COURSE READINESS ASSESSMENT
FOR
PHYSIOLOGY
ATOMS, MOLECULES
AND BONDS
Sections in this module
The atom and the periodic table
II. Electrons and electron shells
III. Ionic bonds
IV. Covalent bonds
I.
I. The atom and the periodic table
The atom

Matter is composed of atoms

Each atom is a single unit of a substance called an element
Element =
A substance that cannot be further broken down

All known elements are arranged in a chart called the
periodic table of the elements
© LeVanHan/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL
The structure of an atom

Almost all atoms consist of three different subatomic particles:
◦ Protons
+
◦ Neutrons
◦ Electrons
-

The center of an atom is called its nucleus. Protons and
neutrons occupy an atom’s nucleus.

The electrons of an atom orbit the nucleus
© Svdmolen/Jeanot/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL
Atom structure:
Subatomic particles

A proton has a positive charge and a mass of 1 Dalton
(“Dalton” is a standard unit of mass for these particles)

A neutron is uncharged and has a mass of 1 Dalton

An electron has a negative charge and a negligible mass
Subatomic
particle
Location
Charge
Mass
Proton (p or +)
In nucleus of atom
Positive
1 Dalton
Neutron (n or 0)
In nucleus of atom
Uncharged
1 Dalton
Electron (e or -)
Around nucleus
Negative
Negligible

Protons determine the identity of an atom (the element). They
also affect an atom’s charge and atomic mass.

Electrons affect an atom’s charge.

Neutrons affect atomic mass. Atoms of the same element with
different numbers of neutrons are called isotopes.
Isotopes

Example: Carbon has naturally occurring isotopes
Carbon-12 (12C), which is the most common form and has 6
neutrons.
Carbon-13 (13C) and carbon-14 (14C), which are less abundant.
13C has 7 neutrons and 14C has 8 neutrons.

Extra neutrons can make an atom unstable. Unstable atoms
can release energy in the form of radioactivity. 14C is an
example of a radioactive isotope.

Carbon dating measures the relative amount of 14C in a sample:
After an organism dies, the 14C in the organism decays over
time. Thus, the older the specimen, the less 14C will be present
in the specimen.

Radiation therapy uses radioactive isotopes to treat cancer:
Radioactivity damages DNA and other molecules in cancer cells,
causing them to die. This type of therapy can be very targeted.
For example, small capsules (“seeds”) containing radioactive
isotopes can be placed near cancer tumors to avoid damaging
healthy tissue.

Radioactive isotopes can be used as medical markers:
A patient can ingest a
radioactive sugar solution.
The radiation can be
detected by a PET scan.
This is a PET scan of a brain
cancer patient, showing the
location of a cancer tumor.
© LeVanHan/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL
Reading the periodic table of the elements
Atomic number
Atomic mass
5B
10.81
Symbol for the element
1st letter is capitalized
2nd letter, if present, lower
case

The atomic number is unique to an element and is an integer

The atomic mass (or mass number) can be displayed with
a decimal or rounded to a whole number.
Periodic table: Atomic number

Each element in the universe has a unique number of protons

If an atom is uncharged, it will have the the same number of
protons and electrons

Thus, the atomic number represents:
◦ The number of protons present in an atom
◦ The number of electrons also present in an atom,
assuming that the atom is uncharged
Periodic table: Atomic mass

Recall that an atom’s nucleus contains its protons and neutrons

Each proton and neutron has a mass of 1Dalton

The atomic mass (or mass number) represents the sum of
the number of protons and neutrons in an atom

Atoms of the same element can sometimes differ in the
number of neutrons. This is why the atomic mass is sometimes
displayed with a decimal.
5B
10.81

Example: The element boron (B)
The atomic number tells us that boron atoms have 5 protons
and 5 electrons, assuming that the atom is uncharged.
The atomic mass tells us the sum of the protons and neutrons
of an atom. Boron’s atomic mass is a value between 10 and 11
Daltons, telling us that boron atoms can have 5 or 6 neutrons.
Concept check
Determine the number of protons, neutrons, and electrons
present in the elements shown:
7N
14
11 Na
23
“Shorthand” way to represent elements
The atomic number and atomic mass of an element can be
reversed from the way they appear in the periodic table. This is
a common shorthand way to describe an element:
10.81
5
B
Concept check
Write the shorthand for the following elements (you can round
the mass to the closest whole number).
Then, determine the number of protons, neutrons and electrons
for each element.
a)Fluorine
(F) has an atomic number 9 and an atomic mass of
18.998 Daltons
a)Hydrogen
(H) has an atomic number 1 and an atomic mass of
1.0079 Daltons
Concept check answers
19
9
1
1
F
9 protons, 9 electrons, and approximately 10 neutrons
H
1 proton, 1 electron, and 0 neutrons
(Hydrogen atoms are unusual – they generally lack neutrons!)
II. Electrons and electron shells
An atom’s electrons can be diagrammed

