The chemical
nature of cells
Introduction
 At
the most basic level, the cell is a highly
organised assemblage of atoms
interacting with each other in a myriad of
chemical reactions.
 In order to understand the molecular
nature of the cell, we need to revisit some
chemical concepts.
Some definitions

Matter
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Atom
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Anything that takes up space and has mass.
An atom a fundamental piece of matter. Everything in the
universe (except energy) is made of matter, and, so, everything
in the universe is made of atoms.
An atom itself is made up of three tiny kinds of particles called
subatomic particles: protons, neutrons, and electrons.
Isotope

Different forms of atoms of the same element. They have the
same number of protons in their nuclei but a different number of
neutrons.
Some more definitions
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Molecule
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Element
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Formed when two or more atoms are joined
chemically.
An element is a substance consisting of only one type
of atom. e.g oxygen gas is O2 (pairs of oxygen
atoms combine to give oxygen gas)
Compound

A compound is a molecule that contains at least two
different types of molecules. e.g. water is H2O (two
hydrogens and one oxygen)
More about elements
 There
are 92 naturally occurring elements.
 Only 11 of these are found in organisms in
more than trace amounts, and four of
these make up more than 99% of
organisms by weight.
 The four most common elements in living
organisms are carbon (C), hydrogen (H),
oxygen (O) and nitrogen (N).
Memory Aid!
Most common elements:
CHONPS
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C for carbon
H for hydrogen
O for oxygen
N for nitrogen
P for phosphorous
S for sulfur
More about atoms

The red and green
circles in the centre
are the proton and
neutron, they both
make up the nucleus.
 The blue circle is the
electron and the black
ring shows its orbit
around the nucleus.
Even more about atoms – particularly
electrons and electron shells

The orbit or region of space around the nucleus in which
electrons are found is referred to as an electron shell.

There are rules as to the number of electrons that can be
held in each shell.

The arrangement of electrons within these shells is
called the electron configuration.

The electrons that occupy the outermost shell are called
valence electrons.

The chemical behaviour of an atom is a function of its
arrangement of electrons – in particular, the number of
valence electrons in its outermost electron shell.

Atoms are most stable when their outermost shell is full.
Making molecules
 Atoms
are most stable when their
outermost electron shell is full.
 In order to achieve this state, atoms tend
to combine with other atoms to form
molecules.
 There are two types of chemical bonds
that can hold the atoms within a molecule
together. These are ionic bonds or
covalent bonds.
Ionic Bonds

Atoms gain or lose
electrons in order to
increase stability.
 Atoms that lose electrons
have a positive charge
and are called cations.
 Atoms that gain electrons
have a negative charge
and are called anions.
 An ionic bond is an
electrical attraction
between two oppositely
charged atoms or groups
of atoms.
Example of ionic bonding
NaCl or sodium chloride

Sodium has 1 valence electron which it loses to become a cation
with a charge of 1+

Chlorine has 7 valence electrons and acquires 1 extra electron from
sodium in order to become an anion with a charge of 1-

The electrostatic attraction between positive sodium and negative
chloride (name changes when it forms an ion) holds the molecule
together.
Covalent Bonds

Atoms can share electrons
in order to increase their
stability
 Covalent bonding is where
the atoms share pairs of
outer shell or valence
electrons.
 Covalent bonding may be
single or multiple,
depending on the number
of pairs the atoms share.
 Sometimes in covalent
bonds, one atom attracts
the shared electrons more
strongly than the other,
resulting in a polar
covalent bond.
Example of covalent bonding
H2O or water

Hydrogen (H) has 1 valence
electron but needs a total of 2
in order to be a stable atom.

Oxygen has 6 valence
electrons but needs a total of 8
in order to be a stable atom.

By sharing electrons, each
hydrogen atom has two
valence electrons, thus filling
their outer orbits. Likewise,
oxygen now has 8 outer orbit
electrons.

This makes for a good
chemical bond and a stable
molecule.
Polar Covalent Bonds
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Polar covalent bonds are formed when one atom attracts
the shared electrons more strongly than the other.
Molecules that contain polar covalent bonds are referred
to as polar molecules and molecules that are formed by
covalent bonds but don’t have polar bonds are called
non-polar molecules.
The measure of the ability of an atom to attract electrons
in a bond is called electronegativity.
The region of the molecule which contains the atom with
the greatest electronegativity will have a slightly negative
charge compared to the rest of the molecule. The rest of
the molecule will have a slightly positive charge. This
opposite charge separated by a distance is called a
dipole.
Electronegativity

The higher the value on the Pauling Electronegativity
Scale, the stronger the atom’s electron attracting power.
Electronegativity Values (Pauling Scale)
C
H
O
N
P
S
2.5
2.1
3.4
3.0
2.2
2.6
Simple analogy for electronegativity
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Sharing electrons in
covalent bonds is like
trying to get a couple to
share the doona equally.
Someone is always going
to have more pulling
power!
One atom is always going
to “hog” the electrons.
In a water molecule
(H2O), the oxygen atom
will always attract the
shared electrons more
than the hydrogen atoms
do.
Interactions between molecules
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Intermolecular bonds are important in maintaining the 3D structure
of large biomolecules such as proteins and nucleic acids.
These interactions also allow a molecule to bind specifically but
transiently with another molecule.
Individual interactions are weak and constantly break and reform at
the physiological temperatures of organisms.
Multiple interactions act together to produce highly stable and
specific associations between parts of a large biomolecule and/or
between different molecules.
Four types of interactions are responsible for intermolecular bonds:
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Hydrogen bonds
Van der Waals interactions (transient dipoles)
Hydrophobic bonds
Ionic interactions
Hydrogen bonds
(dipole-dipole bonding)

Weak electrostatic
attraction between the
negative region (δ-) of
one polar molecule
and the positive
region (δ+) of another
polar molecule.
Van der Waals interactions
(transient dipoles)

At very short distances all
atoms and molecules show
a weak bonding interaction
due to their fluctuating
electrical charges.

