or more times larger or smaller than the
reference unit. Thus, it’s relatively easy to convert from one measurement to another either by
multiplying or dividing by 10 or a multiple of 10.
TABLE 1.1 Prefixes for Metric System Units
Prefix of Unit (Symbol)
nano (n)
micro ()
milli (m)
centi (c)
kilo (k)
Part of Reference Unit
1/1,000,000,000 = 0.000000001 = 10-9
1/1,000,000 = 0.000001 = 10-6
1/1000 = 0.001 = 10-3
1/100 = 0.01 = 10-2
1000 = 103
There are additional prefixes, but these will suffice for now for our work in Biology. A prefix in front of
a unit tells you how many of that unit you have. For example:
1. cm
centimeters (hundredths of a meter)
2. kg
(thousands of grams)
3. ml
(thousandths of a liter)
In this set of exercises, we examine the metric system and compare it to the American Standard system
of measurement (feet, quarts, pounds, and so on).
Per lab room:

source of distilled water (dH20)
metric bathroom scale
hot plate
boiling chips
thermometer on holder
Per student group of two-four.

30-cm ruler with metric and American
(English) Standard units on opposite edges
250-ml beaker made of heat-proof glass
250-ml Erlenmeyer flask
3 graduated cylinders:10-ml, 25-ml, 100-ml
l-quart jar or bottle marked with a fill line
one-piece plastic dropping pipette (not
graduated) or Pasteur pipette and bulb
graduated pipette and safety bulb or filling

1-pound brick of coffee
ceramic coffee mug
1-gallon milk or water bottle
metric tape measure
1-l measuring cup
non-mercury thermometer(s) with
Celsius (oC) and Fahrenheit (oF)
scales (about 220-110 oC)
a triple beam balance
A. Length ( 15 min.)
Length is the measurement of a real or imaginary line extending from one point to another. The
standard unit is the meter, and the most commonly used related units of length are
1000 millimeters (mm) = 1 meter (m)
100 centimeters (cm) = 1 m
1 kilometer (km) = 1000 m
For orientation purposes, the yolk of a chicken egg is about 3 cm in diameter. Since the difference
between metric units is based on multiples of 10, it’s fairly easy to convert a measurement in one unit to
another. Before we do that, let’s review a few simple rule from mathematics: (note: “N” means any
Rule 1:
Any number multiplied by one equals the original number and doesn’t change its
value. That is: N x 1 = N
For example:
153 x 1 = 153
Rule 2:
Any number divided by itself equals one, or
For example:
7/7 =1
x/x =1
Rule 3:
When multiplying units expressed as fractions, units cancel like numbers.
For example:
Here, the kg and cg cancel and we are left with mg, as follows:
= mg
N/N = 1
Now we are ready to perform metric conversions. We will do this by multiplying the given value by one
or more conversion factors to get the value in the desired units. Doing an example best shows this:
1a. Convert 8 km into m:
We know that 1 km = 1000 m. By dividing this equation by 1000 m, we get:
1 km
1000 m
1000 m
1000 m
1 km
1000 m
This can also be written:
1000 m
1 km
Recall, multiplying by “1” does not change the value of the original number. Therefore, we can write:
8 km x
1000 m
1 km
8 km
1000 m
1 km
8000 m
(Why did we use 1000m/1 km and not 1 km/1000m for our conversion factor?)
1b. Convert 6 cm into m:
We know that 1 m = 100 cm. By using the form of one, 100 cm / 1 m, we can multiply as
6 cm
1 m
6 m
0.06 m
100 cm
Try the following:
a) 5 mm = ______m
b) 5.5 m = ______ µm
c) 9 km = _______nm
(Hint: the last example requires two conversion factors.)
2. Measure the length of this page in centimeters to the nearest tenth of a centimeter with the metric
edge of a ruler. Note that nine lines divide the space between each centimeter into 10 millimeters.
The page is _______________ cm long.
Calculate the length of this page in millimeters, meters and kilometers.
