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The Chemical Level of Organization
(Chapter 2)
Introduction to Chemistry
All living things are composed of matter and require energy to grow and function.
Matter = anything that has mass and takes up space.
Mass = the amount of matter, i.e. the amount of “stuff” that makes up something. (This is not
the same as weight since weight varies with gravity. Your amount of matter, how much
stuff comprises you right now, remains constant regardless of where you are standing, but
you weigh more here on earth and far less on the moon. This is because the gravity of the
moon pulls less strongly on your matter, and weight is simply a measurement of the pull of
gravity on matter. You are “weightless” when there is no net pulling force acting on your
matter, and yet you clearly still have mass, i.e. you are still composed of stuff.)
Chemistry = the study of matter.
All matter will exist in one of three different states depending on the environmental conditions:
1.) solid – has a definite shape and volume.
2.) liquid – has a definite volume but takes the shape of its container.
3.) gas – has no definite shape or volume and can be compressed (both shape and
volume are dependent on the container the gas resides in).
Energy = the capacity to do work; the ability to put matter into motion.
All energy exists in one of two forms:
1.) kinetic energy = the energy of motion.
2.) potential energy = stored energy (the “potential” to put matter into motion).
(If you climb a ladder you are kinetic energy while climbing, and potential
energy while standing still at the top. If you then fall off the ladder, you are
kinetic energy while falling and potential energy again once you come to rest on
the ground.) 
Regardless of form, energy comes in four basic types depending on source:
1.) chemical energy = the energy stored in the bonds of chemical substances. (e.g.
food: when you digest a food item you are breaking the chemical bonds of the
food to release energy for movement and cell growth and repair.)
2.) electrical energy = energy of the movement of charged particles. (e.g. nerve
impulses: we will see later that communication throughout your nervous system
involves the movement of ions such as Na+ and K+ into and out of your cells.)
3.) mechanical energy = energy directly involved in moving matter. (e.g. muscles: your
skeletal muscles act to physically move your appendages around in space.)
4.) radiant energy = energy that travels in waves. (e.g. light, heat: you have
photoreceptors in your eye to detect light energy, and you constantly give off
energy to the environment in the form of heat.)
***Energy is never created or destroyed; it just changes form and type!***
Amy Warenda Czura, Ph.D.
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SCCC BIO130 Chapter 2 Handout
Composition of Matter
Atom = the smallest unit of matter.
Atoms are composed of either two or three types of subatomic particles. The three types of
subatomic particles are:
1) protons – have a positive charge and are heavy (1 atomic mass unit [amu]).
2) neutrons – have no charge and are heavy (1 amu).
3) electrons – have a negative charge but are so light they have a negligible
mass (1/2000 amu) that is often ignored.
Atomic mass number of an atom = sum of all its protons and neutrons: electrons are so small they are
not counted in the mass number. Atoms also have an atomic weight, which is a measure of the
actual mass including electrons, and is reported as an average of all known isotopes.
The mass of a human = the sum of all atoms making up the human body.
Atomic Structure
Nucleus - center of the atom, composed of protons and neutrons;
it has a positive charge
Electron orbitals / shells – organized rings or shells or orbits of electrons
moving around the nucleus at different energy levels; this region
has a negative charge.
The opposing charges of the nucleus and electrons hold atom together.
Atomic number of an atom = number of protons in an atom; this remains constant for any given atom
and defines the type of atom.
Element = all atoms of a particular type; they each have the same number of protons and are all thus
the same type of matter. There are 92 naturally occurring elements organized into the periodic
table of the elements by atomic number (number of protons in the nucleus). The table displays
the elements by their one or two letter abbreviation/symbol with the atomic number written
above the symbol and the atomic mass or atomic weight written below.
*you are not
expected to
memorize this
chart: its here
to help with the
topic!
Amy Warenda Czura, Ph.D.
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SCCC BIO130 Chapter 2 Handout
The 3 most common elements found in humans are: Hydrogen (H, Atomic #1, Atomic Mass 1),
Carbon (C, Atomic #6, Mass 12), and Oxygen (O, Atomic #8, Mass 16).
