OBJECTIVE INTRODUCTION
At the beginning of the lesson, the class will do a Think-Pair-Share to discuss
the objective.
CLASS ACTIVITY
1. Write these words on the board:
o Ionic Bonds
o Covalent Bonds
2. Show the animated video – The Chemical Bonds Song. The link is
provided.
3. Explain to students they will be taking notes during the video.
STUDENT ACTIVITY
1. Have students break up into small groups.
2. In each group, have students discuss and compare their notes with each
member of the group.
3. Have students try to explain Ionic Bonding.
4. Have students try to explain Covalent Bonding.
5. Play the song again to allow students an opportunity to add to their
explanations of both bond types.
The teacher will help to clear any misconceptions about chemical bonding. A
major misconception is students may think they can acquire dichotomous
classification of bonds readily, but fail to understand that bonding is primarily
an electrical phenomenon. This is an area where helping students to
appreciate that our descriptions and diagrams are just models will make
learning easier.
Estimated Class Time for the Engagement: 20-30 minutes
EXPLORATION
This student-centered station lab is set up so students can begin to explore
chemical bonding. Four of the stations are considered input stations where
students are learning new information about chemical bonding, and four of the
stations are output stations where students will be demonstrating their
mastery of the input stations. Each of the stations is differentiated to
challenge students using a different learning style. You can read more
about how I set up the station labs here.
EXPLORE IT!
Students will be working in pairs to identify the types of chemical
bonds. Students will follow the directions on the task cards and use
manipulatives to show how valence electrons in atoms combine to form
chemical bonds. Students will be demonstrating the bonds of H2O and NaCl.
Throughout the process, task cards will continue to assist students to
understand the differences between ionic and covalent bonds.
WATCH IT!
At this station, students will be watching a nine-minute video explaining the
differences between covalent and ionic bonding. Students will then answer
questions related to the video and record their answers on their lab station
sheet. For example, what are the differences between covalent and ionic
bonds? Describe the octet rule. What type of bonds would be in CO2?
RESEARCH IT!
The research station will allow students to explore an interactive webpage that
has students take a scientific approach to understanding chemical bonds.
Students will be instructed to complete a few tasks and record answers on
their lab sheets.
READ IT!
This station will provide students with a one page reading about why are
bonds important. In the reading, students will discover many forms of
chemical bonds, what happens when chemical bonds break, and examples of
energy production. There are 4 follow-up questions that the students will
answer to show reading comprehension of the subject.
ASSESS IT!
The assess it station is where students will go to prove mastery over the
concepts they learned in the lab. The questions are set up in a standardized
format with multiple choice answers. Some questions include: What type of
bonds are shown in the pictures? Which two elements are most likely to bond
with each other? Fill in the blank: ____ bonds share valence electrons with
other atoms.
WRITE IT!
Students who can answer open-ended questions about the lab truly
understand the concepts that are being taught. At this station, the students
will be answering three task cards to describe the valence electrons in
nitrogen and how it could bond to other atoms. Students will also answer in
their own words how are ionic bonds are different from covalent bonds, and
why atoms are attracted to each other in an ionic bond.
ILLUSTRATE IT!
Your visual students will love this station. Students will draw two diagrams to
show how atoms bond with each other to create chemical molecules.
Students will identify the diagrams as either a covalent bond or ionic bond.
ORGANIZE IT!
The organize it station allows your students to manipulate cards in the correct
column by whether the information is describing that of covalent bonds or
ionic bonds.
Estimated Class Time for the Exploration: 1-2, 45 minute class periods
EXPLANATION
The explanation activities will become much more engaging for the class once
they have completed the exploration station lab. During the explanation
piece, the teacher will be clearing up any misconceptions about chemical
bonding with an interactive PowerPoint, anchor charts, and interactive
notebook activities. The chemical bonding lesson includes a PowerPoint with
activities scattered throughout to keep the students engaged.
The students will also be interacting with their journals using INB templates for
chemical bonding. Each INB activity is designed to help students
compartmentalize information for a greater understanding of the concept. The
chemical bonding INB template allows students to focus their notes on
identifying the differences between covalent and ionic bonds.
Estimated Class Time for the Exploration: 2-3, 45 minute class periods
ELABORATION
The elaboration section of the 5E method of instruction is intended to give
students choice on how they can prove mastery of the concept. When
students are given choice the ‘buy-in’ is much greater than when the teacher
tells them the project they will have to create. The elaboration project will
allow students to create a presentation to teach about chemical bonding.
Estimated Class Time for the Elaboration: 2-3, 45 minute class periods (can
also be used as an at-home project)
EVALUATION
The final piece of the 5E model is to evaluate student
comprehension. Included in every 5E lesson is a homework assignment,
