BIOLOGY—101
Text: Biology 7th ed. (Campbell and Reece Chapters 2, 3 and 4)
"The most incomprehensible thing about the world is that it is comprehensible."—Albert Einstein
Lecture #2: The Chemical context of Life; Water and the Fitness of the Environment
I. Atomic Structure
A. Electrons are negatively charged subatomic particles
1. Surround the nucleus of an atom
2. Kept in place by attraction to the positively charged protons found in the nucleus.
B. The periodic table.
1. Each horizontal row in the table is called a period.
2. Each column is called a group.
3. Note that the column all the way to the left contains the group 1
a. (sometimes written with a Roman numeral I) elements.
b. The column all the way to the right contains the group 18 elements (sometimes written with a
Roman numeral VIII).
4. All the elements in group one will behave similarly chemically and all the elements in group 18 or 2
or 13 etcetera will behave similarly chemically.
B. The modern model of the atom is the quantum mechanical model.
1. Even though we (and your book) will draws electrons as existing within shells that occupy space at
given distances from the nucleus, we do so only to make understanding how atoms react chemically
more understandable (see the octet rule below).
2. The fact is that electrons do NOT exist in concentric shells surrounding the nucleus. The quantum
mechanic model of the atom tells us that we cannot state exactly where a given electron is or its
directional path. We can only state with some statistical confidence where a given electron is most
likely to exist.
C.Electrons exist in four quantum energy levels.
1.The first quantum level is called n and is the same as the period number!
a. For instance hydrogen is found in the first period and thus has a quantum number n=1.
b. Helium is also in the first quantum level and thus also has a quantum number n=1.
c. The number of electrons that the primary quantum level (n) can hold is designated by the
formula 2(n2).
1.Thus the first quantum level can hold only two electrons 2(1)2=1 and n=2 elements can
hold up to eight electrons 2(2)2=8 etcetera.
2.Each principal quantum number can then be subdivided into the second quantum number called l,
which are sublevels.
a. The second quantum number refers to the angular momentum of the electrons.
b. Simply put the larger the l the farther it can be from the nucleus, i.e. the greater its (potential)
energy.
3. The values of l for any element can be any number to a maximum of l = n-1.
a. Thus if n=1, l=0. If n=2, l=0 or 1; if n=3, l=0, 1, or 2 etc.
b. Here we can see that l can have multiple sublevels starting with 0 and increasing.
c. These sublevels 0, 1, 2, 3 are themselves designated by letters.
Quantum number l
0
1
2
3
Type of sublevel
s
p
d
f
4. The number of electrons that each sublevel can hold is designated by the formula (4l+2).
a. So for instance if n=1 then l = 0 (as noted in the table above this is the s sublevel).
b. So how many electrons can the s sublevel hold? Plugging in the numbers to the above
equation: 4*0+2=2. The s orbital can hold 2 electrons.
c. Calculate the number of electrons that the p orbital can hold.
5. Each sublevel will have different orbitals—this is the third quantum number—ml. This number
has to do with the direction in space in which the electron cloud surrounding the nucleus will be
found.
a.The number of orbitals ml is directly related to the specific sublevel by the equation
(2l+1).
1. So, how many orbitals can the s sublevel have? The s sublevel l=0, using the above
equation: 2*0+1=1. The s sublevel has only one orbital, which as noted above can hold two
electrons, and its shape is said to be spherical (figure 2.9).
a. How many orbitals are found in the p sublevel?
b. Based on the number of electrons that the p sublevel can hold, how many electrons are
there per orbital in the p sublevel? Note in figure 2.9 that the shapes of the p orbitals are
dumbbell.
6. For this course we will not concern ourselves with the fourth quantum number of the electron.
D. The electrons farthest from the nucleus have the highest energy and are the most chemically
reactive. These electrons are called valence electrons
1. There is a simpler way to determine the valence electron configuration of the main group
elements (groups 1 and 2 and 13-18), than the method noted above.
2. If you remove the transition metals (groups 3-12) that leaves only eight groups.
3. Group 1 and 2 are sometimes called the s block elements and group 13-18 the p block
elements.