All atoms of a particular element will all have the same number of protons, but the number of
neutrons can vary between the individual atoms:
Isotope = atoms with equal protons but differing numbers of neutrons; they exhibit the same properties
(i.e. are the same element) but have a different mass number. (As a related side note, you can
calculate the number of neutrons in any atom if you know the atomic number and atomic mass
of an atom: mass = protons + neutrons, so subtracting the atomic number from the atomic mass
will give you the number of neutrons.) Back to isotopes:
e.g. Isotopes of Hydrogen
Hydrogen-1 = 1 proton, 1 electron
Hydrogen-2 = 1 proton, 1 neutron, 1 electron. (This is also known as deuterium and it
is a stable but “heavy” form of the element Hydrogen)
Hydrogen-3 = 1 proton, 2 neutrons, 1 electron. (This is also known as tritium and it is
radioactive with a half-life of 12.3 years)
Some isotopes emit subatomic particles or radiation and are called radioisotopes. There are
three types of radioisotopes based on the type of particles they emit:
1) α (“alpha”) – emit 2 protons + 2 neutrons. α emissions are weak and capable of only
superficial penetration and damage (they can be shielded by a single piece of
paper).
2) β (“beta”) – emit electrons. β emissions are weak and capable of penetrating flesh up
to 1 cm but cause only minor damage (they can be shielded by 3 mm of
aluminum or Plexiglas).
3) γ (“gamma”) – emit high energy electromagnetic waves. γ emissions are very strong
and penetrate completely through the body causing great damage due to the
ionization of biological molecules (these waves are shielded only by thick
concrete or lead).
Half-life = time required for half a given amount of a radioactive isotope to decay. (This time period
can be a matter of a few seconds to many thousands of years, depending on the radioisotope.)
Weak emitters with short half-lives can be used for diagnostic procedures (e.g. Iodine-231 has a
half-life of 8 days and is used to perform thyroid scans. In contrast, nuclear power uses
Uranium-236 which is a very strong emitter with a half-life of 100,000 years).
Radioisotopes are considered hazardous (damaging) for 10-20 half lives!
(Look back up at the U-236 half-life and ponder that number…)
Amy Warenda Czura, Ph.D.
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SCCC BIO130 Chapter 2 Handout
Electrons and Chemical Reactivity
Individual atoms start off electrically neutral: the number of protons (+) is equal to the number
of electrons (-). Thus if you know the atomic number of an atom, you know how many electrons it
has.
Electrons exist around the nucleus in ordered shells or energy
levels often referred to as orbitals. The shells are filled from the lowest
energy levels nearest to the nucleus to highest energy levels on the outside
of the atom. The outermost shell is called the valence shell. The shells
are always filled starting at the nucleus and working out, and each shell
can accommodate a certain maximum number of electrons. The first can
hold 2 electrons, the second can hold 8, the third can hold 8, and shells
beyond this can hold increasingly more. (When discussing the atoms used
to build biological molecules, their atomic numbers (and thus number of
electrons) tend to be small, so we will only need to consider the filling
number for the first three shells here.)
If the valance shell is not full to its maximum capacity, the atom will be chemically reactive.
What this means is that the atom will combine with other atoms to form molecules. This reactivity is
driven by the atom’s “desire” to have a full valence shell.
-Inert atoms have full outer shells and do not form molecules.
e.g. He = 2 electrons (e`) (2 in the first shell, this valence shell is full)
Ne = 10 e` (2 in the first shell, 8 in the second, the second/valence shell is full)
Ar = 18 e` (2 in the first, 8 in the second, 8 in the third, the third/valence shell is full)
-Atoms with a valence shell that is more than half full try to gain electrons.
e.g. Cl = 17 e` (2 in the first shell, 8 in the second, 7 in the third/valence shell. This atom
would either need to gain 1 or lose 7 electrons to have a full outer shell and it is
“easier” to gain one than lose seven.)
-Atoms with a valence shell that is more than half empty try to lose electrons.
e.g. Na = 11 e` (2 in the first shell, 8 in the second, 1 in the third/valence shell. If this atom
gives up one electron that second shell becomes the valence shell and with 8 electrons
in it, it is full.)
Ions = atoms that have an unequal number of protons and electrons. If an atom gains or loses electrons
it is no longer electrically neutral. The type of ion it is depends on its charge:
cation = more protons than electrons and thus it has a positive charge.
anion = more electrons than protons and thus it has a negative charge.
Charged atoms are denoted by indicating a + or – above the chemical symbol (e.g. Na+)
If the charge occurs because of a gain or loss of more than one electron, a number will also
accompany the symbol (e.g. Ca 2+) to indicate how many electrons were gained or lost.