assessment, and modified assessment. Research has shown that homework
needs to be meaningful and applicable to real-world activities in order to be
effective. When possible, I like to give open-ended assessments to truly
gauge the student’s comprehension.
Chemical Bonding: The Nature of the
Chemical Bond
Life on Earth depends on water – we need water to drink, bathe, cool ourselves off on a hot
summer day (Figure 1). In fact, evidence suggests that life on Earth began in the water, more
specifically in the ocean, which has a combination of water and salts, most prominently common
table salt – sodium chloride. But where do water and these common salts appear on the great
organizer of the elements, the periodic table? Well they, and millions of other substances, are not
found on the most famous of all chemistry references: the periodic table. Why not? The answer
is a simple one.
Figure 1:
Life on Earth depends on water, not only for key biological functions but also for pleasure. For example,
this relaxing oasis on the Mediterranean Sea, Cala Tío Ximo beach in Benidorm, Spain.image © Diego
Delso
The periodic table organizes the 118 currently recognized chemical elements, but water and
sodium chloride are not elements. Rather, both are substances that are made up of a combination
of elements in a fixed ratio. Such fixed ratio combinations of those 118 elements are known
as compounds.
In its chemical reactions and physical interactions, sodium chloride doesn’t act like
the elements that make it up (sodium and chlorine); rather, it acts as a completely different and
unique substance. That’s a good thing since chlorine is a poisonous gas that has been used as a
chemical weapon, and sodium is a highly reactive metal that is mildly explosive with water. So
what allows sodium chloride to act in an entirely different way? The answer is that within
table salt, sodium and chlorine are joined together by a chemical bond that creates a
unique compound, very different from the individual elements that comprise it.
The chemical bond can be thought of as a force that holds the atoms of various elements together
in such compounds. It opens up the possibility of millions and millions of combinations of the
elements, and the creation of millions and millions of new compounds. In short, the existence of
the chemical bonds accounts for the richness of chemistry that reaches far beyond just those 118
building blocks.
The history of the chemical bond
When discussing the history of chemistry it’s always dangerous to point to the specific origin of
an idea, since by its very definition, the scientific process relies upon the gradual refinement of
ideas that came before. However, as is the case with a number of such ideas, one can point to
certain seminal moments, and in the case of chemical bonding, a famous early 18th century
publication provides one such moment.
In his 1704 publication Opticks, Sir Isaac Newton makes mention of a force that points to the
modern idea of the chemical bond. In Query 31 of the book, Newton describes ‘forces’ – other
than those of magnetism and gravity – that allow ‘particles’ to interact.
In 1718, while translating Opticks into his native language, French chemist Étienne François
Geoffroy created an Affinity Table. In this fascinating first look at the likelihood of certain
interactions, Geoffroy tabulated the relative affinity that various substances had for other
substances, and therefore described the strength of the interactions between those substances.
While Newton and Geoffroy’s work predated our modern understanding
of elements and compounds, their work provided insight into the nature of chemical interactions.
However, it was over 100 years before the concept of the combining power of elements was
understood in a more modern sense. In a paper in the journal Philosophical Transactions entitled
“On a new series of organic bodies containing metals” (Frankland, 1852), Edward Frankland
describes the “"combining power of elements,” a concept now known as valency in chemistry.
Frankland summarized his thoughts by proposing what he described as a ‘law’:
A tendency or law prevails (here), and that, no matter what the characters of the uniting atoms
may be, the combining power of the attracting element, if I may be allowed the term, is always
satisfied by the same number of these atoms.
Frankland’s work suggested that each element combined with only a limited number of atoms of
another element, thus alluding to the concept of bonding. But it was two other scientists who
performed the most important contemporary research on the concept of bonding.
In 1916, the American scientist Gilbert N. Lewis published a now famous paper
on bonding entitled “The atom and the molecule” (Lewis, 1916). In that paper he outlined a
number of important concepts regarding bonding that are still used today as
working models of electron arrangement at the atomic level. Most significantly, Lewis developed
a theory about bonding based on the number of outer shell, or valence, electrons in an atom. He
suggested that a chemical bond was formed when two atoms shared a pair of electrons (later
renamed a covalent bond by Irving Langmuir). His "Lewis dot diagrams" used a pair of dots to
represent each shared pair of electrons that made up a covalent bond (Figure 2).
Figure 2:
Lewis dot structures for the elements in the first two periods of the periodic table. The structures are
written as the element symbol surrounded by dots that represent the valence electrons.
Lewis also championed the idea of ‘octets’ (groups of eight), that a filled valence shell was
crucial in understanding electronic configuration as well as the way atoms bond together. The
octet had been discussed previously by chemists such as John Newland, who felt it was
important, but Lewis advanced the theory.
Comprehension Checkpoint
Lewis based his theory of bonding on

