Chapter 2
The Chemistry of Microbiology
I. Elements:
 Substances that can not be broken down into
simpler substances by chemical reactions.
 There are 92 naturally occurring elements: Oxygen,
carbon, nitrogen, calcium, sodium, etc.
 Life requires about 25 of the 92 elements
 Chemical Symbols:
 Abbreviations for the name of each element.
 Usually one or two letters of the English or Latin
name of the element
 First letter upper case, second letter lower case.
Example: Helium (He), sodium (Na), potassium
(K), gold (Au).
Main Elements: Over 98% of an organism’s mass is
made up of six elements.
 Oxygen (O): 65% body mass
 Cellular respiration, component of water, and most
organic compounds.
 Carbon (C): 18% of body mass.
 Backbone of all organic compounds.
 Hydrogen (H): 10% of body mass.
 Component of water and most organic compounds.
 Nitrogen (N): 3% of body mass.
 Component of proteins and nucleic acids (DNA/RNA)
 Calcium (Ca): 1.5% of body mass.
 Bones, teeth, clotting, muscle and nerve function.
 Phosphorus (P): 1% of body mass
 Bones, nucleic acids, energy transfer (ATP), phospholipids.
Minor Elements: Found in low amounts. Between
1% and 0.01%.
 Potassium (K): Main positive ion inside cells.

Nerve and muscle function.
 Sulfur (S): Component of most proteins.
 Sodium (Na): Main positive ion outside cells.

Fluid balance, nerve function.
 Chlorine (Cl): Main negative ion outside cells.

Fluid balance.
 Magnesium (Mg): Component of many enzymes
and chlorophyll.
Trace elements: Less than 0.01% of mass:














Boron (B)
Chromium (Cr)
Cobalt (Co)
Copper (Cu)
Iron (Fe)
Fluorine (F)
Iodine (I)
Manganese (Mn)
Molybdenum (Mo)
Selenium (Se)
Silicon (Si)
Tin (Sn)
Vanadium (V)
Zinc (Zn)
II. Structure & Properties of Atoms
Atoms: Smallest particle of an element that retains
its chemical properties. Made up of three main
subatomic particles.
Particle
Location
Mass
Proton (p+) In nucleus
1
Neutron (no) In nucleus
1
Electron (e-) Outside nucleus 0*
* Mass is negligible for our purposes.
Charge
+1
0
-1
Atomic Particles: Protons, Neutrons, and Electrons
Helium Atom
Carbon Atom
Structure and Properties of Atoms
1. Atomic number = # protons
 The number of protons is unique for each element
 Each element has a fixed number of protons in its
nucleus. This number will never change for a given
element.
 Written as a subscript to left of element symbol.
Examples: 6C, 8O, 16S, 20Ca
 Because atoms are electrically neutral (no charge),
the number of electrons and protons are always the
same.
 In the periodic table elements are organized by
increasing atomic number.
Structure and Properties of Atoms:
2. Mass number = # protons + # neutrons
 Gives the mass of a specific atom.
 Written as a superscript to the left of the element
symbol.
Examples: 12C, 16O, 32S, 40Ca.
 The number of protons for an element is always the
same, but the number of neutrons may vary.
 The number of neutrons can be determined by:
# neutrons = Mass number - Atomic number
Structure and Properties of Atoms:
3. Isotopes:
Variant forms of the same element.
 Isotopes have different numbers of neutrons and
therefore different masses.
 Isotopes have the same numbers of protons and
electrons.
 Example: In nature there are three forms or isotopes
of carbon (6C):
 12C:


About 99% of atoms. Have 6 p+, 6 no, and 6 e-.
13C: About 1% of atoms. Have 6 p+, 7 no, and 6 e-.
14C: Found in tiny quantities. Have 6 p+, 8 no, and 6 e-.
Radioactive form (unstable). Used for dating
fossils.
Electrons Determine How Atoms Bond with Other Atoms
A. Energy levels: Electrons occupy different energy levels
around the nucleus.
Level (Shell)
1
2
3
4, 5, & 6
Electron Capacity
2 (Closest to nucleus, lowest energy)
8
8 (If valence shell, 18 otherwise)
18
B. Electron configuration: Arrangement of electrons in
orbitals around nucleus of atom.
C. Valence Electrons: Number of electrons in outer energy
shell of an atom.
III. How Atoms Form Molecules:
Chemical Bonds
Molecule: Two or more atoms combined chemically.
Compound: A substance with two or more elements
combined in a fixed ratio.





Water (H2O)
Hydrogen peroxide (H2O2)
Carbon dioxide (CO2)
Carbon monoxide (CO)
Table salt (NaCl)
 Atoms are linked by chemical bonds.
Chemical Formula: Describes the chemical composition
of a molecule of a compound.
 Symbols indicate the type of atoms
 Subscripts indicate the number of atoms
How Atoms Form Molecules: Chemical Bonds
“Octet Rule”: When the outer shell of an atom is not full, i.e.:
contains fewer than 8 (or 2) electrons (valence e-), the atom
tends to gain, lose, or share electrons to achieve a complete
outer shell (8, 2, or 0) electrons.
Example:
Sodium has 11 electrons, 1 valence electron.
Sodium loses its electron, becoming an ion:
Na
-------> Na+
+ 1 e1(2), 2(8), 3(1)
1(2), 2(8)
Outer shell has 1 eOuter shell is full
Sodium atom
Sodium ion
Number of Valence Electrons Determine the
Chemical Behavior of Atoms
Element
Sodium
Calcium
Aluminum
Carbon
Nitrogen
Oxygen
Chlorine
Neon*
* Noble gas
Valence
Electrons
1
2
3
4
5
6
7
8
Combining
Capacity
1
2
3
4
3
2
1
0
Tendency
Lose 1
Lose 2
Lose 3
Share 4
Gain 3
Gain 2
Gain 1
Stable
Electron Arrangements of Important
Elements of Life
1 Valence electron
4 Valence electrons
5 Valence electrons
6 Valence electrons
How Atoms Form Molecules:
Chemical Bonds
Atoms can lose, gain, or share electrons to satisfy
octet rule (fill outermost shell).
Two main types of Chemical Bonds
A. Ionic bond: Atoms gain or lose
electrons
B. Covalent bond: Atoms share electrons
A. Ionic Bond: Atoms gain or lose electrons. Bonds
are attractions between ions of opposite charge.
Ionic compound: One consisting of ionic bonds.
Na + Cl ----------> Na+ Clsodium chlorine
Table salt
(Sodium chloride)
Two Types of Ions:
Anions: Negatively charged particle (Cl-)
Cations: Positively charged particle (Na+)
B. Covalent Bond - Involve the “sharing” of one or more
pairs of electrons between atoms.
Covalent compound: One consisting of covalent
bonds.
Example: Methane (CH4): Main component of
natural gas.
H
|
H---C---H
|
H
Each line represents on shared pair of electrons.
Octet rule is satisfied: Carbon has 8 electrons,
Hydrogen has 2 electrons
There May Be More Than One Covalent Bond
Between Atoms:
1. Single bond: One electron pair is shared
between two atoms.
Example: Chlorine (Cl2), water (H2O); methane
(CH4)
Cl
Cl
2. Double bond: Two electron pairs share
between atoms.
Example: Oxygen gas (O2); carbon dioxide (CO2)
O=O
3. Triple bond: Three electron pairs shared
between two atoms.
Example: Nitrogen gas (N2)
N=N
Number of Covalent Bonds:
Carbon (4)
Nitrogen (3)
Oxygen (2)
Sulfur (2)
Hydrogen (1)
Two Types of Covalent Bonds: Polar and Nonpolar
A. Electronegativity: A measure of
an
atom’s ability to attract and hold onto
a shared pair of electrons.
Some atoms such as oxygen or
nitrogen have a much higher
electronegativity than others, such as
carbon and hydrogen.
Element
O
N
S&C
P&H
Electronegativity
3.5
3.0
2.5
2.1
Polar and Nonpolar Covalent Bonds
B. Nonpolar Covalent Bond: When the atoms in a
bond have equal or similar attraction for the
electrons (electronegativity), they are shared
equally.
Example: O2, H2, N2, Cl2
C. Polar Covalent Bond: When the atoms in a bond
have different electronegativities, the electrons
are shared unequally. Electrons are closer to the
more electronegative atom creating a polarity or
partial charge.
Example: H2O
 Oxygen has a partial negative charge.
 Hydrogens have partial positive charges.
Other Bonds: Weak chemical bonds are important
in the chemistry of living things.
 Hydrogen bonds: Attraction between the partially
positive H of one molecule and a partially
negative atom of another
Hydrogen bonds are about 20 X easier to break than
a normal covalent bond.
 Responsible for many properties of water.
 Determine 3 dimensional shape of DNA and
proteins.
 Chemical signaling (molecule to receptor).

