The Chemical Basis of Life
Chapter 2
Overview
• Atoms
• Combining Matter
– Physically
– Chemically
• Water
• Acids, Bases, and pH
• Buffers
Matter and Energy
Matter:
– Occupies space
– Has mass: liquid, gas, solid
Energy:
– Capacity to do work
– Measured by effect on matter
Chemistry
Science of the structure of matter
Central to all other sciences
Chemistry is part of all living & non-living
things
Life requires ~25 chemical elements
Humans & other living organisms differ from
non-living things in elemental composition
96% of body weight made up of C, H, O, N
Other 4%: Ca, P, K, S, Na, Cl, Mg & trace
elements essential for life (e.g. Fe)
Elements
Basic units of all matter
Can’t be broken down to simpler substances using
ordinary chemical methods
112 known elements → periodic table
Each element is represented by its atomic
symbol
(1st letter(s) of element’s name)
e.g. carbon = C
hydrogen = H
oxygen = O
In nature, few elements exist in pure form
(tend to form compounds)
Emergent properties:
e.g. NaCl
Na (metal) + Cl (poisonous gas) =
NaCl (table salt)
Atoms
Building blocks of elements
Unique to each element
Give it specific physical and chemical properties
Physical properties:
Colour, texture, boiling point, melting point, etc.
Chemical properties:
The ways that atoms interact with other atoms
Made up of protons (p+), neutrons (no),
electrons (e-)
ep+
no
Protons (p+) have positive charge
Neutrons (no) have no overall charge
Both are heavy particles with approximately
same mass
Electrons (e-) have negative charge
Do not contribute to atomic mass
(1/2000th mass of proton)
In general, # protons = # electrons
No net electrical charge
Generalized Atom
Mass # = p+ + no
a
H
b
Atomic # = p+
Periodic table is ordered by atomic number
The 3 Smallest Atoms
1
1
4
7
2
3
H He Li
p+
1
2
3
no
0
2
4
e-
1
2
3
Atomic Mass
Approximately equal to mass number
(# p+ + # no) because e-s weigh so little
In general, atomic weight is about equal to
mass # of most abundant isotope
e.g. atomic mass of H = 1.008
(indicates that 1H is present in much greater amounts
than 2H or 3H forms)
Isotopes
Different versions of same element
Occur with most natural elements
Differ in # of neutrons
(same atomic # but different mass #)
If stable, nucleus remains intact
If unstable, is radioactive
Radioisotopes
Nuclei decompose spontaneously into more
stable forms
e.g. 14C: half-life of 5700 years
½ atoms turn into 13N
Used to date rocks and biological remains
Releases particles & energy
(breaks chemical bonds in living organisms)
Damaging to live tissue but used in biological
research & medicine
Structure of an Atom
Nucleus contains protons & neutrons
Electrons move around nucleus
= electron cloud
Atomic orbitals organized into shells
e.g. 11Na
Higher-energy shells hold more e-s (2n2) &
are located further from nucleus
Shells fill up in order of increasing energy
e-s can be excited up to higher energy level
for brief periods
Spontaneously return to lower level while
emitting the energy gained via excitation
e-s in outer (valence) shell dictate chemical
behaviour
(these ones interact with those from other
atoms)
Regardless of # of e-s in each shell, # that
can participate in bonding is 8
= octet rule
The Octet Rule
Atoms want to gain, lose, or share e-s so
that have 8 electrons in outer shell
Exception = H
(only has room in 1st energy level for 2 e-s)
Use atomic number to calculate how many
e-s are available for bonding
• 1st energy level = 2 e-s / shell
• 2nd and up = 8 e-s / shell
e.g. 6C:
Has 4 e-s in outer shell; wants to gain 4 e-s to fill
shell for a total of 8 e-s
7N:
8O:
Has 5 e-s in outer shell
Has 6 e-s in outer shell
Needs 3
Needs 2
11Na:
Has 1 e- in outer shell
Needs 7
17Cl:
Has 7 e-s in outer shell
Needs 1