Molecules or compounds can also be described as being anionic or cationic if they have an
unbalanced charge (e.g. the bicarbonate ion HCO3-)
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SCCC BIO130 Chapter 2 Handout
Chemical Bonding
Chemical bonds form between reactive atoms (those with incompletely full valence shells),
holding them together as molecules or compounds. This structure is represented as a molecular
formula by using the chemical symbols for the atoms involved, grouped together with subscript
numbers to indicate how many of each atom is included: e.g. glucose = C6H12O6 (a structure consisting
of 6 carbon atoms, 12 hydrogen atoms and 6 oxygen atoms). If we happened to have two glucose
molecules we would indicate this with a large number before the chemical symbols: e.g. 2 C6H12O6
Molecule = a chemical structure consisting of two or more atoms held together by chemical bonds.
Compound = a chemical structure consisting or two or more different atoms (elements) held together
by chemical bonds.
(e.g. Oxygen, O2, is a molecule but not a compound. Water, H2O, is both a molecule and a
compound. In biology we usually just use the word “molecule” to describe all chemically
bonded atoms regardless of how many or their type.)
Types of Chemical Bonds
1.) Ionic bonds – bonds between ions of equal but opposite charge: the opposite charges attract
holding the molecule together.
e.g. NaCl (Sodium Chloride = table salt)
Na = 11 e`, it has 1 electron in valance shell
Cl = 17 e`, it has 7 electrons in valance shell
Na donates 1 electron to Cl so both have full outer shells, but now Na
has one more proton than electrons becoming a cation, and Cl
has one more electron than protons becoming an anion.
The opposite charges attract each other forming an ionic bond.
Ionic bonds can be relatively weak in biology as they can easily separate
in water, releasing cations and anions. Since cells are composed
largely of water, ionic bonds are not the most common for
constructing important cellular molecules.
2.) Covalent bonds – bonds formed between atoms that share electrons: the electrons travel around the
outermost electron shells of both nuclei to fill both valance shells simultaneously. Electrons
are always shared in pairs and up to three pairs can be shared between atoms.
(diagrammatically, a shared pair of electrons is represented as a line between the chemical
symbols of the atoms doing the sharing):
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SCCC BIO130 Chapter 2 Handout
1.) single covalent bond – share 1 pair of electrons
e.g. H-H (hydrogen), H-O-H (water)
2.) double covalent bond – share 2 pairs of electrons
e.g. O=O (oxygen), O=C=O (carbon dioxide)
3.) triple covalent bond – share 3 pairs of electrons
e.g. N=N (nitrogen)
Usually the sharing results in full outer shells for both atoms, but
sometimes free radicals are formed.
Free radical = ion or molecule that contains unpaired electrons in its outermost shell.
e.g. N=O (N only has 7 valence electrons when O has 8)
Free radicals are highly reactive and quickly bond with something else to gain the missing
electron: this causes damage in biological systems as the electrons are “stolen” from important
molecules like DNA. (Cumulative free radical damage may be involved with the aging process! Antioxidants are molecules that can satisfy the electron needs of “angry” free radicals preventing them
from damaging your molecules).
Covalent bonds come in one of two types depending on how equally or unequally the electrons
are “shared”:
Nonpolar Covalent Bonds - equal sharing of electrons between atoms of the same type
or similar size. These form very stable bonds. (e.g. carbon-based molecules)
Polar Covalent Bonds - unequal sharing of electrons, more common between atoms of
different kinds that are vastly different in size (usually oxygen or nitrogen
bonded with hydrogen); the larger atom holds the electrons more of the time due
to greater attractive force of its nucleus (more protons). This creates a slight
positive charge on one side of the molecule and a slight negative charge on the
opposite side of the molecule. These bonds are weaker than nonpolar covalent
bonds, but are still generally stronger than ionic bonds in biological systems.
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SCCC BIO130 Chapter 2 Handout
3.) Hydrogen Bonds - attractive force between polar covalent molecules. Hydrogen bonds cannot
form molecules like ionic or covalent bonds; it is simply a relatively weak bond force between
two atoms that can hold multiple molecules together, such as between water molecules, or to
maintain the shape of a larger molecule like DNA or a protein.
Hydrogen bonds form from an attraction of
the slightly negative oxygen or nitrogen in a
polar covalent bond to the slightly positive
hydrogen from a neighboring molecule with a
polar covalent bond. (This attraction is
strong enough to create the surface tension on
water which can “float” objects such as a
sewing needle. H-bonding is also responsible
for the “beading up” effect of water.
