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Chapter 2
The Chemical Basis of Life
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 compounds.
(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.
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
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
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+)
Ionic Bonding: Sodium Chloride
A Crystal of Sodium Chloride:
Ions are Held Together by Ionic Bonds
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
--
Single and Double Covalent Bonds
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.
Polar Covalent Bonds: Electrons are
Shared Unequally Creating Partial Charges
Water
Molecule
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).
Hydrogen Bonds: Weak Attractions between
Hydrogen and Partially Negative Atoms
Water
Molecules
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.
H
bonds only last a fraction of a second, but at any
moment most molecules are hydrogen bonded to
others.
Hydrogen Bonds in Water are Responsible for
Many of its Properties
Unique properties of water caused by H-bonds
 Cohesion:
Water molecules stick to each other.
This causes surface tension.
 Film-like
 Used
surface of water is difficult to break.
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
 Ice
when it freezes.
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
Ice Forms Stable Hydrogen Bonds and is
Less Dense than Water
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.
A Salt Crystal Dissolving in Water
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+ + 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:
2 H2 +
O2 -----------> 2H2O
2 Hydrogen Oxygen
Molecules Molecule
2 Water
Molecules
Chemical Reactions Require Making
and Breaking Bonds