______________m _____________ km
Now repeat the above measurement using the English side of the ruler. Measure the
length of this page in inches.
Convert your answer to feet and then yards.
_________________ ft
_____________ yds
Explain why it is much easier to convert units of length in the metric system than in the English
System _____________________________________________________________________
B. Volume (20 min.)
Volume is a measure of the space an object occupies. The metric standard unit of volume is the liter (l),
and the most commonly used subunit, the milliliter (ml). There are 1000 ml in 1 liter. A chicken egg has
a volume of about 60 ml.
The volume of a box is the length (l) x the width (w) x the height (h) (Figure A – 1).
The amount of water contained in a cube with sides 1 cm each in length is 1 cubic centimeter (cc)
(1 cm x 1 cm x 1 cm) which for all practical purposes equals 1 ml (Fig A – 2).
1. How many milliliters are there in 1.7 l?
How many liters are there in 1.7 ml?
V=l x w x h
Figure A – 1 Determining the volume
of a box.
Figure A-2 Relationship between the units of length,
volume, and mass in the metric system.
2. Use Figure A – 3 to recognize graduated cylinders, beakers, Erlenmeyer flasks, and the different
types of pipettes. Some of these objects are made of glass; some plastic. Some will be calibrated in
milliliters and liters; others will not. (“Graduated” means the container is marked with volume
Figure A – 3 Apparati commonly used
to measure volume: (a) pipette filling
device, (b) pipette safety bulb,
(c) Pasteur pipette and bulb,
(d) Erlenmeyer flask,
(e) glass graduated cylinder,
(f) plastic graduated cylinder,
(g) plastic dropping pipette,
(h) beaker, (i-k) graduated pipettes.
(Photo by D. Morton and J. W. Perry.)
3. Pour some water into a 100-ml graduated cylinder and observe the boundary between fluid and air,
the meniscus. Surface tension makes the meniscus curved, not flat. The high surface tension of water is
due to its cohesive and adhesive or “sticky” properties. Draw the meniscus in the plain cylinder outlined
in Figure A-4. The correct reading of the volume is at the lowest point of the meniscus.
Figure A – 4.
Draw a
meniscus in this
plain cylinder.
4. Using the 100-ml graduated cylinder, pour water into a 1-quart jar or bottle. About how many
milliliters of water are needed to fill the vessel up to the line? __________________ml.
5. Pipettes are used to transfer small volumes from one vessel to another. Some pipettes are not
graduated (for example, Pasteur pipettes and most one-piece plastic dropping pipettes); others are
(a) Fill a 250-ml Erlenmeyer flask with distilled water.
(b) Use a plastic dropping pipette or Pasteur pipette with a bulb to withdraw some water.
(c) Count the number of drops needed to fill a 10-ml graduated cylinder to the 1-ml mark.
Record this number in Table A-2.
(d) Repeat steps b and c two more times and calculate the average for your results in Table A-2.
(remember, to find an average, add all the results together, and divide by the number of
(e) Explain why the average of three trials is more accurate than if you only do the procedure
Estimate of the Number of Drops in 1 ml
C. Mass (25 min.)
Mass is the quantity of matter in a given object. The standard unit of mass is the kilogram (kg), and
other commonly used units are the milligram (mg) and gram (g). There are 1,000,000 mg in 1 kg and
1000 g in 1 kg. A chicken egg has a mass of about 60 g. Note that the following discussion avoids the
term weight. Mass is a constant (scalem). Your mass on the Earth is the same as your mass on the Moon.
However, since gravity on the Moon is 1/6 of that on Earth, your weight on the Moon would be less. For
example, a 60 lb. person on earth would weight 10 lbs. on the Moon, but that person would have the
same mass. However for our purposes, we will use mass and weight interchangeably.
1. How many milligrams are there in 1 g?
Convert 1.7 g to milligrams and kilograms.