Combining Matter
Most atoms do not exist in free state
Chemically combine with other atoms to form
molecules
If atoms are the same
= molecule of element e.g. O2
If atoms are different
= compound e.g. H2O
Molecular Formulas
A molecule’s chemical composition is
written as a formula
Symbols for elements
Subscripts for number of atoms of each
element
e.g. H20 = 2 H, 1 O
e.g. 5 H20 = 10 H, 5 O
Ways to Represent Compounds
e.g. methane (CH4)
Structural formula
Ball-and-stick model
Space-filling model
Special Structure: Carbon Ring
=
If icon for ring shows no atoms, assume
that C occupies each corner
Same goes for 5-carbon rings
Mixtures
2 or more substances
No chemical bonding
= physical intermixing
Living material contains 3 types:
Solutions
Colloids
Suspensions
Mixture #1: Solution
Homogeneous
Transparent
Does not settle out
Solvent
– Present in largest quantity
– Usually liquid
– Water is body’s principle solvent
Solute
– Present in smaller quantity
Mixture #2: Colloid
Heterogeneous
Translucent or milky
Does not settle out
Can undergo sol-gel
transformation
e.g. cytosol in cells
Mixture #3: Suspension
Heterogenous
Settles out
e.g. blood settles out into
plasma & cells
Chemical Bonds
Attractive forces between atoms
Inert if outer e- valence
shell is filled
– Do not tend to form
bonds
e.g. He
Reactive if outer shell is
not filled
– React with other atoms
to gain / lose / share e-s
to fill shells
e.g. O
Ionic Bonds
Transfer of e- s from one atom to another
Become ions (charged particles)
Gain e- → negative charge = anion
Lose e- → positive charge = cation
Both become stable & combine to form
ionic compound (a.k.a. salt)
Salts
Release ions other than H+ and OH-
Usually form when acids and bases mix
Dissociate in water into component ions
(electrolytes that can conduct electricity)
Important in living organisms:
e.g. Na+, K+, Ca2+ used in nerve transmission, muscle
contraction
e.g. plant cells use salts to take up water from soil
Covalent Bonds
E- sharing
Each atom fills outer shell part of the time
Can be single, double, or triple bonds
e.g. H2: H-H;
O2: O=O;
N2: NN
Can be polar or non-polar bonds
Atoms can be:
Electropositive:
1-2 valence shell e-s
Tend to lose e-s
Electronegative:
6-7 valence shell e-s
Tend to attract e-s strongly
Electrically balanced
Non-polar Covalent Bonds
Electrically balanced
Equal sharing of e-s
O
C
O
Polar Covalent Bonds
Unequal e- sharing
One element has more protons
= stronger pull on e-s
= has e-s more of the time
= slightly electronegative
Results in molecule with + & - charges at either end
Often occurs when atoms are of different sizes
O
H
+
H
+
Hydrogen Bonds
Not a true bond
= can’t form molecules
Attraction between covalently-bound
H atom & electronegative atom
(can be different molecule or
different area of same molecule)
e.g. between water molecules,
between complementary bases in
DNA
Mixtures vs. Compounds
TYPE
Mixture
Compound
“BOND”
Physical mixing
Chemical
Physical means
Chemical
means
SEPARATION
BY:
Homogeneous
COMPOSITION
or
heterogeneous
Homogeneous
Water’s Life-Giving Properties
The universal solvent
Water is important because:
• Life originated in it
• All known living things depend on water
(metabolic processes, respiration, photosynthesis)
• Maintains cell structure/shape
Characteristics of Water
•
•
•
•
•
•
•
•
Polar molecules
Specific heat capacity
Heat of vaporization
Density of water
Cohesion
Adhesion
Surface tension
Good solvent
All result from H-bonding
Polarity of the Water Molecule
-
One end slightly positive, other
slightly negative
= no net charge
Attracts other water molecules
(cohesion)
Attracts sugar & other polar
(hydrophilic) molecules
Repels oil & other non-polar
(hydrophobic) molecules
+
+
Why is Polarity Important?