Water - A Unique Compound for
Life
Water: The Ideal Compound for Life
 Living cells are 70-90% water
 Water covers 3/4 of earth’s surface
 Water is the ideal solvent for chemical
reactions
 On earth, water exists as gas, liquid,
and solid
I. Polarity of water causes hydrogen bonding
 Water molecules are held together by H-bonding
 Partially positive H attracted to partially negative O
atom.
 Individual H bonds are weak, but the cumulative
effect of many H bonds is very strong.
Unique properties of water caused by H-bonds
 Cohesion: Water molecules stick to each other.
 Adhesion: Water molecules stick to many surfaces.
 Stable Temperature: Water resists changes in
temperature.
 High heat of vaporization: Water must absorb large
amounts of energy (heat) to evaporate.
 Expands when it freezes (water denser than ice)
 Solvent: Dissolves many substances.
II. Biological Consequences of Water’s Polarity
A. Capillary Action: Water tends to rise in narrow
tubes. This is caused by two factors:


Cohesion: Molecules of water “stick together”
Adhesion: Water molecules stick to walls of tubes.
Examples: Upward movement of water through plant vessels
and fluid in blood vessels.
B. Surface tension: Difficulty in “stretching or
breaking”

At water/air interface, difficult to pull water apart

Causes water to “bead” into tiny balls

Used by some insects who live on the surface of water
C. Temperature Regulation
Water has a very high specific heat

Specific Heat: Amount of heat energy needed to raise 1 g of
substance 1 degree Celsius

Specific Heat of Water: 1 calorie/gram/degree C

Organisms can absorb a lot of heat without drastic changes
in temperature.
D. Evaporative Cooling

Vaporization: Transformation from liquid to gas.

Heat of Vaporization: Energy required to convert 1 gram of
a liquid -> gas is high (540 calories/gram)


Sweating is a form of evaporative cooling.
Can regulate temperature w/o great water loss.
E. Ice floats on Water: Life Can Exist in Bodies of
Water
Ice floats because liquid water is more dense than ice
(solid water).

Water gets more dense as it cools to 4oC.

Water gets less dense (expands) as it cools further to form
ice.

Crystalline lattice forms, molecules farther apart
Because ice floats, life can survive and thrive in
bodies of water, even though the earth has gone
through many winters and ice ages
III. Water is the ideal solvent for chemical reactions
 Solution: Homogeneous mixture of 2 or more
substances.

Examples: Salt water, air, tap water.
 Solvent: Dissolving substance of a solution.

Example: Water, alcohol, oil.
 Solute: Substance dissolved in the solvent.

