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The Chemistry of Life
Chapter 3
 Organisms are chemical machines
– one must know chemistry in order to understand
 Any substance in the universe that has
mass and occupies space is comprised of
– Anything that takes up space and has mass
– Can exist as a liquid, solid, or gas
– All matter is made up of atoms
 Atom = smallest particle a substance can be
divided into that can retain its properties
 All atoms have the same structure
– at the core is a dense nucleus comprised of two
types of subatomic particles
 protons (positively charged)
 neutrons (no associated charge)
– orbiting the nucleus is a cloud of another
subatomic particles
 electrons (negatively charged)
 An atom can be characterized by:
– Atomic Number
 the number of protons in the nucleus
 atoms with the same atomic number exhibit the same
chemical properties and are considered to belong to the
same element
 e.g. Carbon = C: atomic number = 6
– Atomic Mass (mass number)
 the number of protons plus neutrons in the nucleus
 electrons have negligible mass (1/1840 dalton)
 e.g. C: atomic mass = 12.011
Atomic Number and Mass
Note: Hydrogen is unique in that it has 1 proton, but 0 neutrons!
Atomic Symbol
Atomic Mass =
Number of Protons +
Number of Neutrons
Atomic Number = The Number of Protons
in the Nucleus
 Electrons determine the chemical behavior
of atoms
– These subatomic components are the parts of
the atom that come close enough to each other
in nature to interact
 Same charges repel each other
 Opposite charges attract each other
 Electrons can be shared
 Electrons are associated with energy
– Potential energy: energy of position
 e.g. Rollercoaster at top of peak
 e.g. As electrons move away from core, they
increase potential energy
– The field of energy around an atom is arranged
as levels called electron shells
– Orbitals are the location electrons are most
likely to be found within this volume of space
Electron Arrangement of Atoms
1 electron
6 electrons
7 electrons
8 electrons
 Electron shells have specific numbers of
orbitals that may be filled with electrons
– atoms that have incomplete electron orbitals
tend to be more reactive
– atoms will lose, gain, or share electrons in order
to fill completely their outermost electron shell
– these actions are the basis of chemical bonding
 How many electron shells?
 How many electrons in first
 How many electrons in
second shell?
 How many electrons in
third shell?
 How many total electrons?
 What is Atomic Number?
 What is Atomic Mass?
Sodium atom = Na
 Basic building block of
 92 naturally occurring
 Only 6 elements make up
most of the body weight of
– C Carbon
– H Hydrogen
– N Nitrogen
– O Oxygen
– P Phosphorus
– S Sulfur
 Ions – atoms that have gained or lost one or
more electrons
– Gaining an electron makes gives a negative
– Losing an electron gives a positive charge
 For Example:
– Sodium ion has 11 protons, 10 electrons
– Does this sodium ion have a positive or negative
– A negative ion could also form if an extra electron
were added
Sodium Ion
 Isotopes – atoms that have the same number of
protons but different numbers of neutrons
– most elements in nature exist as mixtures of different
 C-12: 6 protons, 6 neutrons, 6 electrons
 C-14: 6 protons, 8 neutrons, 6 electrons
 Some isotopes are unstable and break up into
particles with lower atomic numbers
– this process is known as radioactive decay
– Radioactive isotopes can be used in nuclear medicine
and for dating fossils
Figure 3.5 Isotopes of the element carbon
A molecule is a group of atoms held together by energy
Atoms can interact in 3 ways:
e.g. water (H2O), sodium chloride (NaCl), oxygen (O2)
The energy holding two atoms together is called a chemical bond
1. Share one or more electrons
2. Accept extra electrons
3. Donate electrons to another atom
There are 3 principal types of chemical bonds
1. Ionic
2. Covalent
3. Hydrogen
 Ionic bonds involve the
attraction of opposite
electrical charges
• Transfer of electrons from
one atom to another
 Molecules comprised of
these bonds are often
most stable as crystals
 Remember:
– An ION is an atom that with
a charge
Fig. 3.8(a)The formation of the
ionic bond in table salt
 – Covalent bonds form between two atoms
when they share electrons
– The number of electrons shared varies
depending on how many the atom needs to fill
its outermost electron shell
– Covalent bonds are stronger than ionic bonds
 A covalent bond
– Each Hydrogen has 1
electron in shell
– Sharing 2 electrons fills the
shell, increasing stability
 A double covalent bond
– Sharing 2 pairs of electrons
– Oxygen has total of 8
electrons (2 in inner shell, 6
in outer shell)
– Sharing 2 more fills its outer
shell, increasing stability
 Hydrogen bonds form from covalent bonds
created by an atom’s electronegativity
– Create partial charges in atoms that are
unequally sharing electrons
– Are weak bonds
 Electronegativity is the tendency of one
atom’s nucleus to better attract the shared
electrons from another nucleus
Hydrogen bonding of Water
Water molecules contain two covalent bonds
 Hydrogen bonds form in association with polar molecules
– each atom with a partial charge acts like a magnet to bond weakly
to another polar atom with an opposite charge
– the additive effects of many hydrogen bonding interactions can add
collective strength to the bonds
Figure 3.10 Hydrogen bonding water molecules
Hydrogen Bonds Give Water Unique
 Water is essential for life
– The chemistry of life is water chemistry!
 Water is a polar molecule
– The partial charges of hydrogen bonds creates
– Water can form hydrogen bonds
– Hydrogen bonding confers on water many
different special properties
Hydrogen Bonds Give Water Unique
 Heat Storage
– Water temperature changes slowly and holds temperature well
 This is due to the large number of H bonds many water molecules will
form with each other, it takes a lot of energy to break them (and raise
 Ice Formation
– Few hydrogen bonds break at low temperatures
 Water becomes less dense as it freezes because hydrogen bonds
stabilize and hold water molecules farther apart
 High Heat of Vaporization
– At high temperatures, hydrogen bonds can be broken
 water requires tremendous energy to vaporize because of all the
hydrogen bonds that must be broken
Hydrogen Bonds Give Water Unique
 Water molecules are sticky
– Cohesion – when one polar
water molecule is attracted
to another polar water
– Adhesion – when OTHER
polar molecules water are
attracted to a water molecule
Figure 3.12
Hydrogen Bonds Give Water Unique
 Water is highly polar
– in solution, water molecules tend to form the
maximum number of hydrogen bonds
 Hydrophilic molecules are attracted to water and
dissolve easily in it
– these molecules are also polar and can form hydrogen
 Hydrophobic molecules are repelled by water and
do not dissolve
– these molecules are non-polar and do not form hydrogen
Acids and Bases
 When water ionizes, it releases an equal number
of hydrogen ions (H+) and hydroxide ions (OH-).
Water Ionizes
 The amount of ionized hydrogen from water in a
solution can be measured as pH
pH = -log[H+]
 The pH scale is logarithmic, which means that a pH
scale difference of 1 unit actually represents a 10-fold
change in hydrogen ion concentration
– e.g. pH of 4 has 10x greater H+ concentration than pH of 5
– e.g. pH of 4 has 100x greater H+ concentration than pH of 6
 pH difference of 2; 10 x 10 = 100
The pH scale
Water Ionizes
 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
Water Ionizes
 The pH in most living cells and their environments
is fairly close to 7
– proteins involved in metabolism are sensitive to any pH
 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
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
Buffer Example
 Carbonic acid and bicarbonate in human blood
 Interact in a pair of reversible reactions
– CO2 + H2O  H2CO3
– H2CO3 H+ + HCO3-
 If H+ is added, HCO3- can pick up H+ added to
form H2CO3
 If H+ is removed, it disassociates to release
more H+ into blood