4. These main group elements are also designated by Roman numerals.
1. For instance, group one is also designated group IA, group 2 is IIA.
2. Ignoring the transition elements group 13 becomes group IIIA and group 14 becomes
groups IVA, 15=VA, 16=VIA, 17=VIIA and 18=VIIIA.
3. Thus the main group elements can be denoted as groups IA-VIIIA.
4. The first two groups are called the s block and groups IIIA-VIIIA are called the p block.
5. So how can we use this information to determine the electron configuration of the
valence electrons in carbon?
a. Well carbon is a group IVA element. Thus it is in the p block.
b. Importantly, the Roman Numeral IVA designates the number of valence
electrons it has i.e., four—two of these electrons would be in the 2s sublevel and two
would be in the 2p sublevel.
c. Let’s use the electron configuration methodology learned above to determine if
the shortcut is correct. Carbon is in the second period thus the first quantum number is
n=2. The number of sublevels (l) = n-1= 0 and 1. Remember that 0 is the s sublevel
and 1 is the p sublevel. Remember that s can hold 2 electrons in one orbital and p can
hold 6 electrons in three orbitals. Thus the electron configuration for carbon is
1s22s22p2.
6. Neon is also in the n=2 level. It is a group VIIIA element. Thus it has eight valence electrons.
a. As we will see, chemistry occurs so that atoms (such as carbon or oxygen), which do not
have all their orbitals of their sublevels filled, can obtain the electrons they need to fill their
orbitals.
b. They “desire” to obtain eight valence electrons. We will come back to this concept when
we review the octet rule.
E. It is worth repeating that the electrons closest to the nucleus have the lowest energy and are the
electrons that are the least reactive chemically.
1. The electrons farthest from the nucleus have the highest energy and are the most chemically
reactive. These electrons are called valence electrons.
a. The valence electrons are the electrons that determine the chemical properties of an atom.
2. Chemical reactions occur when the valence electrons of two (or more) atoms interact to form a
more stable compound—called a molecule.
3. Atoms interact chemically in such a way to fill the orbitals of their valence electrons.
F. It is also important to note that electrons can be excited (by an input of energy) to higher energy
levels.
1. As the electrons “fall” back towards the nucleus, i.e. fall back to ground state the absorbed
energy is released (figure 2.7).
2. Because electrons behave as a wave and are electromagnetic (as in light waves) the energy
released can sometimes be in the form of visible light. In fact, it is the specific wavelengths i.e. colors,
of light that electrons of specific atoms release that allowed scientists to deduce the quantum mechanic
model of the atom.
G. Most of the mass of an atom is found in its nucleus.
1. One electron has the mass of only 1/2000 a proton!
2. However, most of the size of the atom comes from the distance electrons are found from the
nucleus.
3.For instance, a hydrogen atom has one proton (designated H+) and one electron (e-). Imagine
this symbol: O is the proton of a hydrogen atom. To keep the atom to scale, the electron would be
orbiting approximately 0.5 km away! Remember though that atoms are extremely small, about 5 x
10-8 mm. It would take about 20 million hydrogen atoms placed one after another to span the
distance of this dash: (-).
II. CHEMICAL BONDS
A. A chemical reaction (which results in a chemical bond) is the process of sharing, donating, or
accepting electrons.
B. There are two types of chemical bonds
1. Ionic bond—note that an ion is an atom or molecule with an electrical charge, resulting from a
gain or loss of electrons.
a. There are two types of ions—anions are negatively charged and cations are positively
charged (figures 2.2 and 2.13).
b. Ionic bonds form between ions of opposite charge.
c. In an ionic bond, one atom donates its electrons to another atom.
1. In general, highly electronegative atoms, e.g. those elements in groups 16 and 17
(VIA and VIIA), accept electrons from weakly electronegative atoms, those elements of group
one or two.
2. Atoms that tend to donate electrons are called metals and atoms that tend to accept
electrons are called nonmetals.
d. A salt is another name for an ionic molecule.