Alcohol, which is a nonpolar substance and
thus cannot form hydrogen bonds, will
completely spread out flat when a small drop
is placed on a surface, whereas the all the Hbonds in water (4 per molecule) hold the
water molecules together in a tight sphere.)
Chemical Reactions
Chemical reaction – new chemical bonds form or existing bonds are broken between atoms. During
the reaction, both matter and energy are conserved. In biology, the input material is called the
substrate and the resulting material is called the product.
Metabolism = the sum of all chemical reactions in the cells and tissues of the body (both catabolic:
bond breaking, and anabolic: bond making).
There are three general types of chemical reactions:
1.) Decomposition reaction
e.g. AB → A + B
Bonds are broken and a molecule is converted into smaller parts.
If water is involved, it is called a hydrolysis reaction
e.g. AB + H2O → AH + HOB
catabolism = all decomposition reactions in the body.
2.) Synthesis reaction
e.g. A + B → AB
Bonds are formed to assemble a larger molecule from smaller parts.
If water is involved, it is called a dehydration synthesis or condensation reaction
e.g. AH + HOB → AB + H2O
anabolism = all synthesis reactions in the body.
3.) Exchange reaction
e.g. AB + CD → AC + BD
Rearrangement of parts of molecules; it consists of a decomposition reaction followed
by a synthesis reaction.
Most reactions are reversible and can occur in both directions at the same time in living cells.
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SCCC BIO130 Chapter 2 Handout
Activation Energy = the energy needed to start a reaction. All atoms and molecules are generally
“stable” just as they are. In order to get them to a point where it is favorable for them to react,
regardless of the type of reaction, some amount of energy must be added to cause the atoms or
molecules to leave their stable conformation and recombine into new molecules and atoms.
Most reactions occur too slowly at the temperatures that are compatible with life, so enzymes
are used in biological systems to lower the activation energy requirements and drive reactions forward
without the need for increased heat or pressure.
Enzyme = a protein catalyst, specific for one type of reaction (these are explained further with
proteins under organic molecules).
Catalyst = a compound or molecule that accelerates a reaction without being permanently
altered or consumed by the reaction.
In living systems, complex chemical reactions involving many steps are necessary to convert
substrate into a final product. An enzyme is specific for only one substrate, and thus at each step in a
pathway different enzyme will be required. (The vast majority of proteins in your body are enzymes
that carry out thousands of different chemical reactions, all of which are essential to life.)
All reactions require some energy to start (the activation energy) but some produce more
energy than they consume and some absorb more energy than they produce:
exergonic reaction –net release of energy; gives off more energy than it took (feels hot to the
touch as it gives off energy in the form of heat).
endergonic reaction – net absorption of energy; takes more energy than it produces (feels cold
to the touch as it consumes energy in the form of heat from the environment).
Types of Molecules and Compounds
Two types of molecules / compounds exist in the human body:
1.) inorganic molecules – contain no carbon and hydrogen backbone (but some do contain
carbon or hydrogen), often formed by ionic or polar covalent bonds (but not always!)
e.g. CO2, O2, H2O, NaCl
2.) organic molecules –structurally based on a carbon and hydrogen backbone, usually formed
by covalent bonds (but also defined as those molecules that are created as a result of
some living process).
e.g. C6H12O6 (glucose), CH4 (methane)
Inorganic Molecules
There are four classes of inorganic molecules found in humans (carbon dioxide, oxygen, water, and
acids, bases and salts):
1.) CO2 (carbon dioxide is waste product of aerobic cellular respiration that we exhale: we will
talk about cellular respiration in a future chapter.)
2.) O2 (oxygen is required for aerobic cellular respiration.)
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SCCC BIO130 Chapter 2 Handout
3.) H2O
Water makes up 2/3 of your total body mass and is essential to life. There are four
properties of water that make it so indispensable to living things (life would not
be possible without it!):
1.) Solubility – many molecules dissociate or dissolve in water due to its
polarity (polar covalently bonded resulting in a positive side and
a negative side of the molecule). It serves as a good solvent for
anything that is ionic or polar covalently bonded, and serves as
the basis of solutions that are necessary for life.
solution = uniform mixture of 2 or more substances
solvent = the medium
solutes = the dispersed substance
Electrolytes = ionically bonded inorganic molecules that
dissociate in water, producing ions in solution that will
conduct an electrical current. (Ionic compounds easily
dissociate in water as the cations are attracted to the
negative oxygen side and anions are attracted to the
positive hydrogen side of the molecule.) Electrolyte
solutions are essential for normal life function (e.g. nerve
impulses).