_________________ kg
2. A 1-cc cube, if filled with 1 ml of water, has a mass of 1 g (Figure A-1). The mass of other materials
depends on their density (water is defined as having a density of 1). The density of any substance is its
mass divided by its volume.
Approximately how many liters are present in 1 cubic meter (m3) of water? Since each of the sides of a
cubic meter ( m3) is 100 cm in length, it’s easy to calculate the number of cubic centimeters (that is, 100
cm X 100 cm X 100 cm = 1,000,000 cc). Now just change cubic centimeters to milliliters and convert
1,000,000 ml to liters.
1,000,000 ml = _________ l
What is its mass in kilograms?______________________ kg.
3. Determine the mass of an unknown volume of water (1). Mass may be measured with a triple beam
balance, which gets its name from its three beams (Figure A-5). In this device a movable mass hangs
from each beam.
(a) Slide all of the movable masses to the
left to zero. Note that the middle and
back masses each click into the
Figure A – 5 Triple Beam Balance
leftmost notch.
movable masses
10-g graduations
(b) Clear the pan of all objects and
100-g graduations
make sure it is clean.
0. 1 -g and 1 -g
(c) The balance marks should line up,
indicating that the beam is level and
that the pan is empty. If the balance
marks don’t line up, rotate the zero
adjust knob until they do.
(d) Place a 250-ml beaker on the pan.
The right side of the beam should
rise. Slide the mass on the middle beam
until it clicks into the notch at the
100-g mark. If the right end of the
beam tilts down below the stationary
balance mark, you have added too
much mass. Move the mass back a
notch. If the right end remains tilted up,
additional mass is needed. Add
zero adjust knob
balance marks
increments until the beam tilts down;
then move the mass back one notch.
Repeat this procedure on the back
beam, adding 10 g at a time until the beam tilts down, and then backing up one notch. Next, slide the
front movable mass until the balance marks line up.
(e) The sum of the masses indicated on the three beams gives the mass of the beaker. Unnumbered lines
divide the space between the numbered gram markings on the front beam into 10 sections, each
representing 0.1 g. Record the mass of the beaker to the nearest tenth of a gram in Table A-3.
Figure 2-4 Triple beam balance. (Photo by D. Morton and J.
W. Perry.)
(f) Add an unknown amount of water and repeat the above procedure. Record the mass of the beaker
and water in Table A-3.
(g) Calculate the mass of the water alone by subtracting the mass of the beaker from that of the
combined beaker and water. Record it in Table A-3.
TABLE A-3 Weighing an Unknown Quantity of Water with a
Triple Beam Balance
Masses (g)
Beaker and water
(h) Now measure the volume of the water in milliliters with a graduated cylinder.
What is the volume?__________________ ml
4. Using the triple beam balance, determine the mass (that is, weight) of a brick of coffee in grams.
(1) Modified from C M Wynn and G. A. Joppich, Laboratory Experiments for Chemistry, A Basic
Introduction, 3rd ed. Wadsworth, 1984.
D. Estimating Length, Volume, and Mass (10 min.)
Now that you have experience using metric units, let’s try estimating the measurements of some
everyday items.
1. Estimate the length of your index finger in centimeters. ___________ cm
2. Estimate your lab partner’s height in meters.______________m
3. How many milliliters will it take to fill a ceramic coffee mug?______________ml
4. How many liters will it take to fill a gallon plastic bottle?_______________l
5. Estimate the weight of some small personal item (for example, loose change) in grams.
6. Estimate your or your lab partner’s weight in kilograms. ______________________kg
7. Record your estimates in Table A-4.
8. Now, check your results using either a ruler, metric tape measure, 100-ml, graduated cylinder, 1-l
measuring cup, triple beam balance, or metric bathroom scale, recording your measurements in Table
A-4. Complete Table A-4 by calculating the difference between each estimate and measurement.
Differences Between Estimates and Measurements
Estimate – Measurement
(How good are you at estimating?)
