Chapter 2The 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 elementsChemical 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 AtomsAtoms: Smallest particle of an element that retains its chemical properties. Made up of three main subatomic particles.

Particle Location Mass ChargeProton (p+) In nucleus 1 +1Neutron (no) In nucleus 1

0Electron (e-) Outside nucleus 0* -1

* Mass is negligible for our purposes.

Structure and Properties of Atoms

1. Atomic number = # protons The number of protons is unique for each

elementEach 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) Electron Capacity1 2 (Closest to nucleus, lowest

energy)

2 83 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.

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 atomsSubscripts 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 e-

1(2), 2(8), 3(1) 1(2), 2(8)Outer shell has 1 e- Outer shell is fullSodium atom Sodium ion

Number of Valence Electrons Determine the Chemical Behavior of Atoms

Element Valence CombiningTendency

Electrons CapacitySodium 1 1 Lose 1Calcium 2 2 Lose 2Aluminum 3 3 Lose 3Carbon 4 4 Share 4Nitrogen 5 3 Gain 3Oxygen 6 2 Gain 2Chlorine 7 1 Gain 1Neon* 8 0 Stable

* Noble gas

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+ Cl-

sodium 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=O3. 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 ElectronegativityO 3.5N 3.0

S & C 2.5P & H 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 PolarityA. 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 RegulationWater 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 WaterIce 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 reactionsSolution: 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+ + Cl-

H2SO4 (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+ + OH-

Directly 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+ + Cl-

NaOH (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+ + HCO3-

carbonic 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 Oxygen 2 WaterMolecules Molecule 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.

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 CompoundsHydrocarbons:

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

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) Isobutane (C4H10)

CH3--CH2--CH2--CH3 CH3--CH--CH3 |CH3

Isomers have different physical and chemical properties.

II. Functional Groups Determine Chemical & Physical Properties of Organic MoleculesCompounds 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 atomCarboxyl groupAmino groupVariable 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 Groups

The Macromolecules of Life:

Carbohydrates, Proteins, Lipids, and Nucleic Acids

Most Biological Macromolecules are PolymersPolymer: 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 monosaccharidesor 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 PolymersA. 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 --------> X - O - Y + H2OMonomer 1 Monomer 2 Dimer

Water(Unlinked) (or Polymer) (or Polymer)

Example: Enzyme

Glucose + Fructose ---------> Sucrose + H2O

Breaking PolymersB. 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 - YPolymer Water Monomer 1 Monomer 2

(or Dimer)

Example: Enzyme

Sucrose + 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. MonosaccharidesB. Disaccharides C. Polysaccharides

A. Monosaccharides: “Mono” single & “sacchar” sugar

Preferred source of chemical energy for cellsCan 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=O2. Names end in -ose. Based on number of carbons:

5 carbon sugar: pentose 6 carbon sugar: hexose

3. Can exist in linear or ring forms4. 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. B. DisaccharidesDisaccharides: : ““DiDi”” double & double & ““saccharsacchar”” sugar sugar

Covalent bond formed by condensation reaction Covalent bond formed by condensation reaction between 2 monosaccharides.between 2 monosaccharides.

ExamplesExamples::

1. 1. MaltoseMaltose: Glucose + Glucose. : Glucose + Glucose.

• Energy storage in seeds. Energy storage in seeds.

• Used to make beer.Used to make beer.

2. 2. LactoseLactose: Glucose + Galactose. : Glucose + Galactose.

• Found in milk.Found in milk.

• Lactose intoleranceLactose intolerance is common among adults. is common among adults.• May cause gas, cramping, bloating, diarrhea, etc.May cause gas, cramping, bloating, diarrhea, etc.

3. 3. SucroseSucrose: Glucose + Fructose. : Glucose + Fructose.

• Most common disaccharide (table sugar). Most common disaccharide (table sugar).

• Found in plant sap.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

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.

Proteins Have Important and Varied Functions:Proteins Have Important and Varied Functions:

1.1. EnzymesEnzymes: : Catalysis of cellular reactionsCatalysis of cellular reactions

2. 2. Structural ProteinsStructural Proteins: : Maintain cell shapeMaintain cell shape

3. 3. TransportTransport: : Transport in cells/bodies (e.g. hemoglobin). Transport in cells/bodies (e.g. hemoglobin).

Channels and carriers across cell membrane.Channels and carriers across cell membrane.

4. 4. CommunicationCommunication: : Chemical messengers, hormones, and Chemical messengers, hormones, and

receptors.receptors.

5. 5. DefensiveDefensive: : Antibodies and other molecules that bind to Antibodies and other molecules that bind to

foreign molecules and help destroy them.foreign molecules and help destroy them.

6. 6. ContractileContractile: : Muscular movement.Muscular movement.

7. 7. StorageStorage: : Store amino acids for later use (e.g. egg white).Store amino acids for later use (e.g. egg white).

Protein functionProtein function is is dependentdependent upon its upon its 3-D shape3-D shape..

Polypeptide: Polymer of amino acids connected in a specific sequence

A. Amino acid: The monomer of polypeptides

Central carbon with:H atomCarboxyl groupAmino groupVariable 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 NucleotidesNucleotide: 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 fuelsEnergy storageInsulationStructural components of cell membranesHormones

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)

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 ringsIncludes 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.