Hydrophilic molecules have polar covalent bonds; they do not
form ions but can be suspended in water to form
solutions (e.g. sugars/carbohydrates).
Hydrophobic molecules do not have polar covalent bonds and
thus do not dissolve in water (e.g. fats/lipids).
Hydrophilic and hydrophobic molecules do not associate with
each other in solutions.
2.) Reactivity – water participates in many chemical reactions. Its
polarity allows it to separate into ions and to engage in hydrolysis
and dehydration synthesis reactions.
3.) High heat capacity – water has the ability to absorb and retain
heat due to the hydrogen bonding holding individual
molecules together: it resists boiling off as a gas.
Likewise, water also retains heat and is thus rather
resistant to freezing, which means that it remains a
liquid at a large range of temperatures compatible with
life. (In a glass of water each water molecule engages
in hydrogen bonds with up to four other water
molecules keeping all the molecules connected to each
other. This produces enough force (surface tension)
that you can float a sewing needle on the surface of a
glass of water. If you add soap, this will disrupt the
hydrogen bonds and the needle will sink.)
4.) Lubrication – there is little friction between water molecules; this
allows moving parts of living things to operate smoothly.
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SCCC BIO130 Chapter 2 Handout
4.) Acids, Bases, and Salts
Water can ionize into hydroxide ions (OH-) and hydrogen ions (H+). In pure water
(H2O) they are in equal amounts. pH is the measure of the relative concentration of hydrogen
ions (expressed as the negative logarithm of hydrogen ion concentration in moles/L, but this is
not important). In pure water [H+] = 1 X 10-7 moles/L, which on the pH scale is expressed as 7.
When pH = 7, the solution is neutral, there are equal concentrations of H+ and
OH- ions.
When pH < 7, the solution is acidic, there are more H+ than OH- ions.
When pH > 7, the solution is basic or alkaline, there are more OH- than H+ ions.
Acid = any solute that dissociates in solution and releases hydrogen ions.
e.g. HCl, dissociates into H+ and ClBase = any solute that removes hydrogen ions from a solution or releases hydroxide
ions (these concepts are essentially the same since any excess H+ can combine
with free OH- to create water and cancel each other out).
e.g. NaOH, dissociates into Na+ and OHMixing an acid and a base together produces water and a salt. This reaction is
called neutralization.
e.g. HCl + NaOH → H2O + NaCl
Salt =any ionic compound consisting of any cation except hydrogen and any anion
except hydroxide. Many salts are electrolytes.
Buffer = a compound that can remove or replace hydrogen ions to stabilize pH. Usually
this involves a weak acid and its related salt. All body fluids contain buffers to
prevent rapid and large fluctuations in pH, as even small changes in the pH of
body fluids (Δ 0.4) can lead to death. (Living systems are only designed to
buffer H+ concentrations as almost all metabolism produces acid waste.)
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SCCC BIO130 Chapter 2 Handout
Organic Molecules
At a minimum, organic molecules contain carbon, hydrogen and usually oxygen, but they can
be large and complex involving other elements as well. In the human body, organic molecules are
either solids, or in one of three types of water-based “solutions”:
1.) solution = a uniform mixture
2.) colloid = a solution containing dispersed proteins or other large molecules (e.g. milk)
3.) suspension = a solution containing large particles that will settle out (e.g. blood)
There are four classes of organic molecules: carbohydrates, lipids, proteins, and nucleic acids.
Each group will be discussed below. (You should follow along with the figures in your book to
enhance your understanding of these molecules, but I am not including the figures here because I do
not want anyone to think you are responsible for their individual specific molecular structures.)
1.) Carbohydrates (make up 3% of body mass)
Carbohydrates consist of C:H:O in a 1:2:1 ratio for monosaccharides. You know them
commonly as sugars and starches, and they are generally catabolized in order to produce energy. They
are grouped into three types based on structure:
A.) Monosaccharides – simple sugars, 3-7 carbon atoms
e.g. glucose (C6H12O6): this is the most important human fuel. Fructose
is also C6H12O6 but has a different structure. When two molecules have
the same chemical formula but different shapes they are called isomers.