1. Ionic molecules form salt crystals because of the attractions between the charges of the
individual atoms of the individual molecules that make up the crystal (figure 2.14).
e. We can use the concept of the main group elements to predict if the elements will tend to donate
or accept an electron.
1. For instance, sodium is a group IA element. Thus it has only one valence electron.
2. Chlorine is a group VIIA element and thus it has seven valence electrons. To fill the
valence orbitals chlorine needs to accept a single electron.
3.It can take an electron from sodium, filling its valence orbital and in the process sodium
losses its lone valence electron, which leaves eight valence electrons and a stable
configuration.
C. Covalent bond—atoms that participate in a covalent bonds share valence electrons rather than
donate and accept them as in ionic bonds (figure 2.11).
1. Covalent bonds are very strong bonds and are extremely important in biological
molecules.
2. The maximum number of covalent bonds an atom can form is usually equal to the
number of electrons needed to fill its valence orbitals.
3. For instance, if you examine figure 2.11c, you will notice that in the molecule H2O, the 1s
orbitals of hydrogen contain two electrons (the maximum they can carry) and the oxygen
atom now has 1s22s22p6—exactly the same electron configuration of neon—it has used
chemistry to obtain a stable electron configuration in all the orbitals of the various
sublevels.
4. If you go back to our discussion of electrons above you will note that oxygen needed two
electrons to fill its valence orbitals and each hydrogen atom needed one electron to fill
their respective valence orbitals.
5. Chemists draw a pair of shared electrons as a dash (figure 2.11).
D. The octet rule states that atoms seek an arrangement of electrons that will surround them with eight
electrons in their outermost (valence) energy level.
1. It generally works well for many main group elements (groups 1 and 2 and 13-18, i.e. groups
IA-VIIIA)—especially those that are biologically important such as oxygen, nitrogen and carbon.
2. How the octet rule works is most easily demonstrated by examining (figure 2.8)
3. Notice that neon and argon both have eight valence electrons.
4. These group VIIIA elements are called noble gasses because they do not normally interact
with other atoms chemically. There is no need to since their valence structure is stable.
5. If you examine figure 2.11 and 2.13 closely you will see how the octet rule does work well in
many instances. It is a close approximation of how electrons behave, but it is not perfect.
E. There are two additional types of chemical interactions that do not result from sharing or donating
electrons, but are a secondary result of those processes.
1. Van der Waals Forces—Van der Waals forces are not true chemical bonds.
a. To understand this force, you must remember that the electrons moving around the
atoms of a molecule are not evenly distributed.
b. They randomly fluctuate around the atoms of the molecule.
c. Thus, some regions surrounding the molecule will be slightly positive and others slightly
negative.
d. These fluctuations between charges occur extremely rapidly as the electrons randomly
move about the molecule.
e. The rapid fluctuations from positive to negative in one molecule or region of a molecule
can set up opposite fluctuations in electrons of nearby molecules, establishing weak
attractive forces between the atoms of the two separate molecules.
f.
Of course if the electrons of two different molecules randomly move close together there
is repulsion between the molecules.
g. The net effect of Van der Waals forces are weak attractions and repulsions and thus a
tendency for different molecules or atoms within a molecule to take up optimum
positions that represent the closest packing (due to the attractive affect of Van der Waals
forces) that does not violate the minimum space requirement of individual atoms (due to
repulsive Van der Waals force).
h. Van der Waals forces do not come into play until distances between atoms of different
molecule are less than 1 angstrom. One angstrom is 10-10 meters!
i. Page 42 in your book shows a picture of a gecko.
1. These lizards are well known for their ability to walk up walls and hang upside
down—even on extremely smooth surfaces such as glass.
2. The feet of geckos have thousands of tiny hair-like projections which create a
very large surface area.
3. Each hair will have a week attraction to the surface it is on by Van der Waals
forces. Since there are hundreds of thousands of individual hairs the overall
attractive force is enough to allow the gecko to climb walls!
.