Example: NaCl, sugar, carbon dioxide.
 Aqueous solution: Water is the solvent.
 Solubility: Ability of substance to dissolve in a given
solvent.
Solubility of a Solute Depends on its
Chemical Nature
Two Types of Solutes:
A. Hydrophilic: “Water loving” dissolve easily in
water.




Ionic compounds (e.g. salts)
Polar compounds (molecules with polar regions)
Examples: Compounds with -OH groups (alcohols).
“Like dissolves in like”
B. Hydrophobic: “Water fearing” do not dissolve in
water


Non-polar compounds (lack polar regions)
Examples: Hydrocarbons with only C-H non-polar
bonds, oils, gasoline, waxes, fats, etc.
ACIDS, BASES, pH AND BUFFERS
A. Acid: A substance that donates protons (H+).
 Separate into one or more protons and an anion:
HCl (into H2O ) -------> H+ + ClH2SO4 (into H2O ) --------> H+ + HSO4 Acids INCREASE the relative [H+] of a solution.
 Water can also dissociate into ions, at low levels:
H2O <======> H+ + OH-
B. Base: A substance that accepts protons (H+).
 Many bases separate into one or more positive ions
(cations) and a hydroxyl group (OH- ).
 Bases DECREASE the relative [H+] of a solution (
and increases the relative [OH-] )
H2O <======> H+ + OHDirectly
NH3 + H+ <=------> NH4+
Indirectly NaOH ---------> Na+ + OH( H+ + OH- <=====> H2O )
Strong acids and bases: Dissociation is almost
complete (99% or more of molecules).
HCl (aq) -------------> H+ + ClNaOH (aq) -----------> Na+ + OH(L.T. 1% in this form)
(G.T. 99% in dissociated form)
 A relatively small amount of a strong acid or base will
drastically affect the pH of solution.
Weak acids and bases: A small percentage of
molecules dissociate at a give time (1% or less)
H2CO3 <=====> H+ +
HCO3carbonic acid
Bicarbonate ion
(G.T. 99% in this form)
(L.T. 1% in dissociated form)
C. pH scale: [H+] and [OH-]
 pH scale is used to measure how basic or acidic a
solution is.
 Range of pH scale: 0 through 14.

Neutral solution: pH is 7. [H+ ] = [OH-]

Acidic solution: pH is less than 7. [H+ ] > [OH-]

Basic solution: pH is greater than 7. [H+ ] < [OH-]
 As [H+] increases pH decreases (inversely proportional).
 Logarithmic scale: Each unit on the pH scale represents
a ten-fold change in [H+].
pH of Common Solutions
D. Buffers keep pH of solutions relatively constant
 Buffer: Substance which prevents sudden large
changes in pH when acids or bases are added.
 Buffers are biologically important because most of
the chemical reactions required for life can only
take place within narrow pH ranges.
 Example:

Normal blood pH 7.35-7.45. Serious health problems will
arise if blood pH is not stable.
CHEMICAL REACTIONS
 A chemical change in which substances (reactants)
are joined, broken down, or rearranged to form new
substances (products).
 Involve the making and/or breaking of chemical
bonds.
 Chemical equations are used to represent chemical
reactions.
Example:
2H2 +
O2 -----------> 2H2O
2 Hydrogen
Molecules
Oxygen
Molecule
2 Water
Molecules
Organic Compounds
I. Organic Chemistry: Carbon Based Compounds
 Organic Compounds: Compounds that contain
carbon and are synthesized by cells (except CO
and CO2).
 Diverse group: Several million organic compounds are
known. More are identified daily.
 Common: After water, organic compounds are the most
common substances in cells.


