CHAPTER 2
The Chemical Basis of Life
MOLECULAR BASIS FOR LIFE

All living organisms are composed of matter


Defined as any substance that has mass and takes up
space
Matter is generally found in 1 of 3 states
Gas – loosely packed and highly movable
 Liquid – less densely packed and move more rapidly
 Solid – densely packed with minimal movement

ELEMENTS

Fundamental forms of matter

Can’t be broken apart by normal means

92 occur naturally on Earth

About 25 are essential to life


96% of human body, as well as other living
organisms, from 4 elements (CHON) or more
Trace elements: necessary in minute quantities

Iron (Fe) needed by all life

Iodine (I) needed by all vertebrates
WHAT DOES CHEMISTRY HAVE TO
DO WITH BIOLOGY?
Human
Oxygen
Carbon
Hydrogen
Nitrogen
Calcium
Phosphorus
Potassium
Sulfur
Earth’s Crust
61.0%
Oxygen
23.0
96% Silicon
10.0
Aluminum
2.6
Iron
1.4
Calcium
1.1
Magnesium
0.2
Sodium
0.2
Potassium
46.0%
27.0
8.2
6.3
5.0
2.9
2.3
1.5
Seawater
Oxygen
Hydrogen
Chlorine
Sodium
Magnesium
Sulfur
Calcium
Potassium
85.7%
10.8
2.0
1.1
0.1
0.1
0.04
0.03
ATOMS



Smallest particles that retains properties of an
element
Made up of subatomic particles:

Protons (+) in nucleus

Electrons (-) orbits nucleus

Neutrons (no charge) in nucleus
Protons and neutrons


Mass of about 1 dalton (atomic mass unit, amu)
Electrons

Mass is negligible (1/2000 amu)
COMPOUNDS


Contains 2 or more different elements in a fixed
ratio

Na (metal) + Cl (poisonous gas) = NaCl (table salt)

2 H (gas) + O (gas) = H2O (water)

glucose+fructose = sucrose; glucose+glucose=maltose
Demonstrates emergent properties
READING A PERIODIC TABLE
Elements differ depending on the number of
subatomic particles
 Atomic symbol

1st letter or 2 (usually)
 Familiar with the 1st 20


Atomic number

Determined by number of protons



Neutral atoms contain equal # of electrons
Element specific
Mass number
Determined by number of protons + neutrons
 Atomic mass: total mass of an atom (includes electrons)


About equal to mass number
ISOTOPES
Naturally, elements are a mix of isotopes
 Behave the same as respective element

Same number of protons (atomic number)
 Different numbers of neutrons
 Mass number varies


Can be stable or unstable (radioactive)

Nucleus decays spontaneously, releasing particles of
energy to form other elements (changes proton #)

Nitrogen decays to 14Carbon
RADIOISOTOPES AND THE MEDICAL FIELD

Basic research



Dating fossils
Biological tracers
Medical diagnosis
Brain scanning
 Cancer treatments


Dangers
Radioactive atoms give off energy that destroys
chemical bonds when they collide
 Can damage DNA, which changes genetic
information
 Nuclear warfare

ENERGY OF AN ATOM
Energy is the capacity to cause change
 Potential energy is the energy that matter has
because of its location or structure
 The electrons of an atom differ in their amounts
of potential energy
 An electron’s state of potential energy is called its
energy level, or electron shell
 Only the electrons of an atom are involved in
chemical reactions

ELECTRON ARRANGEMENT

Electron orbitals
Orbitals closest to nucleus are lower energy and are
filled first
 Can hold up to2 or 8 electrons
 Atoms differ in the number of occupied orbitals
(shells)
 Outermost electrons (valence) determine the
properties



Similar valence shells share similar properties
Electrons can move between orbitals



Absorbing energy moves an electron up a level
Losing energy drops it down a level
Absorption or loss must be = to difference in PE
between levels

Lost energy released as heat to environment
ELECTRON SHELL MODEL
electron
SODIUM
11p+ , 11e-
CHLORINE
17p+ , 17e-
CARBON
6p+ , 6e-
OXYGEN
8p+ , 8e-
HYDROGEN
1p+ , 1e-
HELIUM
2p+ , 2e-
proton
neutron
NEON
10p+ , 10e-
Fig. 2-6, p.23
CHEMICAL BONDS
Interactions between electrons that hold atoms
together
 Valence = bonding capacity
 Forms molecules


Can be same or different elements

O2 or H20
Atoms with unfilled outer shells are chemically
reactive
 With filled outer shells are chemically inert
 Created by sharing, donating, or receiving
electrons to complete outer shells
 3 Types

COVALENT BOND

Atoms share a pair or
pairs of electrons to fill
outermost shell (valence
electrons)


Single, double, or triple
covalent bond
Non-polar
Atoms share electrons
equally
 Example: Hydrogen gas
(H-H)


Polar
Electrons spend more
time near most
electronegative element
 Water

IONIC BOND
One atom loses electrons, becomes positively
charged ion (cation)
 Another atom gains these electrons, becomes
negatively charged ion (anion)
 Charge difference attracts the two ions to each
other


Actual bond not formed, but is able to occur
Very weak bond
 Also called salts


E.g. NaCl (table salt)
IONIC BOND
One atom loses electrons, becomes positively
charged ion
 Another atom gains these electrons, becomes
negatively charged ion
 Charge difference attracts the two ions to each
other


Actual bond not formed, but is able to occur
Very weak bond
 Salt e.g NaCl

cation
anion
WEAK BONDS
Ionic bonds in water
 Hydrogen bonds



Hydrogen covalently bonded to one atom is attracted
to another atom
Van der Waals interactions
Only occur when atoms and molecules are close
together
 Momentary uneven distribution of charge


Both individually weak, but cumulatively strong
BIOLOGICAL IMPORTANCE OF MOLECULAR
SHAPE
Determines how molecules are recognized and
how they interact
 Endorphins, codeine, and morphine

Single change in functional group
 All bind to same receptors
 All produce analgesic effects
 Endorphins don’t bind long enough to allow tolerance
or withdrawl symptoms to occur

MAKING AND BREAKING BONDS
Process is a chemical reaction, often reversible
 Starting materials are reactants, ending are products
 # of atoms are conserved on both sides of the reaction



Matter can’t be created or destroyed, only rearranged
Chemical equilibrium: forward and reverse reactions
occurring at the same rate (no net change)