the total number of electrons in an atom.

the number of electrons in an atom's outer shell.
The modern chemical bond
While still in college, a young chemist by the name of Linus Pauling familiarized himself with
Lewis’s work and began to consider how it might be interpreted within the context of the newly
developed field of quantum mechanics. The theory of quantum mechanics, developed in the first
half of the 20th century, had redefined our modern understanding of the atom and so any theory
of bonding would be incomplete if it were not consistent with this new theory (see our
modules Atomic Theory II: Bohr and the Beginnings of Quantum Theory and Atomic Theory III:
Wave-Particle Duality and the Electron for more information).
Pauling’s greatest contribution to the field was his book The Nature of the Chemical
Bond (Pauling, 1939). In it, he linked the physics of quantum mechanics with the chemical
nature of the electron interactions that occur when chemical bonds are made. Pauling’s work
concentrated on establishing that true ionic bonds and covalent bonds sit at extreme ends of
a bonding spectrum, and that most chemical bonds are classified somewhere between those
extremes. Pauling further developed a sliding scale of bond type governed by
the electronegativity of the atoms participating in the bond.
Pauling’s immense contributions to our modern understanding of the chemical bond led to his
being awarded the 1954 Nobel Prize for "research into the nature of the chemical bond and its
application to the elucidation of the structure of complex substances."
Types of chemical bonds
Chemical bonding and interactions between atoms can be classified into a number of different
types. For our purposes we will concentrate on two common types of chemical bonds, namely
covalent and ionic bonding.
Molecular bonds are formed when constituent atoms come close enough together such that the
outer (valence) electrons of one atom are attracted to the positive nuclear charge of its neighbor.
As the independent atoms approach one another, there are both repulsive forces (between the
electrons in each atom and between the nuclei of each atom), and attractive forces (between the
positive nuclei and the negative valence electrons). Some constituents require the addition
of energy, called the activation energy, to overcome the initial repulsive forces. But at various
distances, the atoms experience different attractive and repulsive forces, ultimately finding the
ideal separation distance where the electrostatic forces are reduced to a minimum. This minimum
represents the most stable position, and the distance between the atoms at this point is known as
the bond length.
Covalent bonding
As the name suggests, covalent bonding involves the sharing (co, meaning joint)
of valence (outer shell) electrons. As described previously, the atoms involved in covalent
bonding arrange themselves in order to achieve the greatest energetic stability. And the valence
electrons are shared – sometimes equally, and sometimes unequally – between neighboring
atoms. The simplest example of covalent bonding occurs when two hydrogen atoms come
together to ultimately form a hydrogen molecule, H2 (Figure 3).
Figure 3:
Here the interaction of two gaseous hydrogen atoms is charted showing the potential energy (purple line)
versus the internuclear distance of the atoms (in pm, trillionths of a meter). The observed minimum in
potential energy is indicated as the bond length (r) between the atoms.image © Saylor Academy
The covalent bond in the hydrogen molecule is defined by the pair of valence electrons (one
from each hydrogen atom) that are shared between the atoms, thus giving each hydrogen atom a
filled valence shell. Since one shared pair of electrons represents one covalent bond, the
hydrogen atoms in a hydrogen molecule are held together with what is known as a single
covalent bond, and that can be represented with a single line, thus H-H.
Multiple covalent bonds
There are many instances where more than one pair of valence electrons are shared
between atoms, and in these cases multiple covalent bonds are formed. For example, when four
electrons are shared (two pairs), the bond is called a double covalent bond; in the case of six
electrons being shared (three pairs) the bond is called a triple covalent bond.
Common examples of such multiple bonds are those formed between atoms in oxygen and
nitrogen gas. In oxygen gas (O2), two atoms share a double bond resulting in the structure O=O.
In nitrogen gas (N2), a triple bond exists between two nitrogen atoms, N≡N (Figure 4).
Figure 4: The bonds
between gaseous oxygen and nitrogen atoms. In oxygen gas (O2), two atoms share a double bond resulting
in the structure O=O. In nitrogen gas (N2), a triple bond exists between two nitrogen atoms, N≡N.
Double covalent bonds are shorter and stronger than comparable single covalent bonds, and in
turn, triple bonds are shorter and stronger than double bonds – nitrogen gas, for example, does
not react readily because it is a strongly bonded stable compound.
Comprehension Checkpoint
When four electrons are shared between atoms, _____ bonds are formed.