Electrical charges fluctuate
due to the uneven
distribution of electrons as
they orbit the nucleus of an
atom. The larger the atoms
involved the greater the
fluctuation in charge.
Hydrophobic bonds
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Non-polar molecules
such as fats or oils will
aggregate together when
placed in a polar
substance such as water.
This aggregation is
referred to as a
hydrophobic bonds.
It is not a separate
bonding force but is due
to water molecules
excluding the non-polar
molecules, forcing them
to adhere to one another.
Red molecules are H2O
Ionic interactions

Ionic compounds are generally soluble in water due to
ionic interactions with water molecules.

Atoms in functional groups can donate or accept protons
(H+) forming ions that interact with other charged groups
on atoms on different molecules.
Molecule with NH2 group
attached has gained H+ so
becomes positive ion.
Molecule with COOH
group attached has lost H+
so becomes negative ion.
Inorganic and Organic Molecules

Both living and non-living things are made from
the same chemical elements but their is a
difference in the way that these elements are put
together to make larger molecules.
 Organic compounds contain carbon and
hydrogen (and sometimes other elements such
as oxygen and nitrogen). They are called
organic compounds as the first ones studied
were produced by or found in living organisms.
 All other compounds are called inorganic
compounds.
Biologically Important
Inorganic Molecules

Inorganic molecules important for living organisms include:
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nitrogen – present in all proteins and nucleic acids. Fixed from the
atmosphere by nitrogen-fixing bacteria
minerals – found in the cytosol and structural components of cells and
in the molecules of enzymes and vitamins. Important minerals include
phosphorous, potassium, calcium, magnesium, iron, sodium, iodine and
sulfur. Examples: phosphorous present in phospholipids of cell
membrane and in ATP, magnesium important component of chlorophyll,
iron an important component of haemoglobin
oxygen – most cells require oxygen to release usable energy from food
molecules
carbon dioxide – main source of carbon for the production of organic
molecules despite the fact that it only makes up 0.035% of atmosphere
by volume. Organic molecules (sugars) are eaten by animals and
carbon dioxide released back into atmosphere as an end-product of
cellular respiration.
water - chemical reactions in cells occur in a water environment
More about carbon
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All the chemicals of life on this planet, with the exception of water, are based on the
carbon atom.
It is the valency of carbon that allows it to form the base of all chemicals of life. The
carbon atom has four valence electrons, meaning it can form four stable covalent
bonds with other atoms.
Carbon atoms can also bond with other carbon atoms to form straight and branched
chain and ring structures of various sizes and complexity that form the backbone to
many biological molecules.
Carbon atoms can share more than one pair of electrons between two carbon atoms,
resulting in the formation of double and triple bonds.
Molecules containing only carbon and hydrogen atoms are known as hydrocarbons.
Hydrocarbons are non-polar and hence insoluble in water.
What creates the diversity and chemical properties of carbon based molecules is the
addition of other groups of atoms to the hydrocarbon backbone.
Groups of atoms that confer water solubility and chemical properties to the
hydrocarbon chain are known as functional groups.
These groups of atoms are more reactive than the hydrocarbon portion of
biomolecules and often contain highly electronegative atoms which convey a polarity
to their end of the molecule while the hydrocarbon chain is non-polar.
Water
Water covers about 75% of our planet’s surface and
makes up 70% to 90% of the cell content of living things.
 It has a number of unique properties that support life
on Earth.
 These properties are a direct result of the polar nature of
the water molecule

Why is water so special?
PROPERTY OF WATER
EXPLANATION OF PROPERTY
Cohesion
Liquid water is cohesive due to the constant forming and
reforming of hydrogen bonds that hold the molecules
together.
High specific heat
Heat must be absorbed to break hydrogen bonds before
the water molecules can move faster and the
temperature rises.
This means cell does not overheat despite the heat
energy produced by many reactions happening within the
cell.
Heat of vaporisation
A large amount of heat is required to break the hydrogen
bonds for liquid water to be converted to water vapour.
Solid is less dense than liquid
Hydrogen bonding creates a crystalline lattice in which
the water molecules are spaced further apart, reducing
the density of the solid.
Solvent
Water molecules are attracted to solutes that carry a
charge. The solute is said to have dissolved into
solution.
Both polar molecules and ionic compounds are easily
dissolved in water.
Water as a pH buffer
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pH is a measure of the hydrogen
ions in a solution and hence the
state of acidity or alkalinity of a
solution.
pH scale is 0 to 14.
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Less than 7 indicates higher
concentration of H+ (acidic
solution)
Greater than 7 indicates higher
concentration of OH- (alkaline
solution)
Pure water has pH of 7.0 and is a
neutral solution but water readily
ionizes or break ups to form H+
and OH- ions.
This allows the cellular fluids to
balance changes in pH as binding
of these ions to other substances
within the cell prevents severe
changes in the pH of a cell or fluid.
I’m a little water molecule!
(to the tune of “I’m a little teapot”)
I’m a little water molecule
They call me H2O
I’m not to good at sharing
My H more positive than my O
That’s why they call me polar
Cause my molecules attract each other
In between them, H bonds
I’m a BIG SELF LOVER!