Download Chapter 2

Chapter 2
Biology 25: Human
Biology
Prof. Gonsalves
Los Angeles City College
Loosely Based on Mader’s Human Biology,7th edition
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).
• 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
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.
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.
Electron Arrangements of Important
Elements of Life
1 Valence electron
4 Valence electrons
5 Valence electrons
6 Valence electrons
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
Tendency
Sodium
Calcium
Aluminum
Carbon
Nitrogen
Oxygen
Chlorine
Neon*
* Noble gas
Valence
Combining
Electrons
1
2
3
4
5
6
7
8
Capacity
1
2
3
4
3
2
1
0
Lose 1
Lose 2
Lose 3
Share 4
Gain 3
Gain 2
Gain 1
Stable
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: Involves 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
formed by important elements:
Carbon (4)
Nitrogen (3)
Oxygen (2)
Sulfur (2)
Hydrogen (1)
Two Types of Covalent Bonds: Polar and Nonpolar
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
A. 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, Cl2
Nonpolar Covalent Bonds: Electrons are
Shared Equally
Polar and Nonpolar Covalent Bonds
B. 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: 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 bond are weak, but the cumulative effect
of many H bonds is very strong.
• H bonds only last a fraction of a second, but at any
moment most molecules are hydrogen bonded to
others.
Unique properties of water caused by H-bonds
– Cohesion: Water molecules stick to each other.
This causes surface tension.
• Film-like surface of water is difficult to break.
• Used by some insects that live on water surface.
• Water forms beads.
– Adhesion: Water sticks to many surfaces.
Capillary Action: Water tends to rise in narrow
tubes. This is caused by cohesion and adhesion
(water molecules stick to walls of tubes).
Examples: Upward movement of water through plant
vessels and fluid in blood vessels.
Unique properties of water caused by H-bonds
– Expands when it freezes.
• Ice forms stable H bonds, each molecule is bonded to
four neighbors (crystalline lattice). Water does not
form stable H bonds.
• Ice is less dense than water.
• Ice floats on water.
• Life can survive in bodies of water, even though the
earth has gone through many winters and ice ages
Unique properties of water caused by H-bonds
– Stable Temperature: Water resists changes in
temperature because it has a 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/oC
• High heat of vaporization: Water must absorb large
amounts of energy (heat) to evaporate.
– Heat of Vaporization of Water: 540 calorie/gram.
• Evaporative cooling is used by many organisms to
regulate body temperature.
– Sweating
– Panting
Unique properties of water caused by H-bonds
– Universal Solvent: Dissolves many (but not all)
substances to form solutions.
Solutions are homogeneous mixtures of two or
more substances (salt water, air, tap water).
All solutions have at least two components:
• Solvent: Dissolving substance (water, alcohol, oil).
– Aqueous solution: If solvent is water.
• Solute: Substance that is dissolved (salt, sugar, CO2).
– Water dissolves polar and ionic solutes well.
– Water does not dissolve nonpolar solvents well.
Solubility of a Solute Depends on its
Chemical Nature
Solubility: Ability of substance to dissolve in a
given solvent.
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”
Solubility of a Solute Depends on its
Chemical Nature
Two Types of Solutes:
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+ + 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+ + 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
(G.T. 99% in this form)
Bicarbonate ion
(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+].
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:
2 H2 +
O2 -----------> 2H2O
2 Hydrogen
Molecules
Oxygen
Molecule
2 Water
Molecules
Organic Chemistry: Carbon Based Compounds
A. Inorganic Compounds: Compounds without carbon.
B. Organic Compounds: Compounds synthesized by cells
and containing carbon (except for CO and CO2).
– Diverse group: Several million organic compounds
are known and more are identified every day.
– 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.
Carbon: unique element for basic building block of
molecules of life
• Carbon has 4 valence electrons: Can form four
covalent bonds
– Can form single , double, triple bonds.
– Can form large, complex, branching
molecules and rings.
– Carbon atoms easily bond to C, N, O, H, P, S.
• Huge variety of molecules can be formed based
on simple bonding rules of basic chemistry
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 (Natural gas).
• (2C) Ethane:
CH3CH3
• (3C) Propane:
CH3CH2CH3 (Gas grills).
• (4C) Butane:
CH3CH2CH2CH3 (Lighters).
• (5C) Pentane:
CH3CH2CH2CH2CH3
• (6C) Hexane:
CH3CH2CH2CH2CH2CH3
• (7C) Heptane:
CH3CH2CH2CH2CH2CH2CH3
• (8C) Octane:
CH3CH2CH2CH2CH2CH2CH2CH3
Functional groups play pivotal role in chemical &
physical properties of organic molecules
Compounds that are made up solely of carbon
and hydrogen are not very reactive.
Functional groups:
– One or more H atoms of the carbon skeleton
may be replaced by a functional group.
– Groups of atoms that have unique chemical
and physical properties.
•
Usually a part of molecule that is chemically active.
•
Similar activity from one molecule to another.
•
Together with size and shape, determine unique
bonding and chemical activity of organic molecules.
Functional Groups Determine Chemical &
Physical Properties of Organic Molecules
Four Important Functional Groups:
–
•
Hydroxyl (-OH)
•
Carbonyl (=C=O)
•
Carboxyl (-COOH)
•
Amino (-NH2)
Notice that all four functional groups
are polar.
I. 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.
Making and Breaking Polymers
– There are two main chemical
mechanisms in the production and
break down of macromolecules.
• Condensation or Dehydration Synthesis
• Hydrolysis
– In the cell these mechanisms are
regulated by enzymes.
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.
III. 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 (glucose)
Can be synthesized by plants from light, H2O and CO2.
Store energy in chemical bonds.
Carbon skeletons used to synthesize other molecules.
Characteristics:
1. May have 3-8 carbons. -OH on each carbon; one with C=0
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.
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.
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.
2. Chitin: Polymer of an amino sugar (with NH2
group)
• Forms exoskeleton of arthropods (insects)
• Found in cell walls of some fungi
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
Types of Fats
– 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)
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 acid
• 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.
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 SPECIFIC
3-D conformations
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
– H atom
– Carboxyl group
– Amino group
– Variable R-group
Protein Function is dependent upon Protein Structure
(Conformation)
CONFORMATION: The 3-D shape of a protein is
determined by its 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-D shape of a
polypeptide chain.
4. Quaternary structure: Only in proteins with 2 or
more polypeptides. Overall 3-D shape of all chains.
• Example: Hemoglobin (2 alpha and 2 beta
polypeptides)
Nucleic acids store and transmit hereditary information
for all living things
There are two types of nucleic acids in living things:
A. Deoxyribonucleic Acid (DNA)
• Contains genetic information of all living
organisms.
• 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 that determine
the primary structure of proteins
• 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) and 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