double

quadruple
Ions and ionic bonding
Ionic bonding occurs when valence electrons are shared so unequally that they spend more time
in the vicinity of their new neighbor than their original nuclei. This type of bond is classically
described as occurring when atoms interact with one another to either lose or gain electrons.
Those atoms that have lost electrons acquire a net positive charge and are called cations, and
those that have gained electrons acquire a net negative charge and are referred to as anions. The
number of electrons gained or lost by a constituent atom commonly conforms with Lewis’s
valence octets, or filled valence shell principle.
In reality even the most classic examples of ionic bonding, such as the sodium chloride bond,
contain characteristics of covalent bonding, or sharing of electrons of outer shell electrons. A
common misconception is the idea that elements tend to bond with other elements in order to
achieve these octets because they are 'stable' or, even worse, 'happy', and that’s what elements
'want'. Elements have no such feelings; rather, the actual reason for bond formation should be
considered in terms of the energetic stability arising from the electrostatic interaction of
positively charged nuclei with negatively charged electrons.
Substances that are held together by ionic bonds (like sodium chloride) can commonly separate
into true charged ions when acted upon by an external force, such as when they dissolve in
water. Further, in solid form, the individual atoms are not cleanly attracted to one individual
neighbor, but rather they form giant networks that are attracted to one another by the electrostatic
interactions between each atom’s nucleus and neighboring valence electrons. The force of
attraction between neighboring atoms gives ionic solids an extremely ordered structure known as
an ionic lattice, where the oppositely charged particles line up with one another to create a rigid,
strongly bonded structure (Figure 5).
Figure 5: A sodium
chloride crystal, showing the rigid, highly organized structure.
The lattice structure of ionic solids conveys certain properties common to ionic substances.
These include:



High melting and boiling points (due to the strong nature of the ionic bonds throughout the lattice).
An inability to conduct electricity in solid form when the ions are held rigidly in fixed positions within
the lattice structure. Ionic solids are insulators. However, ionic compounds are often capable of
conducting electricity when molten or in solution when the ions are free to move.
An ability to dissolve in polar solvents such as water, whose partially charged nature leads to an attraction
to the oppositely charged ions in the lattice.
The special properties of ionic solids are discussed in further detail in the module Properties of
Solids.
Comprehension Checkpoint
Atoms that lose electrons and acquire a net positive charge and are called

anions.

cations.
Lewis dot diagrams
Lewis used dots to represent valence electrons. Lewis dot diagrams (see Figure 1) are a quick
and easy way to show the valence electron configuration of individual atoms where
no bonds have yet been made.
The dot diagrams can also be used to represent the molecules that are formed when
different species bond with one another. In the case of molecules, dots are placed between
two atoms to depict covalent bonds, where two dots (a shared pair of electrons) denote a single
covalent bond. In the case of the hydrogen molecule discussed above, the two dots in the Lewis
diagram represent a single pair of shared electrons and thus a single bond (Figure 6).
Figure 6: Two
hydrogen atoms are connected by a covalent bond. This can be represented by two dots (left) or a single
bar (right).
When is it ionic? When is it covalent?
If ionic bonding and covalent bonding sit at the extreme ends of a bonding spectrum, how do we
know where any particular compound sits on that spectrum? Pauling’s theory relies upon the
concept of electronegativity, and it is the differences in electronegativity between the atoms that
is crucial in determining where any bond might be placed on the sliding scale of bond type.
Pauling’s scale of electronegativity assigns numbers between 0 and 4 to each chemical element.
The larger the number, the higher the electronegativity and the greater the attraction that element
has for electrons. The difference in electronegativity between two species helps identify
the bond type. Ionic bonds are those in which a large difference in electronegativity exists
between two bonding species. Large differences in electronegativity usually occur when metals
bond to non-metals, so bonds between them tend to be considered ionic.
When the difference in electronegativity between the atoms that make up the chemical bond is
less, then sharing is considered to be the predominant interaction, and the bond is considered to
be covalent. While it is by no means absolute, some consider the boundary between ionic and
covalent bonding to exist when the difference in electronegativity is around 1.7 – less of a
difference tends toward covalent, and a larger difference tends towards ionic. Smaller differences
in electronegativity usually occur between elements that are both considered non-metals, so
most compounds that are made up from two non-metal atoms are considered to be covalent.
Comprehension Checkpoint
If there is a big difference in electronegativity between two different elements, the bond
between them will be

ionic.