2. Before we introduce the last chemical attractive force, we need to introduce the concept of
polar versus non-polar covalent bonds:
a. If the electrons shared by a given molecule spend equal time near each nucleus, not only is the
molecule as a whole electrically neutral, but each end or pole of the molecule is electrically
neutral.
b.
Such an electrically symmetrical bond is called a nonpolar covalent bond and the compound
formed with such nonpolar bonds is a nonpolar molecule.
1.Hydrocarbons, which are long chains of carbon atoms covalently bound to
hydrogen atoms, are examples of nonpolar molecules.
c. But electron-sharing in covalent bonds is not always equal.
d.
In some molecules, one nucleus may initially have a larger positive charge and, therefore,
attract the electrons more strongly than another nucleus in the molecule.
e. This results in a polar covalent bond. Even though the molecule as a whole has no charge—it
is electrically neutral—it does have weakly-charged ends or poles (figure 2.12).
f.
This occurs because one atom in the molecule will more strongly attract the electrons than
another atom.
g. The most electronegative atom is fluorine. But oxygen, nitrogen and chlorine are also extremely
electronegative.
h. Atoms in group 1, which includes hydrogen, are not strongly electronegative.
i.
A strongly electronegative atom will preferentially hold electrons, and a weakly
electronegative atom will easily loose electrons.
j. The overall result of a polar covalent bond is that one atom will be weakly negative, (written δ-)
due to its attraction to electrons and the other weakly positive (written δ+) do to the loss of electrons.
k. Note that we are not dealing with full + or – charges. These are covalent bonds and so we are
still sharing electrons but the electrons are not shared equally.
3. This leads us to the last type of chemical interaction, hydrogen bonds.
Hydrogen bond—an attractive force formed between molecules when a hydrogen atom is covalently
bound to either O, N, or F.
a. This bond between H—O, H—N, or H—F would be by definition a polar covalent bond.
b.
The intramolecular polar covalent bond results in the intermolecular attraction between
individual molecules that we call the hydrogen bond.
c.
It should be repeated that the hydrogen bond is not a true bond but a chemical
attraction between individual molecules as a result of polar covalent bonds (figures 2.15
and 3.2)
d. The classic example of a polar molecule and the resulting hydrogen bonding between
individual molecules is water—H2O.
III. WATER (H2O)
A. As noted above water is a polar molecule.
B. The oxygen atom has eight protons, but the hydrogen atoms have only one proton. Thus,
oxygen is more electronegative and will more strongly attract the hydrogen atom’s electron.
C. Thus the oxygen atom will have a weakly negative charge. The hydrogen atoms become
slightly positive as they lose their grip on their electrons (figure 2.12).
D. The polar nature of water results in hydrogen bonding—the weak attraction between a
hydrogen atom that bears a partial positive charge from one water molecule and an oxygen
atom that bears a partial negative charge from a separate water molecule.
E. Note: hydrogen bonding can also occur between hydrogen and nitrogen, this will be important
when we discuss DNA (figures 2.15 and 3.2).
F. The polar nature of water and the resulting H-bonding gives water many unique and very
important properties, vital for life.
1. Water is a good solvent. It is capable of dissolving a wide range of substances. Note:
water or other liquid substances containing dissolved substances are called solutions.
2. Solute—the substance that is dissolved in a solvent.
3. Aqueous solution—a solution where water is the solvent.
F. Water easily dissolves other polar or charged molecules. Remember the rule “like dissolves
like” and the saying “oil and water don’t mix.”
1. Na+Cl- (table salt) dissolves in water.
2. The weakly positive hydrogen atoms surround the Cl- anion and the weakly negative oxygen
atoms surround the Na+ cation, enclosing the sodium and chloride ions and separating the
crystal so that it dissolves in the liquid (figure 3.6).
3. Note: table salt is thus said to be a hydrophilic molecule—which literally means ‘waterloving’ molecule.
G. However oil, which is a hydrocarbon, is a nonpolar molecule.
1. The electrons of the C—C covalent bonds and C—H covalent bonds are shared fairly equally
and so there is no resultant week positive or negative charges.