E. Temperature (About 20 min.)
The degree of hot or cold of an object is termed temperature. More specifically, it is the average
kinetic energy of molecules. Heat always flows from high to low temperatures. This is why hot objects
left at room temperature always cool to the surrounding or ambient temperature, while cold objects
warm up. Consequently, to keep a heater hot and the inside of a refrigerator cold requires energy.
Thermometers are instruments used to measure temperature.
We in the United States are accustomed to the Fahrenheit scale to measure temperature. When
you realize that Fahrenheit, over 200 years ago, created the lowest temperature he could (by mixing
equal parts snow and salt) and called that temperature zero, it is easy to see why this scale is so far out of
date. Water on Fahrenheit’s scale freezes at 32 and boils at 212– advanced for his time, antiquated and
almost silly for ours. This discourse is not a “put down” of Fahrenheit; to the contrary he was a very
bright man who pioneered the use of mercury in thermometers because of its constant rate of expansion
over temperature ranges. We are merely saying that our clinging to a cumbersome, irrational system
based on 200 year old technology is unfortunate.
Celsius took pure water at sea level and said it freezes at 0 and boils at 100 and he put 100
divisions between.
The Fahrenheit scale has 180 divisions between the freezing and boiling points of water (212 –
32) whereas Celsius has 100. Therefore, the Celsius degrees are almost twice as big as those of
Fahrenheit. 180F to 100C = 1.8 to 1 = 9 to 5.
1. Using a thermometer with both Celsius (oC) and Fahrenheit (oF) scales, measure room temperature
and the temperature of cold and hot running tap water. Record these temperatures in Table A-5.
2. Fill a 250-ml, beaker with ice about three-fourths full and add cold tap water to just below the ice.
Wait for 3 minutes, measure the temperature, and record it in Table A-5. Remove the thermometer
and discard the ice water into the sink.
3. Observe as the instructor fills a beaker with warm tap water to about three-fourths full and adds three
boiling chips. A thermometer holder will be used to clip a thermometer onto the rim of the beaker so
that the bulb of the thermometer is halfway into the water. The instructor will boil the water in the
beaker by placing it on a hot plate. After the water boils, record its temperature in Table A-5.
CAUTION: Do not touch the hot beaker, the boiling water, or the edges of the hot plate.
4. To convert Celsius degrees to Fahrenheit degrees, multiply by 9/5 and add 32.
Is 4oC the temperature of a hot or cool day?
What temperature is 4oC in degrees Fahrenheit? ________________oF
5. To convert Fahrenheit degrees to Celsius degrees, subtract 32 and multiply by 5/9.
What is body temperature, 98.6 oF, in degrees Celsius? ___________________oC
6. In summary, the formulas for these temperature conversions are:
9/5 (oC) + 32
water freezes
refrigerator temperature
room temperature
body temperature
water boils
Table A-5
= 5/9 (oF – 32)
Comparison of Celsius and Fahrenheit Temperatures
Cold running tap water
Hot running tap water
Ice water
Boiling water
Lab Topic 1B_________________________________________________
Metric System Conversions: An Alternative
M. Lakrim, Ph. D.
Activity 1: Conversion within the Metric System
To convert units to multiples or to fractions, use the following examples and table:
Example 1: Fractions (larger to smaller)
Convert 1 gram into milligrams (1 g = x mg) [See table below].
Write the number “1” in the column “Units–gram”
2nd: Fill each column with zero “0” until you reach the desired column “milli”
3rd: Read the whole number and report it on your answer sheet
1 g = 1,000 mg
Example 2: Multiples (small to larger)
Convert 12.5 meters into kilometers (12.5 m = x km) [See table below].
Write the number “12.5: on the column “Units–meter”.