B.) Disaccharides – two monosaccharides bound together
e.g. sucrose (table sugar) = glucose + fructose. Disaccharides are created
by dehydration synthesis of monosaccharides and are catabolized by
hydrolysis to release energy.
C.) Polysaccharides – linkages of multiple (more than two) monosaccharides (by
dehydration synthesis) to create large insoluble molecules. e.g.
cellulose, starch, and glycogen (animal starch; the storage form of
glucose, produced by skeletal muscle and the liver).
2.) Lipids (make up 12-24% of body mass)
Lipids consist of C, H, and O but there is always much less O than C. You know them
commonly as fats, oils and waxes. Gram for gram, there is approximately 2X the energy of a
carbohydrate in a lipid of equal size. Some lipids are used for energy, some are used for structural
parts of cells, and some are used for cellular communication. Lipids are grouped into five classes
based on structure:
A.) Fatty acids – consist of a carboxyl group (COOH) bound to a long carbon and hydrogen
chain. COOH is hydrophilic and will associate with water but the larger carbon chain is
hydrophobic and will not, making the molecule insoluble. Fatty acids come in two
types base on the type of covalent bonds in the chain:
a.) saturated fatty acids – each carbon engages in four single covalent bonds.
(These kinds of fats are solid at room temperature, like butter, and are
more likely to cause solid plaques in your arteries!)
b.) unsaturated fatty acids – one or more of the carbons engages in a double
covalent bond. (These fats are liquid at room temperature, like olive oil,
and are generally considered “healthier”. Usually plant sources of fatty
acids are the unsaturated kind, but in industry they often add hydrogen
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SCCC BIO130 Chapter 2 Handout
thus making these fats saturated; this is called a trans fat or hydrogenated
fat and it will cause the same health issues that naturally saturated animal
fats do).
B.) Eicosanoids – special lipids derived from arachidonic acid (a lipid that is only produced
from dietary sources of the Omega fatty acids: eat your fish and nuts!) Eicosanoids are
used for cellular communication and are never burned for energy. There are two main
types:
a.) leukotrienes – used by cells to signal injury
b.) prostaglandins – used for cell-to-cell signaling to coordinate events (e.g.
the processes of pain and inflammation following an injury)
C.) Glycerides – consist of fatty acids bound to a glycerol group (C3H8O3) by a dehydration
synthesis (this is why you need water to “burn” fat: without adequate water your body
cannot perform hydrolysis reactions on your triglyceride (fat) deposits!) Glycerides are
named for the number of fatty acids they have:
monoglycerides = glycerol + 1 fatty acid
digylcerides = glycerol + 2 fatty acids
triglycerides = glycerol + 3 fatty acids. Triglycerides are what make up the fat
deposits on animals. These fat deposits are important for: 1) energy
storage (all excess molecules from your food are converted into
triglycerides), 2) insulation, and 3) mechanical protection (you have fat
pads in places that need cushioning like your knees and your eye
sockets.)
D.) Steroids – complex 4-ring structures with attached carbon chains, e.g. cholesterol. Steroids
are also not burned for energy but they are important to a variety of unrelated basic
functions in the body:
1. cell membrane formation and maintenance for normal cell growth, cell
division, and osmotic stability of the cell (cholesterol).
2. regulation of sexual function (steroid based sex hormones, e.g. testosterone).
3. tissue metabolism and mineral balance (steroid based metabolism hormones).
4. processing of dietary fats (structural component of bile salts).
E) Phospholipids and Glycolipids
Cell membranes are composed predominantly of phospholipids and glycolipids (but
they also include cholesterol, see point 1 above).
Phospholipids – 1 phosphate group (PO3) + 1 non-lipid group + 1 diglyceride
Glycolipid – 1 carbohydrate + 1 diglyceride
For both phospholipids and glycolipids, the head group (non-diglyceride part) is
hydrophilic and the lipid tail is hydrophobic: the large head group interacts with water whereas
the small tail group tries to shield itself from water. This results in the spontaneous formation
of a micelle when these lipids are added to water.
micelle = sphere formed in water with head groups on the outside and tails
on the inside. (Laundry detergent depends on this principle: greasy
hydrophobic dirt associates with the lipid tails of the detergent and gets
bundled into the middle of a micelle which then floats away in the water thus
removing the dirt from your clothes!)