Over 98% of the dry weight of living cells is made up of organic
compounds.
Less than 2% of the dry weight of living cells is made up of
inorganic compounds.
 Inorganic Compounds: Compounds without
carbon.
Organic Compounds are Carbon Based
Carbon Has 4 Valence Electrons and
Can Form 4 Covalent Bonds
Organic Compounds are Incredibly Diverse
Organic molecules can vary dramatically in:



Length (1-100s of C atoms)
Shape (Linear chain, branched, ring)
Type of bonds:




Single
Double
Triple bonds
Other elements that bond to C:





Nitrogen (N)
Oxygen (O)
Hydrogen (H)
Sulfur (S)
Phosphorus (P)
Carbon Skeletons of Organic Compounds
Diversity of Organic Compounds
 Hydrocarbons:
 Organic molecules that contain C and H only.
 Good fuels, but not biologically important.
 Undergo combustion (burn in presence of oxygen).
 In general they are chemically stable.
 Nonpolar: Do not dissolve in water (Hydrophobic).
Examples:








(1C) Methane:
(2C) Ethane:
(3C) Propane:
(4C) Butane:
(5C) Pentane:
(6C) Hexane:
(7C) Heptane:
(8C) Octane:
CH4
CH3CH3
CH3CH2CH3
CH3CH2CH2CH3
CH3CH2CH2CH2CH3
CH3CH2CH2CH2CH2CH3
CH3CH2CH2CH2CH2CH2CH3
CH3CH2CH2CH2CH2CH2CH2CH3
Hydrocarbons have C and H only
Isomers: Compounds with same chemical formula
but different structures
 Structural Isomers: Differ in atom arrangement:
Example: Isomers of C4H10
Butane (C4H10)
CH3--CH2--CH2--CH3
Isobutane (C4H10)
CH3--CH--CH3
|
CH3
 Isomers have different physical and chemical
properties.
II. Functional Groups Determine Chemical & Physical
Properties of Organic Molecules
 Compounds that are made up solely of carbon and
hydrogen (hydrocarbons) are not very reactive.
 In an organic compound, the groups of atoms that usually
participate in chemical reactions are called functional
groups.

Groups of atoms that have unique chemical and physical
properties.

Biologically important functional groups:

Hydroxyl (-OH)

Carbonyl (=C=O)

Carboxyl (-COOH)

Amino (-NH2)
Notice that all are polar.
A. Hydroxyl Group (-OH)
 Polar group: Polar covalent bond between O and H.
 Can form hydrogen bonds with other polar groups.
 Generally makes molecule water soluble.
Found in:
 Alcohols: Organic molecules with a simple hydroxyl
group. Examples:

Methanol (wood alcohol, toxic)

Ethanol (drinking alcohol)

Propanol (rubbing alcohol)
 Sugars
 Water soluble vitamins
B. Carbonyl Group (=CO)
 Polar group
 O can be involved in H-bonding.
 Generally makes molecule water soluble.
Found in:
 Aldehydes: Carbonyl is located at end of molecule
 Ketone: Carbonyl is located in middle of molecule
Examples:

Sugars (Aldehydes or ketones)

Formaldehyde (Aldehyde)

Acetone (Ketone)
Sugars Have Both -OH and =CO Functional Groups
C. Carboxyl Group (-COOH)
 Polar group
 Generally makes molecule water soluble
 Acidic because it can donate H+ in solution
Found in:
 Carboxylic acids: Organic acids, can increase
acidity of a solution. Examples:

Acetic acid: Sour taste of vinegar.

Ascorbic acid (Vitamin C): Found in fruits and
vegetables.

Amino acids: Building blocks of proteins.
D. Amino Group (-NH2)
 Polar group
 Generally makes molecule water soluble
 Weak base because N can accept a H+
 Amine: General term given to compound with (-
NH2)
Found in:
 Amino acids: Building blocks of proteins.
 Urea in urine. From protein breakdown.
Amino acid Structure:

Central carbon with:




H atom
Carboxyl group
Amino group
Variable R-group
Amino Acid Structure:
H
|
(Amino Group) NH2---C---COOH (Carboxyl group)
|
R
(Varies for each amino acid)
Amino Acids Have -NH2 and -COOH Group
The Macromolecules of
Life:
Carbohydrates, Proteins, Lipids, and
Nucleic Acids
Most Biological Macromolecules are Polymers
 Polymer: Large molecule consisting of many identical
or similar “subunits” linked through covalent bonds.
 Monomer: “Subunit” or building block of a polymer.
 Macromolecule: Large organic polymer. Most
macromolecules are constructed from about 70 simple
monomers.