Put only one figure at each column:
“2” in column “Units–meter”
“1” in column “deca”
“5” in column “deci:
2nd: Fill the remaining columns up with zeroes until you reach the desired
column “kilo”
3rd: Read the whole number and report it on your answer sheet
12.5 m = 0.0125 km
gram, liter, meter
Exercise 1. Conversions
Using the table below, convert the following:
12 g =
2 km =
3.5 km =
13.5 l =
20 ml =
30 mg =
100 dl =
Gram, Liter, Meter
The Process of Scientific Inquiry
From Jean Dickey, “Laboratory Inverstigations for Biology” (2nd ed. 2003)
* Before to coming to lab, you should read through all of Lab Topic 2A
Scientific inquiry is a particular way of answering questions. It can’t be used for
all types of questions. The questions that can be answered by science must meet
specific guidelines and scientific investigations must be carried out using certain
rules. When an investigation is designed properly and meets these guidelines,
then the results are acceptable to other scientists and are added to the body of
scientific knowledge. If an investigator cannot show that his or her experiment
was done according to the guidelines, then the results of that experiment will not
be recognized as valid by other scientists.
The purpose of such guidelines can be understood by comparing them to sports
records. For example, a new record set in a track and field event only counts if
the meet was approved by the governing body that sets the guidelines. The site
and equipment used are scrutinized to be sure that they are within the regulations
and the athlete is tested for use of illicit substances. Only when these required
conditions are met is the record certified as valid.
In this laboratory you will learn about the basic elements of scientific inquiry
and how to apply this process to solving problems.
Exercise 2.1: The Black Box
Exercise 2.2: Defining a Problem
Exercise 2.3: The Elements of an Experiment
Exercise 2.4: Designing an Experiment
The Black Box
After completing this exercise, you should be able to
1. Explain the scientific inquiry method, which you apply to various examples
in this exercise.
You will use the “black box” exercise as a model of how scientific inquiry is
carried out. Each lab team has a container with one or more objects sealed inside.
Each team also has an empty container of the same type and a plastic bag
holding objects that might be inside the sealed container. Your task is to devise a
way to find out what is in the box without opening it. The steps listed below
give you some idea of how to proceed. Answer the questions to keep a record of
what you did.
1. Make observations. Investigate the container by any means available to you
except opening the container.
What are your observations?
How did you make your observations?
What other methods that are not available to you right now might be used to
make observations?
Why is making observations an important first step in solving this problem?
2. Make a guess about the contents of the box.
What did you base your guess on?
3. For now, you still can’t open the sealed container. How can you test whether
your guess is correct?
4. Use the method you described above to check your guess. Record your
results below Was your guess correct? How sure are you?
5. If you aren’t sure you know yet what is in the box, what should you do next?
6. Short of opening the box, what’s the best you can do to find out what’s in it?
7. Suppose you tell your instructor what you have concluded is in the box, and
he or she says that you are wrong. What are some things that could have led you
to make the wrong conclusion?
8. Summarize the methods you used to solve the problem of the black box.
The steps you used to determine the contents of the black box are similar to the
procedure followed in one type of scientific investigation. The investigator poses
a question-for example, “What is in the box?” From the question and preliminary
observations, the investigator makes an educated guess (known as a hypothesis)
about the answer. She then devises an experiment to test the hypothesis,
performs the experiment, and draws a conclusion from its results. The hypothesis
may be revised, and further experiments may be done if the results are not
conclusive. Eventually the investigator reaches a point where she is confident
that her conclusions are correct.
In Exercises 2.2 and 2.3 you will learn to recognize the elements of a good
scientific investigation. In this and later laboratories you will design your own
Defining a Problem
After completing this exercise, you should be able to
1. Identify questions that can be answered through scientific inquiry and explain
what characterizes a good question.
2. Identify usable hypotheses and explain what characterizes a good scientific
Every scientific investigation begins with the question that the scientist wants to
answer. The questions addressed by scientific inquiry are based on observations
or on information gained through previous research, or on a combination of both.
Just because a question can be answered doesn’t mean that it can be answered
scientifically. Discuss the following questions with your lab team and decide
which of them you think can be answered by scientific inquiry.
What is in the black box?
Are serial killers evil by nature?
What is the cause of AIDS?
Why is the grass green?
What is the best recipe for chocolate chip cookies?