covalent.
How covalent is covalent?
Once differences in electronegativity have been considered, and a bond has been determined as
being covalent, the story is not quite over. Not all covalent bonds are created equally. The only
true, perfectly covalent bond will be one where the difference in electronegativity between the
two atoms within the bond is equal to zero. When this occurs, each atom has exactly the same
attraction for the electrons that make up the covalent bond, and therefore the electrons are
perfectly shared. This typically occurs in diatomic (two-atom) molecules such as H2, N2, O2, and
those of the halogen compounds when the atoms in the bond are identical.
However, most covalent bonds occur between elements where even though
the electronegativity difference is lower than 1.7, it is not zero. In these cases, the electrons are
still considered shared, that is, the bond is still considered covalent, but the sharing is not perfect.
Polarity and dipoles in covalent molecules
Most covalent bonds are formed between atoms of differing electronegativity, meaning that the
shared electrons are attracted to one atom within the bond more than the other. As a result, the
electrons tend to spend more time at one end of the bond than the other. This sets up what is
known as a dipole, literally meaning ‘two poles’. One end of the bond is relatively positive (less
attraction for electrons), and one end of the bond is relatively negative (more attraction for
electrons). If this difference in electron affinity exists across the molecule, then the molecule is
said to be polar – meaning that it will have two different, and opposite, partial charges at either
end. Water (H2O) is an excellent example of a polar molecule. Electrons are not shared evenly
since hydrogen and oxygen have different electronegativities. This creates dipoles in each H-O
bond, and these dipoles do not cancel each other out, leaving the water molecule polar overall
(Figure 7). (Read more about these bonds in our module Properties of Liquids.)
Figure 7: In panel A, a molecule of water, H2O, is shown with uneven electron sharing resulting in a
partial negative charge around the oxygen atom and partial positive charges around the hydrogen atoms.
In panel B, three H2O molecules interact favorably, forming a dipole-dipole interaction between the
partial charges.
When the electrons in a bond are perfectly shared, there is no dipole, and neither end of the bond
carries any partial charge. When no such overall charge exists, the molecule is said to be nonpolar. An example of such a non-polar molecule is hydrogen, H2. In larger molecules with
multiple covalent bonds, each bond will have either no dipole or a dipole with varying degrees of
partial charge. When all of these dipoles are taken into consideration in three dimensions, the
uneven distribution of charge caused by the dipoles may cancel out, making the molecule nonpolar.
Alternatively, there may be a partial electrical charge across the molecule, making it
a polar molecule. An example of a multiple atom non-polar molecule is carbon
dioxide. Electrons are not shared evenly across the C=O bonds since carbon and oxygen have
different electronegativities. This creates dipoles in each C=O bond, but because these are
aligned oppositely across a linear molecule, with the oxygen atoms on either side of the carbon
atom, they cancel via symmetry to leave the carbon dioxide molecule non-polar (Figure 8).
Figure 8:
Electrons are not shared evenly across the C=O bonds in CO2 and thus it contains two dipoles. Since these
two dipoles are opposite to one another across a linear molecule, they cancel via symmetry to leave the
carbon dioxide molecule non-polar.image © Molecule: FrankRamspott/iStockphoto
Other types of bonding and the future
We have limited our discussion to ionic and covalent bonding and the sliding scale of bond type
that exists between them. However, many other types of interactions and bonds
between atoms exist, notably metallic bonding (the attractions that hold metal atoms together in
metallic elements), and intermolecular forces (the interactions that exist between, rather than
within, covalently bonded molecules). These each involve similar electrostatic interactions to the
ones described in ionic and covalent bonds, but even those extensions are far from the end of the
bonding story.
In 2014, researchers found the first experimental evidence for a new type of interaction
between atoms that had been predicted in the 1980s (Fleming et al., 2014). Named a "vibrational
bond," the theory describes a lightweight element (in this case, an isotope of hydrogen)
oscillating or "bouncing" between two much heavier atoms (in this case, bromine) and
effectively holding the larger atoms together. Donald Fleming, a chemist based at the University
of British Columbia in Canada, described the new bond as being "like a Ping Pong ball bouncing
between two bowling balls." As research continues, we can expect to understand interactions at
the molecular level with increasing sophistication, and with it, a greater understanding of what
we call chemical bonding.
Summary
The millions of different chemical compounds that make up everything on Earth are composed
of 118 elements that bond together in different ways. This module explores two common types
of chemical bonds: covalent and ionic. The module presents chemical bonding on a sliding scale
from pure covalent to pure ionic, depending on differences in the electronegativity of the
bonding atoms. Highlights from three centuries of scientific inquiry into chemical bonding
include Isaac Newton’s ‘forces’, Gilbert Lewis’s dot structures, and Linus Pauling’s application
of the principles of quantum mechanics.
Key Concepts

When a force holds atoms together long enough to create a stable, independent entity, that force
can be described as a chemical bond.

The 118 known chemical elements interact with one another via chemical bonds, to create brand
new, unique compounds that have entirely different chemical and physical properties than the
elements that make them up.

It is helpful to think of chemical bonding as being on a sliding scale, where at one extreme there
is pure covalent bonding, and at the other there is pure ionic bonding. Most chemical bonds lie
somewhere between those two extremes.