2. Instead of surrounding the atoms of an oil molecule, the weakly charged atoms of water
exclude oil, creating layers (for instance, an oil sheen on the surface of a lake).
3. Oil and other nonpolar molecules are said to be hydrophobic—which literally means ‘water
fearing’.
H. Cohesion—water molecules have a tendency to stick together. That is because the oxygen atom of
one water molecule is weakly attracted to the hydrogen atoms of another molecule. This interaction
creates a lattice structure (figure3.2).
1.Cohesion is very important in biological processes, e.g. transpiration. Water is released
(evaporates) from the leaves of trees—a process called transpiration. As water is lost from a
tree’s leaf, it is replaced by water moving up in the trunk from the roots through tubes in the
tree’s tissue. The hydrogen bonds of water are STRONGER than the weight of the water, even
pulling the water up very tall trees (figure 3.3).
I. Surface tension—cohesion at the surface of the water creates surface tension. Surface tension is
what allows insects to walk on water. The water molecules are sticking together forming a weak
electromagnetic force (figure 3.4).
a. To understand this even further, remember that what keeps your hand from going through a
brick wall is the electromagnetic forces—the covalent bonds—of the brick repelling the
electromagnetic forces in the covalent bonds of your hand. Obviously, electromagnetism is
strong, even stronger than gravity. “It isn’t the fall that will kill you, it is the sudden stop” i.e.
the convergence of strong electromagnetic forces between your body and the ground.
b. So the week electromagnetic force of hydrogen bonds creates a cohesive force at the
surface of a body of water called surface tension. This electromagnetic force is strong
enough for insects and other organisms to walk on water. It is also why a belly flop hurts!
J. Adhesion—a term that describes water’s tendency to adhere to polar (charged) surfaces. Put water
into a narrow test tube and you will see water move up the sides of the glass (glass is polar).
1. Thus water tends to stick to glass by adhesion and the water sticking to the sides of the
glass pulls up the water molecules below by cohesion.
2. This results in what is called a meniscus—we will discuss the concept of a meniscus further
in lab.
K. Water’s hydrogen bonds help moderate temperature changes.
1. First you need to understand that temperature is the average kinetic energy of the molecules of
a substance (i.e., how fast on average are the molecules of a substance moving). The faster they
move the hotter the substance and vice versa.
2. Note: This is different than the heat of a substance, which is the total amount of kinetic energy
of a substance.
3. Hydrogen bonds resist molecular movement in water. This should make sense. The hydrogen
bonds create an attractive cohesive force between water molecules that must be overcome
before the water molecule’s kinetic energy can increase, i.e. the hydrogen bonds must first be
broken before water can increase in heat and thus temperature. Energy absorbed by water
must first break the H-bonds before the molecules will begin to speed up and subsequently
increase the temperature. H-bonds are in effect storing energy.
a. As water cools, it “stores” energy in the form of H-bonds.
L. The effect of hydrogen bonds can be measured and quantified relative to other liquids:
1. Specific heat—the amount of heat that must be absorbed or lost for 1 g of that substance to change
its temperature by 1oC. Water has a very high specific heat. It takes 1 cal/g/C. In other words, it takes
one calorie to raise one gram of water one degree celsius. The specific heat of alcohol is just over half
that (0.6 cal/g/C).
a. This has a number of effects. For instance, water resists temperature changes. So if you live
near a large body of water, as the air temperature cools in the fall, heat is released from the
water to the atmosphere. The kinetic energy of the water is thus lost to the air and more and
more hydrogen bonds form between individual water molecules. Conversely, as the air heats
up in spring and summer, water absorbs this heat. The water itself is slow to heat because first
enough energy must be absorbed to break the hydrogen bonds and then the kinetic energy of
the water can increase.
b. The high specific heat of water moderates coastal air temperature. Coastal communities
generally have a much more moderate climate than inland communities. Coastal communities
are warmer than inland areas in the fall and winter and milder than inland areas in the spring
and summer. This is the reason that right along the coast in South Carolina we have palmetto
palm trees. If you move inland even slightly (sometimes only a few hundred feet) you will not
see palms growing naturally. All the palmettos you see growing in our state that are not right
along the coast have been planted and do not reproduce—it is too cold and temperature
extremes too great.