If water were linear (non-polar), not bent (polar):
– It would not liquify except at high pressures
– It would probably not remain liquid over more than
about a 20°C. temperature range
• Polarity helps water stay liquid because molecules
so strongly attracted to each other
– It would dissolve very few other substances
• Polarity of water molecules can cause temporary
polarity in non-polar molecules; virtually everything
will dissolve to a small extent in water
In consequence, life could not exist anywhere
Water & Heat
H-bonds make it difficult to separate
molecules
H-bonds are constantly forming &
breaking
When temperature is stable, H bonds
form at the same rate that they break
Heat of Vaporization
When temperature increases:
H bonds break & stay broken
Individual molecules escape into air
= evaporation
Heat energy changes liquid H2O into
gaseous form
High boiling point (100°C)
When water cools:
H-bonds reform
H-bonds release heat energy as
temperature drops
Specific Heat Capacity
= energy required to raise given amount of
substance by 1°C
Water has high specific heat capacity:
At high temperatures, water absorbs heat as
H-bonds break
(can absorb a lot before temperature
measurably rises)
As water cools, heat released from formation
of H-bonds slows down cooling
Water’s high specific heat capacity:
• Helps regulate Earth’s climate by
buffering large changes in
temperature
• Helps moderate internal temperature
Density of Water
Water reaches max. density at 4°C
(becomes less dense at lower temps)
When temp decreases below 0°C:
Molecules don’t move enough to break H-bonds
so become locked
= ice
Water expands as freezes due
to hexagonal configuration of
molecules caused by H-bonds
Causes molecules to be
further apart than normal
Lower density causes ice to
“float” or form sheets at top of
water column
Insulates lakes & other bodies of
water in the winter
Cohesion and Adhesion
Cohesion:
– Water sticks to itself
– H-bonds cause attraction
between water molecules
Adhesion:
– Water sticks to other things
– Due to electrostatic forces of
molecules/H-bonds
e.g. transpiration in plants:
– Adhesion = water sticks to xylem
– Cohesion = holds water column
together
Surface Tension
How hard it is to break a liquid’s surface
Causes liquid to act as elastic sheet
Caused by H-bonds between water molecules
Liquid compresses to have smallest surface
area possible
e.g. water beading
Water as a Solvent
Ions & other polar molecules
dissolve readily in water
H2O molecules cluster
around ions / molecules in
sphere of hydration
Acids and Bases
Acid:
Dissociates in H2O
Releases H+ ions = proton donor
Concentration of protons
determines acidity of a
solution
Acids and Bases
Base:
Takes up H+ ions = proton accepter
Dissociates in H2O
Releases hydroxyl (OH-) ions
These bind to protons in solution, produce
water, & lower acidity of solution
Neutral:
Acid and base form H2O and salt
e.g. HCl + NaOH = H2O + NaCl
Strong Acids
Dissociate completely & irreversibly in
water
e.g. 100 HCl molecules in H2O becomes
100 H+ and 100 Cl(reaction occurs in one direction only)
Dramatically affect pH
Weak Acids
Dissociate partially in water
e.g. HAc  H+ + Ac(molecules of intact acid are in dynamic
equilibrium with dissociated ions )
Do not affect pH as much as strong acids
Important in body’s chemical buffer
systems
pH (potential of hydrogen)
Relative concentration of H+ ions in a solution
pH scale 0-14
Each pH unit is 10-fold change in [H+]
At pH = 7, [H+] = [OH-]
= neutral
Body’s internal environment
= pH 7.3-7.5
pH Scale
More on Acids and Bases
Strong acids and bases can cause severe
chemical burns
e.g. battery acid (pH ~ 1.0)
In high concentrations, can kill organisms
in an ecosystem
Acid Precipitation
Rain, snow, or fog with pH < 5.6
Caused by S oxides & N oxides in air
(from N-containing fertilizers & burning
of fossil fuels)
Oxides react with water vapour in air to
form H2SO4 & HNO3
Acid Precipitation in the US
Eastern US: pH 2-3 (rain)
Los Angeles: pH 1.7 (fog)
Effects on Terrestrial Systems
Has damaged / destroyed
forests in US, Canada,
Europe
Physical damage from
acid contact
Essential minerals in soil
washed away
Effects on Aquatic Systems
Kills aquatic life
Especially prevalent in spring:
Combo of snow melt & breeding season
Buffers
Buffers resist changes in pH by:
• Acting as acids (releasing H+) when pH 
• Acting as bases (binding H+) when pH 
Buffer Systems
It is imperative for cells to respond to
changes in pH
Changes disrupt cellular processes &
functioning of biological molecules
Buffer systems help resist large and
abrupt swings in pH
Bicarbonate Buffer System
Maintains blood pH (7.3 - 7.5)
If pH increases, carbonic acid releases H+
to neutralize excess OHH+ combines with OH- to form water
OH- + H2CO3 → HCO3- + H2O
When pH begins to drop, bicarbonate
consumes excess H+ to shift reaction
back towards acid
HCO3- + H+ → H2CO3
System is constantly buffering pH changes
A Final Word on Buffers
Buffer systems work within narrow range
When range is exceeded, extremely
severe effects
If blood pH drops to 7.0:
= respiratory acidosis, coma
If blood pH rises to 7.8:
= alkalosis, tetany