When a chemical bond is formed between two elements, the differences in the electronegativity
of the atoms determine where on the sliding scale the bond falls. Large differences in
electronegativity favor ionic bonds, no difference creates non-polar covalent bonds, and
relatively small differences cause the formation of polar-covalent bonds.

NGSS


HS-C4.3, HS-C6.2, HS-PS1.A3, HS-PS1.B1
Further Reading

The Periodic Table of Elements
The Periodic Table of Elements
In 1869, the Russian chemist Dmitri Mendeleev first proposed that the
chemical elements exhibited a "periodicity of properties." Mendeleev had tried to organize the
chemical elements according to their atomic weights, assuming that the properties of the
elements would gradually change as atomic weight increased. What he found, however, was that
the chemical and physical properties of the elements increased gradually and then suddenly
changed at distinct steps, or periods. To account for these repeating trends, Mendeleev grouped
the elements in a table that had both rows and columns.
The Periodic Table of Elements
Arrangement of the modern periodic table
The modern periodic table of elements is based on Mendeleev's observations; however, instead
of being organized by atomic weight, the modern table is arranged by atomic number (z). As
one moves from left to right in a row of the periodic table, the properties of the elements
gradually change. At the end of each row, a drastic shift occurs in chemical properties. The next
element in order of atomic number is more similar (chemically speaking) to the first element in
the row above it; thus a new row begins on the table.
For example, oxygen (O), fluorine (F), and neon (Ne) (z = 8, 9 and 10,respectively) all are stable
nonmetals that are gases at room temperature. Sodium (Na, z = 11), however, is a silver metal
that is solid at room temperature, much like the element lithium (z = 3). Thus sodium begins a
new row in the periodic table and is placed directly beneath lithium, highlighting their chemical
similarities.
Rows in the periodic table are called periods. As one moves from left to right in a given period,
the chemical properties of the elements slowly change. Columns in the periodic table are
called groups. Elements in a given group in the periodic table share many similar chemical and
physical properties.
Comprehension Checkpoint
Why does sodium appear directly below lithium in the periodic table?

Sodium comes after lithium alphabetically.

Sodium is similar to lithium in terms of chemical properties.
Electron configuration and the table
The "periodic" nature of chemical properties that Mendeleev had discovered is related to
the electron configuration of the atoms of the elements. In other words, the way in which an
atom's electrons are arranged around its nucleus affects the properties of the atom.
Niels Bohr's theory of the atom tells us that electrons are not located randomly around an
atom's nucleus, but they occur in specific electron shells (see our Atomic Theory II module for
more information). Each shell has a limited capacity for electrons. As lower shells are filled,
additional electrons reside in more-distant shells.
The capacity of the first electron shell is two electrons and for the second shell the capacity is
eight. Thus, in our example discussed above, oxygen, with eight protons and eight electrons,
carries two electrons in its first shell and six in its second shell. Fluorine, with nine electrons,
carries two in its first shell and seven in the second. Neon, with ten electrons, carries two in the
first and eight in the second. Because the number of electrons in the second shell increases, we
can begin to imagine why the chemical properties gradually change as we move from oxygen to
fluorine to neon.
Sodium has eleven electrons. Two fit in its first shell, but remember that the second shell can
only carry eight electrons. Sodium's eleventh electron cannot fit into either its first or its second
shell. This electron takes up residence in yet another orbit, a third electron shell in sodium. The
reason that there is a dramatic shift in chemical properties when moving from neon to sodium is
because there is a dramatic shift in electron configuration between the two elements. But why is
sodium similar to lithium? Let's look at the electron configurations of these elements.
Electron Configurations for Selected Elements
As you can see in the illustration, while sodium has three electron shells and lithium two, the
characteristic they share in common is that they both have only one electron in their
outermost electron shell. These outer-shell electrons (called valence electrons) are important in
determining the chemical properties of the elements.
An element's chemical properties are determined by the way in which its atoms interact with
other atoms. If we picture the outer (valence) electron shell of an atom as a sphere encompassing
everything inside, then it is only the valence shell that can interact with other atoms – much the
same way as it is only the paint on the exterior of your house that "interacts" with, and gets wet
by, rain water.
An atom's valence shell “covers” inner electron shells
The valence shell electrons in an atom determine the way it will interact with neighboring atoms,
and therefore determine its chemical properties. Since both sodium and lithium have one valence
electron, they share similar chemical properties.
Comprehension Checkpoint
The chemical properties of an element are determined by the number of electrons in

the electron shell closest to the nucleus of the atom.