2. Vaporization (evaporation)—process in which molecules move fast enough (have enough kinetic
energy) to enter into a gaseous state.
a. Heat of vaporization is the quantity of heat a liquid must absorb for 1 g of it to be converted
from the liquid to the gaseous state.
b. Water has a high heat of vaporization. To evaporate 1 g of water, about 580 calories of heat are
needed—almost double that needed to evaporate 1 g of alcohol or ammonia.
3. Evaporative cooling—as a substance evaporates, the surface of the liquid that remains behind
cools down. This is because the “hottest” molecules—that is, those with the greatest kinetic energy are
the most likely to evaporate, leaving behind “cooler” molecules—those with the least kinetic energy.
That’s why we sweat, dogs pant and plants transpire! It is also why you should not muzzel your dog on
a hot day. They can’t sweat—they do not have any sweat glands on their skin and so evaporative
cooling occurs through their tongue. If you muzzle your dog on a hot day it can easily overheat.
4. Ice is less dense than liquid water. Water is one of the few substances that is less dense in the solid
form than in the liquid form (figure 3.5).
a. At above 4oC, water behaves like other liquids, expanding as it warms and contracting as it
cools (due to the kinetic energy of the molecules).
b. Water begins to freeze when its molecules are no longer moving vigorously enough to break
their hydrogen bonds. Thus, at 0oC, the water molecules are locked into a crystalline structure
at the exact distance allowed by hydrogen bonding, with each water molecule hydrogen bonded
to the maximum of four partners. There is not enough kinetic energy for the water molecules to
break free of their hydrogen bonds. Water molecules can’t move closer to one another because
of the hydrogen bonding—these bonds lock them in place at a fixed distance to one another.
c. Thus, ice is 10% less dense than liquid water—i.e., 10% fewer molecules for the same volume.
Of course, this fact is important to fish living in temperate ponds and lakes.
M. pH
1. pH is the concentration of hydrogen ions (protons—written H+) in moles per liter—for a
discussion on moles study the appropriate section in lab one in your lab book.
2. The concentration of H+ in pure water is about 1/10,000,000 of a mole per liter. One mole of
H+ weighs 1 gram so in pure water there is 1/10,000,000 of a gram of H+. Obviously, water
splits into H+ and OH- (hydroxide ions) very rarely, as noted, only 1 out of 10,000,000 water
molecules is apart at any instant.
3. Note that 1/10,000,000 can be rewritten as 1/107. Of course 1/107 can be rewritten as 10-7.
To simplify matters, pH is actually expressed as the negative log (base 10) of the hydrogen ion
concentration: pH = -log [H+]. Thus the pH of pure water is –log 10-7, which equals – (-7), which
equals 4. Simply put, pH is expressed as the number of zeros in the denominator (7 in pure
water because 10,000,000 has 7 zeros). As the H+ concentration increases 10 times, the pH
goes down one (1/1,000,000 at pH of 6, and 1/100,000 at pH of 5).
a. Thus, each pH unit represents a 10x change in the hydrogen ion concentration. Thus a pH
of 2 has 10 times more protons than pH 3 and is 10 times more acidic. pH 12 has 10 times
fewer protons than pH 11 but has 10 times more hydroxide ions (OH-).
b. At ph 7 a substance is said to be neutral with equal numbers of:
H+ + OH-  H2O; below 7 is acid; above 7 is alkaline (basic).
4. See the pH scale (figure 3.8).
N. Buffers are substances that resist pH changes by accepting H+ when they are in excess and by
donating H+ when they are scarce. A buffer system consists of a weak acid and a weak base. One
such buffer system consists of the weak acid carbonic acid (H2CO3) and the weak base sodium
bicarbonate (NaHCO3): the H2CO3/ NaHCO3 system.
