Transcript Chapter 2: Chemical Principles

Chapter 2
Chemical Principles
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
 Hydrogen
of all organic compounds.
(H): 10% of body mass.
 Component
 Nitrogen
(N): 3% of body mass.
 Component
 Calcium
 Bones,
of proteins and nucleic acids (DNA/RNA)
(Ca): 1.5% of body mass.
teeth, clotting, muscle and nerve function.
 Phosphorus
 Bones,
of water and most organic copounds.
(P): 1% of body mass
nucleic acids, energy transfer (ATP).

Minor Elements: Found in low amounts.
Between 1% and 0.01%.
 Potassium
 Nerve
(K): Main positive ion inside cells.
and muscle function.
 Sulfur
(S): Component of most proteins.
 Sodium (Na): Main positive ion outside cells.
 Fluid
balance, nerve function.
 Chlorine
 Fluid
(Cl): Main negative ion outside cells.
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
Proton (p+) In nucleus
Neutron (no) In nucleus
Electron (e-) Outside nucleus
Mass
Charge
1
1
0*
+1
0
-1
* Mass is negligible for our purposes.
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 an atom can 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 bond 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:
Molecules of water “stick together”
 Adhesion: Water molecules stick to walls of tubes.
 Cohesion:
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
 Used
water to “bead” into tiny balls
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:
 Solvent:
Dissolving substance of a solution.
 Example:
 Solute:
Salt water, air, tap water.
Water, alcohol, oil.
Substance dissolved in the solvent.
 Example:
 Aqueous
NaCl, sugar, carbon dioxide.
solution: Water is the solvent.
 Solubility: Ability
given solvent.
of substance to dissolve in a
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
(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
 Acidic
 Basic
 As
solution: pH is 7. [H+ ] = [OH-]
solution: pH is less than 7. [H+ ] > [OH-]
solution: pH is greater than 7. [H+ ] < [OH-]
[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
2Hydrogen Oxygen
Molecules 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)
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 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
 Can
group: Polar covalent bond between O and H.
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
O
group
can be involved in H-bonding.
 Generally
makes molecule water soluble.
Found in:
 Aldehydes:
 Ketone:
Carbonyl is located at end of molecule
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
 Acidic
makes molecule water soluble
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
 Weak
makes molecule water soluble
base because N can accept a H+
 Amine:
General term given to compound with
(-NH2)
Found in:
 Amino
 Urea
acids: Building blocks of proteins.
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 Groups
The Macromolecules of Life:
Carbohydrates, Proteins, Lipids, and
Nucleic Acids
3. 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
X - O - Y + H2O
Dimer
Water
(Unlinked)
(or Polymer)
(or Polymer)
Example:
Enzyme
Glucose + Fructose ---------> Sucrose +
(Monomer) (Monomer)
(Dimer)
H2O
Water
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:
Enzyme
Sucrose
(Dimer)
+
H2O ---------> Glucose + Fructose
Water
(Monomer) (Monomer)
Synthesis and Hydrolysis of Sucrose
IV. Carbohydrates: Molecules that store energy and
are used as building materials
 General
 Simple
Formula: (CH2O)n
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.
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
V. 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:
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:
 Example:
Can disrupt weak interactions.
Function of an enzyme depends on
pH, temperature, and salt concentration.
VI. 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)
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
 Nucleic acid: A polymer 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 3D 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)
VII. 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
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)
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