The electrons of an atom can be drawn in circular orbits
around the nucleus called “shells”, as shown for carbon:

Since these diagrams are meant to show electrons, the
individual protons and neutrons in the nucleus are not
shown
© Pumbaa/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL
Drawing “electron shell” diagrams

Electron shells are drawn as circles around the atom’s nucleus

Electrons in a shell are drawn as dots:
◦ The first electron shell is closest to the nucleus. It holds a
maximum of 2 electrons.
◦ If an atom has more than 2 electrons, additional shells are
drawn outside of the first one.
Outer shells hold a maximum of 8 electrons. *
* This rule changes for elements with higher atomic numbers, so
we can limit ourselves to elements with an atomic number of
20 or lower.

Carbon has an atomic number of 6. Thus, a carbon atom has
a total of 6 electrons, which are shown as gray dots in the
diagram below:

The nucleus of the carbon atom is shown as a green circle

2 electrons occupy the first electron shell, which is closest to
the nucleus. This first shell is completely full.

The remaining 4 electrons are placed in an outer shell, one
by one.

Fill the outer shells with electrons one by one, moving
clockwise or counterclockwise.

Since an outer shell holds 8 electrons, you can go two
rounds around each outer shell. The purpose of doing this is
to draw electrons in pairs.

For example, fluorine (F) has an atomic number of 9, meaning
that a fluorine atom has 9 electrons. Here is a diagram of
the electrons in this atom:
F
Valence electrons

The outermost shell drawn in an electron diagram is called
the valence shell

Thus, the electrons in this shell are called valence electrons
Examples:
Carbon has 4 valence electrons.
Fluorine has 7 valence electrons.
F
Building a molecule

A molecule consists of two or more atoms held together by
chemical bonds.
Examples:
Oxygen gas (O2) consists of two oxygen atoms.
Sodium chloride (NaCl) consists of one sodium and one chlorine
atom.

When a molecule is made of two or more elements, it can be
called a compound.
Example: NaCl is a compound, but O2 is not!
Building a molecule: Electrons!

Atoms with a full valence shell do not form bonds with other
atoms. This is because all of their valence electrons are paired.
Paired electrons are happy electrons! 

Atoms with unpaired valence electrons can form chemical bonds.
Unpaired valence electrons are unhappy because they want a
partner! 

Each unpaired valence electron can form a bond.
Examples:
Carbon has 4 unpaired valence electrons.
It can form 4 bonds.
Fluorine has 1 unpaired valence electron.
It can form 1 bond.
F
III. Ionic bonds
The ionic bond

An ionic bond is formed when one atom donates an unpaired
valence electron to another atom
◦ The donor atom (that gives away the electron) has a valence
shell that is almost empty
◦ The recipient atom (that gets the electron) has a valence
shell that is almost full

When an atom gains or loses electrons, it becomes an ion,
which is an atom with a charge

An ionic bond is the attraction between ions of opposite
charges
Atoms that form ionic bonds

Typically, metals and non-metals form ionic bonds with each other.
These are elements found at opposite ends of the periodic table.
Non-metals
Metals

In a molecule of sodium chloride (NaCl), the sodium and
chlorine atom are held together by an ionic bond
◦ The sodium atom (atomic number = 11) is the electron
donor:
Its valence shell has one unpaired electron,
making the shell almost empty
◦ The chlorine atom (atomic number = 17) is the electron
recipient:
Its valence shell has 7 electrons,
making the shell almost full
© Tra/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL

The sodium atom will donate its unpaired valence electron to
chlorine:
© Tra/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL

Sodium’s outermost shell is empty, so that the valence shell is now
the second one, which is full
Sodium lost an electron. Now it has 11 protons and 10 electrons,
making it a sodium ion (Na+) with a +1 or “+” charge

Chlorine’s outermost shell is now full.
Chlorine gained an electron. Now it has 17 protons and 18
electrons, making it a chloride ion (Cl-) with a -1 or “-” charge

Since the sodium ion and chloride ion have opposite charges,
they will be attracted to each other. This is the ionic bond that
holds them together.