1. If a strong acid such as hydrochloric acid HCl is added to an aqueous solution the weak base
NaHCO3 can take up excess protons in the following chemical reaction:
HCl + NaHCO3  H2CO3 + NaCl
2. If a strong base such as sodium hydroxide NaOH (found in oven cleaner) is added to an
aqueous solution the weak acid can release protons in the following chemical reaction:
NaOH + H2CO3  NaHCO3 + H2O
3. Blood stays at a relative constant pH of very close to 7.4 because of buffers. This is very
important because slight changes in pH can have very deleterious effects on biological
molecules.
4. Buffer systems such as the one noted above keep pH at a constant range by accepting or
donating H+ in solution.
IV. CHEMICAL REACTIONS
A. In Biochemistry, Structure = Function—believe it or not, a very small change in the structure of a
biological molecule can completely change its function.
B. Molecules are rearranged through chemical reactions
1.Chemical reactions are simply the breaking and re-forming of chemical bonds, leading to changes
in chemical matter. The acid base reactions noted above are excellent examples of the
rearrangement of chemical bonds resulting in different products from the original reactants.
2. General form of a chemical reaction: Reactants  Products
3. In a chemical reaction, you do not change the number of each atom present from the reactants
to the products; you simply change the position or arrangement of the atoms
a. For instance, the creation of sugar and oxygen from carbon dioxide and wate during
photosynthesis: 6CO2 + 6H20 + sunlight  C6H12O6 + 6O2
V. BIOLOGICAL CHEMISTRY: CARBON BASED
A. Carbon based molecules are called organic molecules (or organic compounds). That is because
carbon based molecules constitute the central chemicals of all living things on this planet.
B. Carbon atoms are so important to biological chemistry because of their valence electrons. A
carbon atom has four electrons in its valence shell (remember the octet rule states that an atom
can hold eight electrons). Therefore it can bond up to four other atoms through single covalent
bonds (figures 4.3 and 4.4).
C. In other words carbon atoms can bind to each other to form long chains, and still have enough
valence electrons to bind to other atoms where additional chemistry can occur (figure 4.5). To
understand what is meant by this, you need to understand the concept of functional groups and
how they facilitate chemical reactions between organic compounds creating even larger
molecules. We will examine functional groups shortly.
D. Stuctural Isomers—compounds with the same molecular formula, but a different covalent
chemical structure e.g., C—C—OH ethanol and C—O—C dimethyl ether (figure 4.7a).
E. Geometric isomers have the same chemical covalent structures but different spatial
arrangements of their atoms. With geometric isomers we have either a cis or trans arrangement of
atoms surrounding two carbon molecules, themselves held together by a double covalent bond
(figure 4.7b). We will see geometric isomers again when we discuss fats. Ever hear a commercial
touting a food product as having no trans fatty acids?
1. Remember a slight change in the arrangement of atoms in a molecule can greatly change its
function.
VI. BIOLOGICAL CHEMISTRY: FUNCTIONAL GROUPS
Along with the carbon skeleton, organic molecules have attached molecules called functional groups
that greatly determine the properties of an organic compound.
A. Functional groups usually participate in chemical reactions.
B.
A hydrocarbon can have more than one functional group.
C.
Functional groups are far less stable than the carbon backbone and are more likely to
participate in chemical reactions, i.e. the functional groups are where the interesting
chemistry such as dehydration synthesis and hydrolysis occurs. As we will see in the next
lecture, dehydration synthesis and hydrolysis allows small molecules to combine and form
larger molecules and then to be broken back down to smaller constituents once again.
D.
The functional groups you should know include:

Hydroxyl—results in molecules called alchohols.

Carboxyl—results in molecules called carboxylic acids.

Carbonyl—results in molecules called ketones and aldehydes.

Sulfhydryl—results in molecules called thiols.

Phosphate—results in molecules called organic phosphates.

Amino—results in molecules called amines.

Methyl – nonpolar

Hydryl- nonpolar
Examples of these functional groups can be found in your book (figure 4.10).
YOU MUST BE ABLE TO LIST AND DRAW THESE FUNCTIONAL GROUPS!