Only about 70 monomers are used by all living things on
earth to construct a huge variety of molecules

Structural variation of macromolecules is the basis for the
enormous diversity of life on earth.
Relatively few monomers are used by cells to make a huge
variety of macromolecules
Macromolecule
Monomers or Subunits
1. Carbohydrates
20-30 monosaccharides
or simple sugars
2. Proteins
20 amino acids
3. Nucleic acids (DNA/RNA) 4 nucleotides (A,G,C,T/U)
4. Lipids (fats and oils)
~ 20 different fatty acids
and glycerol.
Making Polymers
A. Condensation or Dehydration Synthesis reactions:
 Process in which one monomer is covalently linked
to another monomer (or polymer).
 The equivalent of a water molecule is removed.
 Anabolic Reactions: Make large molecules from
smaller ones. Require energy (endergonic)
General Reaction:
Enzyme
X - OH + HO - Y -------->
Monomer 1 Monomer 2
(Unlinked)
(or Polymer)
X - O - Y + H2O
Dimer
Water
(or Polymer)
Example:
Enzyme
Glucose + Fructose ---------> Sucrose +
H2O
Breaking Polymers
B. Hydrolysis Reactions: “Break with water”.
 Break down polymers into monomers.
 Bonds between subunits are broken by adding water.
 Catabolic Reactions: Break large molecules into
smaller ones. Release energy (exergonic)
General Reaction:
Enzyme
X - O - Y + H2O ----------> X - OH + HO - Y
Polymer
Water
Monomer 1 Monomer 2
(or Dimer)
Example:
Sucrose
+
Enzyme
H2O ---------> Glucose + Fructose
Synthesis and Hydrolysis of Sucrose
I. Carbohydrates: Molecules that store energy
and are used as building materials
 General Formula: (CH2O)n
 Simple sugars and their polymers.
 Diverse group includes sugars, starches, cellulose.
 Biological Functions:



Fuels, energy storage
Structural component (cell walls)
DNA/RNA component
 Three types of carbohydrates:
A. Monosaccharides
B. Disaccharides
C. Polysaccharides
A. Monosaccharides: “Mono” single & “sacchar” sugar
 Preferred source of chemical energy for cells
 Can be synthesized by plants from light, H2O and CO2.
 Store energy in chemical bonds.
 Carbon skeletons used to synthesize other molecules.
Characteristics:
1. Have 3-8 carbons. -OH on each carbon; one with C=O
2. Names end in -ose. Based on number of carbons:


5 carbon sugar: pentose
6 carbon sugar: hexose
3. Can exist in linear or ring forms
4. Isomers: Many molecules with the same molecular
formula, but different atomic arrangement

Example: Glucose and fructose are both C6H12O6
Fructose is sweeter than glucose
Monosaccharides Can Have 3 to 8 Carbons
Linear and Ring Forms of Glucose
B. Disaccharides: “Di” double & “sacchar” sugar
 Covalent
bond formed by condensation reaction
between 2 monosaccharides.
Examples:
1. Maltose: Glucose + Glucose.
• Energy storage in seeds.
• Used to make beer.
2. Lactose: Glucose + Galactose.
• Found in milk.
• Lactose intolerance is common among adults.
• May cause gas, cramping, bloating, diarrhea, etc.
3. Sucrose: Glucose + Fructose.
• Most common disaccharide (table sugar).
• Found in plant sap.
Maltose and Sucrose are Disaccharides
C. Polysaccharides: “Poly” many (8 to 1000)
Functions: Storage of chemical energy and
structure.
 Storage polysaccharides: Cells can store simple
sugars in polysacharides and hydrolyze them
when needed.
1. Starch: Glucose polymer (Helical)

Form of glucose storage in plants (amylose)

Stored in plant cell organelles called plastids
2. Glycogen: Glucose polymer (Branched)

Form of glucose storage in animals (muscle and liver
cells)
 Structural Polysaccharides: Used as structural
components of cells and tissues.
1. Cellulose: Glucose polymer.