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SCCC BIO130 Chapter 2 Handout
3.) Proteins (make up 20% of body mass)
Proteins are composed of C, O, H and N and have a variety of important functions in the body:
Functions:
1) support – structural proteins
2) movement – contractile proteins in muscle
3) transport – transport proteins in blood
4) buffering – regulate pH of all body fluids
5) metabolic regulation – enzymes to carry out metabolic chemical reactions
6) coordination and control – hormones
7) defense – skin, antibodies, clotting factors
Proteins are assembled from building blocks called amino acids. There are 20 different kinds
of amino acids. All have a central carbon with groups bound to it on its
four sides:
1) an amino group (NH2)
2) a carboxyl group (COOH)
3) a hydrogen
4) a variable R group (the nature of this structure differs for each of
the different types of amino acids. Some are as simple as a
single H, others are large complex ring structures).
Amino acids are linked together, the carboxyl group of one to amino group
of the next via dehydration synthesis, creating a special bond called a
peptide bond. A string of 3 or more amino acids bonded together is called a polypeptide. (When
skin lotions proclaim they contain “amino-peptide complexes” all they are saying is that there are some
chains of amino acids in there, essentially small proteins.) Proteins tend to consist of anywhere from
300 to 3000+ peptide-bonded amino acids linked together in a long chain. The final protein is folded
into some special shape: the correct 3-dimensional shape is absolutely essential to the function of the
protein. This folding occurs as part of the basic structure of a protein. Every protein has at least three,
but possibly four levels of structure:
4 levels of protein structure (follow along on the figure on the next page):
1) primary structure – the linear sequence of amino acids, the polypeptide, how many
amino acids were bonded together and in what order.
2) secondary structure – local folding due to hydrogen bonding between the –H and
–NH2 groups on the central carbons of neighboring amino acids. Two shapes
are possible: the α helix or the β pleated sheet (many of both can occur
simultaneously in the same protein).
3) tertiary structure – global coiling or folding due to R group interactions. (Some R
groups are charged and opposites attract, ones that contain sulfur can bind to
each other, hydrophobic ones want to get inside the protein away from water,
and hydrophilic ones want to be on the outside close to water: the chain squirms
and folds and knots itself up until all the R groups are “happy”.)
All the folding results in two basic shapes:
1) fibrous – sheets or strands, non-soluble, usually structural proteins
2) globular – compact or rounded, soluble, usually enzymes and
transport proteins
If a protein is a single polypeptide type, it is complete once its tertiary structure
is complete. If it is a multi-subunit protein, it has one more step:
4) quaternary structure – interaction between multiple tertiary folded peptides for
multi-subunit proteins.
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Once a protein is correctly folded, the result is called its native conformation. If this protein is
subjected to unfavorable environmental conditions (wrong pH, temperature, hydration, pressure, etc.),
it will unfold, and this is called denaturation. (The native conformation of egg white is a clear colloid
solution: when denatured by heat it becomes white and stiff and cannot perform “egg functions”!)
Enzymes
Enzymes are the most abundant proteins in the body. They function to catalyze metabolic
reactions. They are capable of lowering the activation energy of a reaction so that it can occur easily in
the body without the need for heat. (Without chemical reactions constantly occurring you would not
be alive. Essentially you are just a big bag of enzymes!) The correctly folded enzyme has an active
site which binds a specific substrate (or substrates) to carry out only one particular chemical reaction,
either anabolic or catabolic.
Characteristics of enzymes:
1) specificity – each enzyme binds only one kind of substrate (two or more things if it
functions in a synthesis reaction): the active site is specific and makes the
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enzyme specific.
2) have saturation limits – the concentration of substrate required to have the
maximum rate of reaction: more substrate above this point will not drive the
reaction any faster.
3) regulation – cofactors or other regulators can turn enzymes on or off. This provides
short term control over reaction rates.
cofactor = ion or molecule that binds to an enzyme to activate it; usually
a metal ion (e.g. Mg2+)
coenzyme = non-protein organic molecule that acts as a cofactor (e.g. a
vitamin)
(Most of what you take in your multivitamins (which are both vitamins
and minerals) you use as cofactors and coenzymes to allow the enzymes
produced by your cells to function correctly: you aren’t alive without
your enzymes functioning after all so you better take your vitamins!)
Temperature and pH are very important to the function of enzymes. Going outside of the
permissible range (loss of homeostasis in the body) denatures the enzyme: it unfolds and no longer
binds substrate and thus no longer catalyzes chemical reactions. No chemistry, no life. The process of
homeostasis ultimately functions to keep the human body within the parameters necessary for enzyme
function.