the outermost electron shell.
Electron configuration shorthand
For elements in groups labeled A in the periodic table (IA, IIA, etc.), the number
of valence electrons corresponds to the group number. Thus Li, Na, and other elements in group
IA have one valence electron. Be, Mg, and other group-IIA elements have two valence electrons.
B, Al, and other group-IIIA elements have three valence electrons, and so on. The row, or period,
number that an element resides in on the table is equal to the number of total shells that contain
electrons in the atom. H and He in the first period normally have electrons in only the first shell;
Li, Be, B, and other period-two elements have two shells occupied, and so on. To write the
electron configuration of elements, scientists often use a shorthand in which the element's
symbol is followed by the element's electron shells. A few examples are shown below.
Element Configuration Shorthand
Element
Configuration Shorthand
Hydrogen
H
1e-
Lithium
Li
2e-
1e-
Fluorine
F
2e-
7e-
Sodium
Na
2e-
8e-
1e-
For further details, the table linked below shows the electron configurations of the first
eleven elements.
Interactive Animation:Atomic and ionic structure of the first 12 elements
Summary
The modern periodic table is based on Dmitri Mendeleev’s 1896 observations that chemical
elements can be grouped according to chemical properties they exhibit. This module explains the
arrangement of elements in the period table. It defines periods and groups and describes how
various electron configurations affect the properties of the atom.


Chemical Reactions
Chemical Reactions
Chemical reactions happen absolutely everywhere. While we sometimes associate chemical
reactions with the sterile environment of the test tube and the laboratory - nothing could be
further from the truth. In fact, the colossal number of transformations make for a dizzying,
almost incomprehensible array of new substances and energy changes that take place in our
world every second of every day.
In nature, chemical reactions can be much less controlled than you’ll find in the lab, sometimes
far messier, and they generally occur whether you want them to or not! Whether it be a fire
raging across a forest (Figure 1), the slow process of iron rusting in the presence of oxygen and
water over a period of years, or the delicate way in which fruit ripens on a tree, the process of
converting one set of chemical substances (the reactants) to another set of substances (the
products) is one known as a chemical reaction.
Figure 1: A controlled fire in Alberta, Canada, set to create a barrier for future
wildfires.image © Cameron Strandberg, Rocky Mountain House
Though chemical reactions have been occurring on Earth since the beginning of time, it wasn’t
until the 18th century that the early chemists started to understand them. Processes like
fermentation, in which sugars are chemically converted into alcohol, have been known for
centuries; however, the chemical basis of the reaction was not understood. What were these
transformations and how were they controlled? These questions could only be answered when
the transition from alchemy to chemistry as a quantitative and experimental science took place.
Historical context
Beginning in the early Middle Ages, European and Persian philosophers became fascinated with
the way that some substances seemed to “transmute” (or transform) into others. Simple stones,
such as those that contained sulfur, seemed to magically burn; and otherwise
unimpressive minerals were transformed, like the ore cinnabar becoming an enchanting
silvery liquid metal mercury when heated. Alchemists based their approach on Aristotle’s ideas
that everything in the world was composed of four fundamental substances - air, earth, fire, and
water (Figure 2).
Figure 2: Aristotle believed that everything in the world was composed of four fundamental
substances - air, earth, fire, and water.
As such, they proposed, and spent generations trying to prove, that less expensive metals like
copper and mercury could be turned into gold. Despite their misguided approach, many early
alchemists performed foundational chemical experiments - transforming one substance into
another, and so it is difficult to point to a specific date or event as the birth of the idea of an
ordered, quantifiable chemical reaction. However, there are some important moments in history
that have helped to make sense of it.
Lavoisier: Law of Mass Conservation
Antoine Lavoisier was a French nobleman in the 1700s who began to experiment with
different chemical reactions. At the time, chemistry still couldn’t be described as being a true,
quantitative science. Most of the theories that existed to explain the way that substances changed
relied upon Greek philosophy, and there was precious little experimental detail attached to the
alchemist’s tinkering.
However, during the second half of the 18th century, Lavoisier performed many
quantitative experiments and observed that while substances changed form during a chemical
reaction, the mass of the system – or a measure of the total amount of “stuff” present – did not
change. In doing so, Lavoisier championed the idea of conservation of mass during
transformations (Figure 3). In other words, unlike the alchemists before him who thought that
they were creating matter out of nothing, Lavoisier proposed that substances are neither created
nor destroyed, but rather change form during reactions. Lavoisier’s ideas were published in the
seminal work Traité élémentaire de Chimie in 1789 (Lavoisier, 1789), which is widely hailed as
the birth of modern chemistry as a quantitative science.
Figure 3: Lavoisier's Law of Mass Conservation, which states that substances are neither
created nor destroyed, but rather change form during reactions. In this example, the reactants
(zinc and two hydrogen chloride molecules) convert into different products (zinc chloride and
dihydrogen), but no mass is lost or created.
Proust: Law of Constant Composition
Joseph Proust was a French actor who followed in Lavoisier’s footsteps. Proust performed
dozens of chemical reactions, starting with different amounts of various materials. Over time he
observed that no matter how he started a certain chemical reaction, the ratio in which
the reactants were consumed was always constant. For example, he worked extensively with
copper carbonate and no matter how he changed the ratio of starting reactants, the copper,
carbon, and oxygen all reacted together in a constant ratio (Proust, 1804). As a result, in the last
few years of the 18th century, Proust formulated the law of constant composition (also referred
to as the law of definite proportions, Figure 4).
He realized that any given chemical substance (that we now define as a compound) always
consisted of the same ratio by mass of its elemental parts regardless of the method of
preparation. This was a huge step forward in modern chemistry since it had been previously
believed that the substances formed during chemical reactions were random and disordered.
Figure 4: An example of Proust's Law of Constant Composition, which states that any
compound always consists of the same ratio by mass of its elemental parts, regardless of the
method of preparation.
Dalton: Law of Multiple Proportions
The English chemist John Dalton helped make sense of the laws of conservation of mass and
definite proportions in 1803 by proposing that matter was made of atoms of unique substances
that could not be created or destroyed (see our module Early Ideas about Matter for more
information).
Dalton extended Proust’s ideas by recognizing that it was possible for two elements to form
more than one compound, but that whatever the compound was, it would always contain
elements combined in whole number ratios (Dalton, 1808). This observation is known as
the law of multiple proportions (Figure 5) and with his atomic theory, helped to cement
Lavoisier’s observations.
Figure 5: Dalton's Law of Multiple Proportions, which states that two elements may form
more than one compound, but whatever the compound was, it would always contain elements
combined in whole number ratios
These advancements, taken together, laid the groundwork for our modern understanding
of chemical reactions, chemical equations, and chemical stoichiometry, or the process of
expressing the relative quantities of reactants and products in a chemical reaction.
Comprehension Checkpoint
____ first theorized that while substances changed form during a chemical reaction, the
mass of the system did not change.