When will the Big Earthquake hit San Francisco?
How can the maximum yield be obtained from a peanut field?
Does watching television cause children to have shorter attention spans?
How did you decide what questions can be answered scientifically?
A scientific question is usually phrased more formally as a hypothesis, which is
simply a statement of the scientist’s educated guess at the answer to the question.
A hypothesis is usable only if the question can be answered “no.” If it can be
answered “no,” then the hypothesis can be proven false. The nature of science is
such that we can prove a hypothesis false by presenting evidence from an
investigation that does not support the hypothesis. But we cannot prove a
hypothesis true. We can only support the hypothesis with evidence from this
particular investigation. For example, you used hypotheses to investigate the
contents of your sealed box. A reasonable hypothesis might have been, “The
sealed box contains a penny and a thumbtack.” This hypothesis could be proven
false by doing an experiment: putting a penny and a thumbtack in a similar box
and comparing the rattle it makes to the rattle of the sealed box. If the objects in
the experimental box do not sound like the ones in the sealed box, then the
hypothesis is proven false by the results of the experiment, and you would move
on to a new hypothesis. However, if the two boxes do sound alike, then this does
not prove that the sealed box actually contains a penny and a thumbtack. Rather,
this investigation has supplied a piece of evidence in support of the hypothesis.
You could test almost any hypothesis you made by putting objects in the empty
box. What one hypothesis could not be proven false by experimentation?
You may now open the sealed container. Was your final conclusion about
its contents correct.?
If your conclusion has now been disproven, explain how you reached an
erroneous conclusion. (You may have found that your conclusion was wrong in
spite of accurate observations and careful experimentation. Conclusions reflect
the best evidence available at the time.)
Can you think of any areas of scientific inquiry where a new technology or
technique might challenge or disprove hypotheses that are already supported by
experimental evidence?
The scientific method applies only to hypotheses that can be proven false
through experimentation (or through observation and comparison, a different
means of hypothesis testing). It is essential to understand this in order to
understand what is and is not possible to learn through science. Consider, for
example, this hypothesis: More people behave immorally when there is a full
moon than at any other time of the month. The phase of the moon is certainly a
well-defined and measurable factor, but morality is not scientifically measurable.
Thus there is no experiment that can be performed to test the hypothesis. Propose
a testable hypothesis for human behavior during a full moon.
Which of the following would be useful as scientific hypotheses? Give the
reason for each answer.
Plants absorb water through their leaves as well as through their roots.
Mice require calcium for developing strong bones.
Dogs are happy when you feed them steak.
An active volcano can be prevented from erupting by throwing a virgin into it
during each full moon.
The higher the intelligence of an animal, the more easily it can be domesticated.
The earth was created by an all-powerful being.
HIV (human immunodeficiency virus) can be transmitted by cat fleas.
The Elements of an Experiment
After completing this exercise, you should be able to
1. Define and give examples of dependent, independent, and standardized
2. Identify the variables in an experiment.
3, Explain what control treatments are and why they are used.
4. Explain what replication is and why it is important.
Once a question or hypothesis has been formed, the scientist turns his attention
to answering the question (that is, testing the hypothesis) through
experimentation. A crucial step in designing an experiment is identifying the
variables involved. Variables are things that may be expected to change during
the course of the experiment. The investigator deliberately changes the
independent variable. He measures the dependent variable to learn the effect
of changing the independent variable. To eliminate the effect of anything else
that might influence the dependent variable, the investigator tries to keep
standardized or extraneous variables constant.
Dependent Variables
The dependent variable is what the investigator measures (or counts or
records). It is what the investigator thinks will vary during the experiment. For
example, she may want to study peanut growth. One possible dependent variable
is the height of the peanut plants. Name some other aspects of peanut growth that
can be measured.
All of these aspects of peanut growth can be measured and can be used as
dependent variables in an experiment. There are different dependent variables
possible for any experiment. The investigator can choose the one she thinks is
most important, or she can choose to measure more than one dependent variable.
