THE CHEMISTRY OF LIFE
CHAPTER 2
ATOMS
• Any substance that has mass and occupies
space is called matter.
• Matter is composed of small particles called
atoms.
• An atom is the smallest particle that still
retains the chemical properties of a
substance.
MASS VS. WEIGHT
• Mass – the amount of a substance.
• Weight – the force gravity exerts on a
substance.
• An objects mass is the same on Earth or the
moon.
• An objects weight will be greater on Earth
because the Earth’s gravitational force is
greater.
BASIC ATOMIC STRUCTURE
• Core nucleus containing protons (+) and neutrons.
• Cloud of electrons (-) surrounding the nucleus.
• Protons, neutrons, and electrons are called
subatomic particles.
IONS
• Neutral atom (no charge) one orbiting electron (-)
for every proton (+) in the nucleus.
• Atoms where the protons and electrons do not
balance are called ions.
• Ions may be positively or negatively charged.
ELECTRONS
• Protons and neutrons contribute most of the
mass of an atom – electrons have very little
mass.
• Electrons determine chemical behavior of
atoms because they come into contact
with electrons of other atoms.
ELECTRONS CARRY ENERGY
• It takes work to hold electrons in orbit.
• Electrons have energy of position – potential
energy.
ELECTRONS CARRY ENERGY
• An electron can be transferred from one
atom to another.
• Loss of an electron is called oxidation.
• Gain of an electron is called reduction.
ELECTRONS CARRY ENERGY
• We use simplistic models to visualize atoms.
• The volume of space around a nucleus
where an electron is most likely to be found
is called the orbital of that electron.
ELECTRONS CARRY ENERGY
• Each energy level has a specific number of
orbitals, which can hold up to 2 electrons.
• The first energy level has 1 orbital.
• The outer energy levels have 4 orbitals each.
• Atoms with incomplete orbitals tend to be
more reactive.
ATOMIC NUMBER
• The number of protons in the nucleus is the
atomic number.
• Atomic number of sodium is 11.
ATOMIC MASS
• Neutrons and protons are similar in mass.
• Each weighs about 1 dalton.
• The number of neutrons plus the number of
protons is called the atomic mass.
• Atomic mass of sodium is about 22 (22.989)
daltons.
• Electrons have negligible mass.
ELEMENTS
• Atoms with the same atomic number (# of
protons) have the same chemical properties
and belong to the same element.
ISOTOPES
• Atoms that have the same number of
protons, but different number of neutrons
are called isotopes.
• Same atomic number, different mass.
ISOTOPES
• Example – Carbon: 6 protons, 6 electrons.
• Usually has 6 neutrons = Carbon-12
• Sometimes has 7 or 8 neutrons.
• Carbon-14 is unstable; nucleus tends to break up
(radioactive decay).
DATING FOSSILS
• Fossils are created
when remains,
footprints, or other
traces become buried
in sand or sediment.
• Over time, calcium is
mineralized as the
sediment turns to rock.
• A fossil is any record of
prehistoric life –
generally over 10,000
years old.
DATING FOSSILS
• Scientists date the rocks containing fossils to
get an idea of how old the fossils are.
• Rocks can be dated by measuring the
degree of radioactive decay of certain
radioactive atoms among rock-forming
minerals.
DATING FOSSILS
• A radioactive element contains so many
neutrons and protons that it is unstable and
likely to fly apart.
• The rate of decay for a particular element is
constant and so it can be used to date
rocks.
DATING FOSSILS
• Carbon-14 (14C) radioisotopic dating is often
used to date fossils less than 50,000 years
old.
• A certain proportion of naturally occurring
carbon is carbon-14 rather than carbon-12.
• After an organism dies, its 14C decays over
time.
DATING FOSSILS
• It takes 5,600 years for half of the 14C to be
converted to 14N.
• This is called the half-life of the isotope.
• Constant
• A sample with ¼ the original proportion of
14C would be 11,200 years old (2 half-lives).
• 40K decay into 40Ar has a half-life of 1.3 billion
years.
MOLECULES
• A molecule is a group of atoms held
together by energy.
• The force holding two atoms together is
called a chemical bond.
• Ionic bond
• Covalent bond
• Hydrogen bond
IONIC BONDS
• Ionic bonds occur when atoms are
attracted to each other by opposite
electrical charges.
• Strong (not as strong as covalent bonds)
• Not directional
• Form crystals
COVALENT BONDS
• Covalent bonds are formed when two atoms
share electrons.
• Atoms seek to fill outermost sphere of orbiting
electrons.
• Strongest type of bond.
• Very directional – bonds form between two specific
atoms rather than a generalized attraction of one
atom for its neighbors.
HYDROGEN BONDS
• When covalent bonds form, one atom may
attract the electrons more strongly than the
other(s).
• Shared electrons spend more time around
the “stronger” atom giving it a slightly
negative charge, while the “weaker”
atom(s) have a slightly positive charge.
HYDROGEN BONDS
• These charges are not as strong as those on
an ion.