The major component of plant cell walls.

CANNOT be digested by animal enzymes.
Only microbes have enzymes to hydrolyze cellulose,
found in digestive systems of:



Cows, goats, and rabbits
Termites
2. Chitin: Polymer of an amino sugar (with NH2
group)


Forms exoskeleton of arthropods (insects)
Found in cell walls of some fungi
Three Different Polysaccharides of Glucose
II. Proteins: Large three-dimensional
macromolecules responsible for most cellular
functions
 Polypeptide chains: Polymers of amino acids
linked by peptide bonds in a specific linear
sequence.
 Protein: Macromolecule composed of one or more
polypeptide chains folded into a specific threedimensional conformation.
Proteins Have Important and Varied Functions:
1. Enzymes: Catalysis of cellular reactions
2. Structural Proteins: Maintain cell shape
3. Transport: Transport in cells/bodies (e.g. hemoglobin).
Channels and carriers across cell membrane.
4. Communication: Chemical messengers, hormones, and
receptors.
5. Defensive: Antibodies and other molecules that bind to
foreign molecules and help destroy them.
6. Contractile: Muscular movement.
7. Storage: Store amino acids for later use (e.g. egg white).
Protein function is dependent upon its 3-D shape.
Polypeptide: Polymer of amino acids connected in a
specific sequence
A. Amino acid: The monomer of polypeptides

Central carbon with:




H atom
Carboxyl group
Amino group
Variable R-group
Amino Acid Structure:
H
|
(Amino Group) NH2---C---COOH (Carboxyl group)
|
R
(Varies for each amino acid)
A Protein’s Specific Shape (Conformation)
Determines its Function
Conformation: The 3-D structure of a protein.
Determined by the amino acid sequence.
Four Levels of Protein Structure
1. Primary structure: Linear amino acid sequence,
determined by gene for that protein.
2. Secondary structure: Regular coiling/folding of
polypeptide.


Alpha helix or beta sheet.
Caused by H-bonds between amino acids.
3. Tertiary structure: Overall 3-dimensional shape of a
polypeptide chain.
4. Quaternary structure: Only found in proteins with 2
or more polypeptides.
Overall 3-D shape of all polypeptide chains.

Example: Hemoglobin (2 alpha and 2 beta
polypeptides)
What determines a protein’s Conformation ?
A. Primary structure: Exact location of each amino
acid along the chain determines folding pattern
Example: Sickle Cell Hemoglobin protein

Mutation changes amino acid #6 on the alpha
chain.

Defective hemoglobin causes red blood cells
to assume sickle shape, which damages tissue
and capillaries.

Sickle cell anemia gene is carried in 10% of
African Americans.
B. Chemical & Physical Environment: Presence of
other compounds, pH, temperature, salts.

Denaturation: Process which alters native
conformation and therefore biological
activity of a protein

pH and salts: Disrupt hydrogen, ionic bonds.

Temperature: Can disrupt weak interactions.

Example: Function of an enzyme depends on
pH, temperature, and salt concentration.
III. Nucleic Acids: Store and Transmit
Hereditary Information for All Living Things
 There are two types of nucleic acids in cells:
A. Deoxyribonucleic Acid (DNA)


Has segments called genes which provide information
to make each and every protein in a cell
Double-stranded molecule which replicates each time
a cell divides.
B. Ribonucleic Acid (RNA)



Three main types called mRNA, tRNA, rRNA
RNA molecules are copied from DNA and used to
make gene products (proteins).
Usually exists in single-stranded form.
DNA and RNA are Polymers of Nucleotides
Nucleotide: Subunits of DNA or RNA.
Nucleotides have three components:
1. Pentose sugar (ribose or deoxyribose)
2. Phosphate group to link nucleotides (-PO4)
3. Nitrogenous base (A,G,C,T or U)

Purines: Have 2 rings.