Independent Variables
The independent variable is what the investigator deliberately varies during the
experiment. It is chosen because the investigator thinks it will affect the
dependent variable. Name some factors that might affect the number of peanuts
produced by peanut plants.
In many cases, the investigator does not manipulate the independent variable
directly He collects data and uses the data to evaluate the hypothesis, rather than
doing a direct experiment. For example, the hypothesis that more crimes are
committed during a full moon can be tested scientifically. The number of crimes
committed is the dependent variable and can be measured from police reports.
The phase of the moon is the independent variable. The investigator cannot
deliberately change the phase of the moon, but can collect data during any phase
he chooses.
Although many hypotheses about biological phenomena cannot be tested by
direct manipulation of the independent variable, they can be evaluated
scientifically by collecting data that could prove the hypothesis false. This is an
important method in the study of evolution, where the investigator is attempting
to test hypotheses about events of the past.
The investigator can measure as many dependent variables as she thinks are
important indicators of peanut growth. By contrast, she must choose only one
independent variable to investigate in an experiment. For example, if the
scientist wants to investigate the effect that the amount of fertilizer has on peanut
growth, she will use different amounts of fertilizer on different plants; the
independent variable is the amount of fertilizer. Why is the scientist limited to
one independent variable per experiment?
Time is frequently used as the independent variable. The investigator
hypothesizes that the dependent variable will change over the course of time. For
example, she may want to study the diversity of soil bacteria found during
different months of the year, However, the units of time used may be anywhere
from seconds to years, depending upon the system being studied.
What was the independent variable in your black box investigation?
What was (or were) the dependent variable(s)?
Identify the dependent and independent variables in the following examples
(circle the dependent variable and underline the independent variable):
Height of bean plants is recorded daily for 2 weeks.
Guinea pigs are kept at different temperatures for 6 weeks. Percent weight gain
is recorded.
The diversity of algal species is calculated for a coastal area before and after an
oil spill.
Light absorption by a pigment is measured for red, blue, green, and yellow light.
Batches of seeds are soaked in salt solutions of different concentrations, and
germination is counted for each batch.
An investigator hypothesizes that the adult weight of a dog is higher when it has
fewer littermates.
Standardized Variables
A third type of variable is the standardized or extraneous variable.
Standardized variables are factors that are kept equal in all treatments, so that
any changes in the dependent variable can be attributed to the changes the
investigator made in the independent variable.
Since the investigator’s purpose is to study the effect of one particular
independent variable, she must try to eliminate the possibility that other
variables are influencing the outcome. This is accomplished by keeping the other
variables at constant levels, in other words, by standardizing these variables, For
example, if the scientist has chosen the amount of fertilizer as the independent
variable, she wants to be sure that there are no differences in the type of fertilizer
used. She would use the same formulation and same brand of fertilizer
throughout the experiment. What other variables would have to be standardized
in this experi-ment?
A hypothesis is a formal, testable statement. The investigator devises an
experiment or collects data that could prove the hypothesis false. He should also
think through the possible outcomes of the experiment and make predictions
about the effect of the independent variable on the dependent variable in each
situation. This thought process will help him interpret his results. It is useful to
think of a prediction as an if/then statement: If the hypothesis is supported, then
the results will be ….
For example, a scientist has made the following hypothesis: Increasing the
amount of fertilizer applied will increase the number of peanuts produced. He
has designed an experiment in which different amounts of fertilizer are added to
plots of land and the number of peanuts yielded per plot is measured.
What results would be predicted if the hypothesis is supported? (State how the
dependent variable will change in relation to the independent variable.)
What results would be predicted if the hypothesis is proven false?
Levels of Treatment
Once the investigator has decided what the independent variable for an
experiment should be, he must also determine how to change or vary the
independent variable. The values set for the independent variable are called the
levels of treatment, For example, an experiment measuring the effect of
fertilizer on peanut yield has five treatments. In each treatment, peanuts are
grown on a 100 m2 plot of g…

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