The Chemistry of Microbiology

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• Chapter 2
• The Chemistry of Microbiology

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• 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).

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• Main Elements Over 98 of an organisms 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.

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• 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.

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• 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)

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• 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 Charge
• Proton (p) In nucleus 1
  1
• Neutron (no) In nucleus 1
  0
• Electron (e-) Outside nucleus 0
  -1
• Mass is negligible for our purposes.

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Atomic Particles Protons, Neutrons, and
Electrons Helium Atom Carbon Atom
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• 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.

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• 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

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• 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.

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• Electrons Determine How Atoms Bond with Other
  Atoms
• A. Energy levels Electrons occupy different
  energy levels around the nucleus.
• Level (Shell) Electron Capacity
• 1 2 (Closest to nucleus, lowest energy)
• 2 8
• 3 8 (If valence shell, 18 otherwise)
• 4, 5, 6 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.

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• 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

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• 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 -------gt Na 1 e-
• 1(2), 2(8), 3(1) 1(2), 2(8)
• Outer shell has 1 e- Outer shell is full
• Sodium atom Sodium ion

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• Number of Valence Electrons Determine the
  Chemical Behavior of Atoms
• Element Valence Combining Tendency
• Electrons Capacity
• Sodium 1 1 Lose 1
• Calcium 2 2 Lose 2
• Aluminum 3 3 Lose 3
• Carbon 4 4 Share 4
• Nitrogen 5 3 Gain 3
• Oxygen 6 2 Gain 2
• Chlorine 7 1 Gain 1
• Neon 8 0 Stable
• Noble gas

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Electron Arrangements of Important Elements of
Life
1 Valence electron
4 Valence electrons
5 Valence electrons
6 Valence electrons
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• 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

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• 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 ----------gt Na Cl-
• sodium chlorine Table salt
• (Sodium chloride)
• Two Types of Ions
• Anions Negatively charged particle (Cl-)
• Cations Positively charged particle (Na)

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• 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

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• 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)
• OO
• 3. Triple bond Three electron pairs shared
  between two atoms.
• Example Nitrogen gas (N2)
• N N

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• Number of Covalent Bonds
• Carbon (4)
• Nitrogen (3)
• Oxygen (2)
• Sulfur (2)
• Hydrogen (1)

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• Two Types of Covalent Bonds Polar and Nonpolar
• A. Electronegativity A measure of an atoms
  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 Electronegativity
• O 3.5
• N 3.0
• S C 2.5
• P H 2.1

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• 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.

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• 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).

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• Water - A Unique Compound for Life

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• Water The Ideal Compound for Life
• Living cells are 70-90 water
• Water covers 3/4 of earths surface
• Water is the ideal solvent for chemical reactions
• On earth, water exists as gas, liquid, and solid

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• 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.

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• 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.

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• II. Biological Consequences of Waters 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

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• 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 -gt gas is high (540
  calories/gram)
• Sweating is a form of evaporative cooling.
• Can regulate temperature w/o great water loss.

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• 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

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• 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.

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• 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.

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• 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 ) -------gt H Cl-
• H2SO4 (into H2O ) --------gt H HSO4-
• Acids INCREASE the relative H of a solution.
• Water can also dissociate into ions, at low
  levels
• H2O ltgt H OH-

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• 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 ltgt H OH-
• Directly NH3 H lt------gt NH4
• Indirectly NaOH ---------gt Na OH-
• ( H OH- ltgt H2O )

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• Strong acids and bases Dissociation is almost
  complete (99 or more of molecules).
• HCl (aq) -------------gt H Cl-
• NaOH (aq) -----------gt 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 ltgt H HCO3-
• carbonic acid
  Bicarbonate ion
• (G.T. 99 in this form) (L.T. 1 in
  dissociated form)

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• 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 gt
  OH-
• Basic solution pH is greater than 7. H lt
  OH-
• As H increases pH decreases (inversely
  proportional).
• Logarithmic scale Each unit on the pH scale
  represents a ten-fold change in H.

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pH of Common Solutions
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• 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.

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• 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 -----------gt 2H2O
• 2 Hydrogen Oxygen 2 Water
• Molecules Molecule Molecules

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• Organic Compounds

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• 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.

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Organic Compounds are Carbon Based
Carbon Has 4 Valence Electrons and Can Form 4
Covalent Bonds
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• 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)

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• Carbon Skeletons of Organic Compounds

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• 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 CH4
• (2C) Ethane CH3CH3
• (3C) Propane CH3CH2CH3
• (4C) Butane CH3CH2CH2CH3
• (5C) Pentane CH3CH2CH2CH2CH3
• (6C) Hexane CH3CH2CH2CH2CH2CH3
• (7C) Heptane CH3CH2CH2CH2CH2CH2CH3
• (8C) Octane CH3CH2CH2CH2CH2CH2CH2CH3

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Hydrocarbons have C and H only
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• Isomers Compounds with same chemical formula but
  different structures
• Structural Isomers Differ in atom arrangement
• Example Isomers of C4H10
• Butane (C4H10) Isobutane (C4H10)
• CH3--CH2--CH2--CH3 CH3--CH--CH3
• CH3
• Isomers have different physical and chemical
  properties.

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• 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 (CO)
• Carboxyl (-COOH)
• Amino (-NH2)
• Notice that all are polar.

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• 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

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• 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)

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Sugars Have Both -OH and CO Functional Groups
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• 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.

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• 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.

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• 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)

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• Amino Acids Have -NH2 and -COOH Groups

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• The Macromolecules of Life
• Carbohydrates, Proteins, Lipids, and Nucleic Acids

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• 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.

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• 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.

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• 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 --------gt X - O - Y
  H2O
• Monomer 1 Monomer 2 Dimer
  Water
• (Unlinked) (or Polymer)
  (or Polymer)
• Example
• Enzyme
• Glucose Fructose ---------gt Sucrose
  H2O

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• 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 ----------gt X - OH
  HO - Y
• Polymer Water Monomer 1 Monomer 2
• (or Dimer)
• Example
• Enzyme
• Sucrose H2O ---------gt Glucose
  Fructose

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Synthesis and Hydrolysis of Sucrose
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• 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

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• 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
  CO
• 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

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Monosaccharides Can Have 3 to 8 Carbons
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Linear and Ring Forms of Glucose
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• 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.

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Maltose and Sucrose are Disaccharides
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• 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)

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• 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

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Three Different Polysaccharides of Glucose
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• 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
  three-dimensional conformation.

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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.
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• 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)

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• A Proteins 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.

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• 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)

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• What determines a proteins 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.

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• 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.

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• 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.

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• 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.

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• 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 (------------gt)
• Strand 2 3 to 5 direction (lt------------)
• 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 AT and GC pairing.
• 
  ---
• SEQUENCE of nucleotides in a DNA molecule dictate
  the amino acid SEQUENCE of polypeptides

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DNA is a Double Helix Held Together by H-Bonds
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• 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)

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• 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

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• 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.

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• Lipids Fats, phospholipids, and steroids
• Simple Lipids Continued
• Saturated fats Hydrocarbons saturated with H.
  Lack -CC- double bonds.
• Solid at room temp (butter, animal fat, lard)
• Unsaturated fats Contain -CC- 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)

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Fats (Triglycerides) Glycerol 3 Fatty Acids
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• 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

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• 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.

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Cholesterol The Basic Steroid in Animals