Compounds like these are called salts. Salts like sodium chloride
dissolve in water because the water breaks the ionic bond
between the sodium and chloride ions.
The breaking apart of the sodium and chloride ions is called
dissociation.
IV. Covalent bonds
The covalent bond

A covalent bond when two atoms share unpaired valence
electrons with each other

Many molecules that are important for life are held together
by covalent bonds
Examples:
Smaller molecules like water (H2O) and oxygen gas (O2)
Larger molecules like DNA and proteins

In biology, covalent bonds are considered stronger than ionic
bonds
Atoms that form covalent bonds

Typically, non-metals form covalent bonds with each other.
Many molecules important for life contain the element carbon (C).
Non-metals

Let’s use hydrogen gas as an example. Hydrogen gas (H2) consists
of two hydrogen atoms.

Hydrogen has an atomic number of 1, so each atom has one
electron:

H
H


Both hydrogen atoms are unhappy because their valence
electrons are unpaired!

To remedy this, the hydrogen atoms will share their unpaired
electrons with each other, creating an electron pair. This shared
electron pair is the covalent bond.
H

H
Notice that the valence shells are overlapping and the shared
electrons (one from each hydrogen atom) are in the overlap
Showing covalent bonds

Covalent bonds can be shown several different ways:
H
H
Electron shell diagram
H-H
Bond as solid line
H:H
Lewis dot structure
Space-filling model
Multiple covalent bonds are possible
between atoms

Double covalent bond: Sharing two pairs of electrons
O
O
O
O
O C O

Triple covalent bond: Sharing three pairs electrons
C O
Nonpolar covalent bonds

A nonpolar covalent bond is a covalent bond where the pair
of electrons is equally shared between the two atoms

Nonpolar covalent bonds occur between
◦ Two of the same non-metal atom
Examples:
Hydrogen gas (H2), oxygen gas (O2)
◦ Carbon and hydrogen

Thus, a molecule that contains only nonpolar covalent bonds
is a nonpolar molecule
Polar covalent bonds

A polar covalent bond is a covalent bond where the pair of
electrons is not quite equally shared between the two atoms

Polar covalent bonds occur between two different non-metal
atoms, where one atom has a stronger “pull” on shared
electrons * than the other atom.
Examples:
Hydrogen and oxygen in a molecule of water (H2O)
Hydrogen and nitrogen in a molecule of ammonia (NH3)
* An atom’s “pull” on electrons is called electronegativity.

In a water molecule, the oxygen atom has a greater “pull” on
shared electrons than either hydrogen. Thus, each covalent bond
in a water molecule is a polar covalent bond.
δ-
δ+
H
O
H
δ+

Polar covalent bonds give a molecule slight charges. In water:
◦ The shared electrons spend more time with oxygen, giving this
atom a slight negative charge (δ-, δ= Greek letter delta)
◦ The shared electrons spend less time with hydrogen, leaving it
with a slight positive charge (δ+)

Polar molecules are molecules with these slight charges!
© Maksim/ Wikimedia Commons / CC-BY-SA-3.0 / GFDL
Some molecules contain both ionic and
covalent bonds

For example, sodium bicarbonate
(sodium hydrogen carbonate, baking soda):
NaHCO3
This molecule dissociates in water to form three ions:
NaHCO3  Na+ + H+ + CO32sodium ion
hydrogen ion
carbonate ion
The carbon and oxygen atoms in the carbonate ion are held
together by covalent bonds.
The sodium and hydrogen are held together with the carbonate
ion by ionic bonds.
Summary of Atoms, Molecules, and Bonds

Atoms are single units of elements and consist of three types
of subatomic particles: protons, neutrons, and electrons

The atomic number is unique to an element and specifies
the number of protons in an atom. It also specifies the
number of electrons in an uncharged atom. The atomic
mass is the sum of an atom’s protons and neutrons.

Unpaired valence electrons help atoms form chemical bonds
◦ An ionic bond forms when one atom donates an
electron to another atom
◦ A covalent bond forms when two atoms share electrons.
Covalent bonds can be nonpolar or polar.