• The molecules end up with a positive end
and a negative end and are said to be
polar.
• A hydrogen bond occurs between the
positive end of one polar molecule and the
negative end of another polar molecule.
HYDROGEN BONDS
• Hydrogen bonds are weak.
• As a result of weakness, they are highly
directional – polar molecules must be very close
for the weak attraction to be effective.
• Act like Velcro, forming a tight bond as a result
of many weak interactions.
HYDROGEN BONDS
• Hydrogen bonds are important in biology
because they stabilize the shapes of
biological molecules by causing certain
parts to be attracted to other parts.
• Protein shape is dictated by hydrogen
bonds.
HYDROGEN BONDS & WATER
• Water is essential for life.
• H2O – An oxygen covalently bonded to two
hydrogens.
• Water is polar – can form hydrogen bonds.
• Weak, short-lived hydrogen bonds form
between water molecules.
• Cumulative effect of transient bonds is
important.
HEAT STORAGE
• Temperature of a substance is a measure of
how fast molecules are moving.
• Due to many hydrogen bonds, a large input
of thermal energy is necessary to disrupt the
organization of water and raise its
temperature.
• Water heats slowly, holds temperature
longer.
• Contributes to internal body temp
regulation.
ICE FORMATION
• When temperature is
low enough, few
hydrogen bonds
break.
• Water molecules form
a crystal-like structure
– ice.
• Ice is less dense than
water – molecules
are spaced out by
hydrogen bonds.
HIGH HEAT OF VAPORIZATION
• When temperature is high enough, many
hydrogen bonds break and liquid is
changed to vapor.
• A great deal of heat energy is required to
do this.
COHESION
• Polar water molecules are attracted to other
polar molecules.
• Attraction between water molecules is called
cohesion.
ADHESION
• When the other polar molecule is not water,
the attraction is called adhesion.
• Capillary action is a result of adhesion.
• Adhesion is why things get wet when dipped
in water.
HIGH POLARITY
• Water molecules tend to form the maximum
number of hydrogen bonds possible.
• Polar molecules are called hydrophilic or
water-loving molecules.
HIGH POLARITY
• When salt
dissolves,
molecules break
off and are
surrounded by
water molecules.
• Polar molecules
that dissolve in
water are said to
be soluble.
HIGH POLARITY
• Nonpolar molecules like oil do not form
hydrogen bonds and are not water-soluble.
• When nonpolar substances are put in water,
the water forms hydrogen bonds with each
other – leaving out the nonpolar molecules.
• Nonpolar substances are said to be
hydrophobic or water-fearing.
WATER IONIZES
• Covalent bonds within a water molecule
sometimes break spontaneously.
H2O
OH–
hydroxide
ion
+
H+
hydrogen
ion
• This process of spontaneous ion formation is
called ionization.
• It is not common because of the strength of
covalent bonds.
PH
• A convenient way to express the hydrogen
ion concentration of a solution.
pH = _ log [H+]
• The pH scale is logarithmic
• A difference of one unit represents a ten-fold
change in H+ concentration.
H+ Ion
Concentration
100
Examples of
Solutions
pH Value
Hydrochloric acid
0
10-1
1
10-2
2
Stomach acid
Lemon juice
10-3
3
Vinegar, cola, beer
10-4
4
Tomatoes
10-5
5
Black coffee
Normal rain water
10-6
6
Urine
Saliva
10-7
7
10-8
8
Pure water
Blood
Seawater
10-9
9
Baking soda
10-10
10
10-11
11
Great Salt Lake
Milk of magnesia
Household ammonia
10-12
12
Household bleach
10-13
13
Oven cleaner
10-14
14
Sodium hydroxide
H+
H+
+
+
H+ H H
+
OH– H
OH–
OH–
H+
OH–
H+
OH–
H+ H+
OH–
OH– H+
OH–
–
OH– OH
H+
OH–
PH
• Pure water has a pH of 7.
• There are equal amounts of [H+] relative to [OH–].
• Acid—any substance that dissociates in
water and increases the [H+].
• acidic solutions have pH values below 7.
• Base—any substance that combines with
[H+] when dissolved in water.
• basic solutions have pH values above 7.
PH
• The pH in most living cells and their
environments is fairly close to 7.
• Proteins involved in metabolism are sensitive
to any pH changes.
• Acids and bases are routinely
encountered by living organisms.
• From metabolic activities (i.e., chemical
reactions).
• From dietary intake and processing.
• Organisms use buffers to minimize pH
disturbances.
WATER IONIZES
• Buffer—a chemical substance that takes up
or releases hydrogen ions.
• Buffers don’t remove the acid or the base
affecting pH but minimize their effect on it.
• Most buffers are pairs of substances, one an acid
and one a base.
BUFFERS
• Living cells have a pH near 7.
• A constant pH must be maintained for
metabolic activities to work properly
(homeostasis).
• Buffers act as a reservoir for hydrogen ions,
donating them or taking them as needed.
BUFFERS
• Key buffering pair in human blood: carbonic
acid and bicarbonate (a base).