Conjugated Proteins
Conjugated proteins consist of a protein bound to some other organic molecule. Two common
classes are:
Glycoproteins – small carbohydrates attached to large proteins. (e.g. mucus)
Proteoglycans – large polysaccharides linked by polypeptides.
(e.g. the “glue” in connective tissue)
4.) Nucleic Acids (make up less than 1% of body mass)
Nucleic acids are composed of C, H, O, N, and P. They function to store and process
information at the molecular level. There are two types of nucleic acid molecules:
A) DNA – deoxyribonucleic acid
B) RNA – ribonucleic acid
DNA codes for RNA. RNA codes for protein (it dictates the primary structure: how many
amino acids and in what order). You are just a big bag of protein. Your DNA exists simply to provide
the master instructions for all your proteins.
Nucleic acid molecules are made of building blocks called nucleotides that are linked together
by dehydration synthesis reactions. Each nucleotide consists of three parts:
1) a pentose sugar (ribose for RNA or deoxyribose for DNA)
2) a phosphate group
3) a nitrogenous base (there are five types organized into two groups based on their core
chemical structure):
Purines (double ring structure)
- adenine (A)
- guanine (G)
Pyrimidines (single ring structure)
- cytosine (C)
- thymine (T)
- uracil (U)
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The nucleotides are assembled by bonding the sugar of one to the phosphate of the next leaving
the nitrogenous base hanging off the side. The linear strand of bonded sugar and phosphates is called
the backbone of the molecule. The order of the nucleotides when they are bonded together is the basis
of your genetic information. If you think of the nucleotides as just the letters of their bases, you can
imagine assembling these letters into words that spell out the instructions to build proteins. We will
discuss how your genetic information works in a future lecture.
RNA
RNA consists of a single chain of bonded nucleotides with the bases A, U, G, C: there is no T
base used in RNA. The chains can be hundreds to thousands of nucleotides long. There are three
different types of RNA. All are involved in the process of protein synthesis. One serves as the
instructions to assemble the amino acids into polypeptides and the other two make up parts of the
physical machinery that carries out the protein synthesis in a cell.
Three types of RNA:
1) messenger RNA (mRNA) – the protein blueprint or
instructions
2) transfer RNA (tRNA) – acts like a boat or truck to
carry amino acids to the place where proteins
are being synthesized (the ribosome).
3) ribosomal RNA (rRNA) – forms the site of protein
synthesis in the cell. This site or “factory” is
called a ribosome (we will discuss ribosomes
and cellular protein synthesis in a future
lecture).
DNA
DNA consists of a pair of very long nucleotide chains wound
around each other in a shape called a double helix. The two strands are
held in contact with each other via hydrogen bonding between
nitrogenous bases on opposing strands. Because only certain bases will
hydrogen bond with each other, this process is called base pairing. The
A base pairs with T, and the G base pairs with C (A – T and G – C): no other configurations are
allowed and there are no U bases in DNA. (We will discuss DNA in more detail in a future lecture.)
High Energy Compounds
The most important high energy compounds are assembled using RNA nucleotides and thus
are being included here with nucleic acids. High energy compounds have bonds that can be broken
easily by cells to release energy as needed (to carry out various chemical reactions). They are the
mechanism by which food energy is stored for use by cells (food molecules have a lot of energy in
their bonds but this is hard to extract quickly by individual cells). During digestion and cellular
respiration, energy from food is transferred to high energy compounds for quick and easy access. Most
high energy compounds consist of a phosphate group attached to an organic molecule in a process
called phosphorylation. Energy contained in the phosphate bond can be release by simply removing
the phosphate using water to break the bond.
Phosphorylated Adenosine (adenosine = adenine + ribose) is the most important high energy
compound for a cell. It exists in two forms depending on its degree of phosphorylation:
1) adenosine diphosphate (ADP) – 2 phosphates; low energy form
2) adenosine triphosphate (ATP) – 3 phosphates; high energy form
ADP + PO4 + energy → ATP + H2O
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When a cell has extra energy from food molecules it phosphorylates ADP to build up stock of
ATP. Later when the cell needs energy to perform some function, it breaks the phosphate bond, which
converts ATP back into ADP releasing the energy for use. (GTP and UTP also store energy for cell
functions, but are less commonly used by cells than ATP. We will discuss ATP at great length when
we cover cellular respiration in a future lecture.)
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