Adenine (A)
Guanine (G)
Pyrimidines: Have one ring.


Cytosine (C)
Thymine (T) in DNA or uracil (U) in RNA.
James Watson and Francis Crick determined the
3-D shape of DNA in 1953
 Double helix: The DNA molecule is a double helix.
 Antiparallel: The two DNA strands run in opposite
directions.


Strand 1: 5’ to 3’ direction (------------>)
Strand 2: 3’ to 5’ direction (<------------)
 Complementary Base Pairing: A & T (U) and G & C.
 A on one strand hydrogen bonds to T (or U in RNA).
 G on one strand hydrogen bonds to C.
 Replication: The double-stranded DNA molecule
can easily replicate based on A=T and G=C pairing.
-- SEQUENCE of nucleotides in a DNA molecule
dictate the amino acid SEQUENCE of polypeptides
DNA is a Double Helix Held Together by H-Bonds
A Gene is a specific segment of a DNA molecule
with information for cell to make one
polypeptide
DNA
(transcribed into single stranded RNA
“copy”)
!
!
mRNA
(single stranded “copy” of the gene)
!
!
Polypeptide (mRNA message translated into
polypeptide)
IV. Lipids: Fats, phospholipids, and steroids
Diverse groups of compounds.
Composition of Lipids:
 C, H, and small amounts of O.
Functions of Lipids:
 Biological fuels
 Energy storage
 Insulation
 Structural components of cell membranes
 Hormones
Lipids: Fats, phospholipids, and steroids
1. Simple Lipids: Contain C, H, and O only.
A. Fats (Triglycerides).


Glycerol : Three carbon molecule with three hydroxyls.
Fatty Acids: Carboxyl group and long hydrocarbon
chains.
 Characteristics of fats:




Most abundant lipids in living organisms.
Hydrophobic (insoluble in water) because nonpolar.
Economical form of energy storage (provide 2X the
energy/weight than carbohydrates).
Greasy or oily appearance.
Lipids: Fats, phospholipids, and steroids
Simple Lipids: Continued
 Saturated fats: Hydrocarbons saturated
with H. Lack -C=C- double bonds.

Solid at room temp (butter, animal fat, lard)
 Unsaturated fats: Contain -C=C- double
bonds.



Usually liquid at room temp (corn, peanut,
olive oils)
Trans fats: Fats that are artificially created by
chemically saturating unsaturated fats.
Margarine (partially hydrogenated oils)
Fats (Triglycerides): Glycerol + 3 Fatty Acids
2. Complex Lipids: In addition to C, H, and O,
also contain other elements, such as
phosphorus, nitrogen, and sulfur.
A. Phospholipids: Are composed of:



Glycerol
2 fatty acids,
Phosphate group
 Amphipathic Molecule


Hydrophobic fatty acid “tails”.
Hydrophilic phosphate “head”.
Function: Primary component of the plasma
membrane of cells
B. Steroids: Lipids with four fused carbon rings
Includes cholesterol, bile salts, reproductive, and
adrenal hormones.

Cholesterol: The basic steroid found in animals




Common component of animal cell membranes.
Precursor to make sex hormones (estrogen, testosterone)
Generally only soluble in other fats (not in water)
Too much increases chance of atherosclerosis.
C. Waxes: One fatty acid linked to an alcohol.


Very hydrophobic.
Found in cell walls of certain bacteria, plant and
insect coats. Help prevent water loss.
Cholesterol: The Basic Steroid in Animals