Lavoisier

Dalton
Types of chemical reactions
There is a staggering array of chemical reactions. Chemical reactions occur constantly within our
bodies, within plants and animals, in the air that circulates around us, in the lakes and oceans that
we swim in, and even in the soil where we grow crops and build our homes. In fact, there are so
many chemical reactions that occur that it would be difficult, if not impossible, to understand
them all. However, one method that helps us to understand them is to categorize chemical
reactions into a few, general types. While not a perfect system, placing reactions together
according to their similarities helps us to identify patterns, which in turn allows predictions to be
made about as yet unstudied reactions. In this module, we will consider and provide some
context for a few categories of reactions, specifically: synthesis, decomposition, single
replacement, double replacement, REDOX (including combustion), and acid-base reactions.
No matter the type of reaction, one universal truth applies to all chemical reactions. For
a process to be classified as a chemical reaction, i.e., one where a chemical change takes place, a
new substance must be produced. The formation of a new substance is nearly always
accompanied by an energy change, and often with some kind of physical or observable change.
The physical change can be of different types, such as the formation of bubbles of a gas,
a solid precipitate, or a color change. These changes are clues to the existence of a chemical
reaction and are important triggers for further research by chemists.
Synthesis reactions
Prior to Lavoisier’s work, it was poorly understood that there were different gases made up of
different elements. Instead, various gases were commonly mischaracterized as types of "air" or
air missing parts – for example, terms commonly used were "inflammable air," or
"dephlogisticated air." Lavoisier thought differently and was convinced that these were different
substances. He conducted experiments where he mixed inflammable air with dephlogisticated air
and a spark, and he found that the substances combined to produce water. In response, he
renamed inflammable air "hydrogen" from the Greek hydro for "water" and genes for "creator."
In so doing, Lavoisier was identifying a synthesis reaction. In general, a synthesis reaction is one
in which simpler substances combine to form another more complex one. Hydrogen and oxygen
(which Lavoisier also renamed dephlogisticated air) combine in the presence of a spark to form
water, summarized by the chemical equation shown below (for more on chemical equations see
the section called Anatomy of a chemical equation), it represents a simple synthesis reaction.
Equation 1
2H2(g) + O2(g) → 2H2O(l)
Decomposition reactions
In 1774, the scientist Joseph Priestley turned his curiosity to a mineral called cinnabar – a brick
red mineral. When he placed the mineral under sunlight amplified by a powerful magnifying
glass, he found that a gas was produced which he described as having an “exalted nature”
because a candle burned in the gas brightly (Priestley, 1775). Without realizing it, Priestley had
discovered oxygen as a result of a decomposition reaction. Decomposition reactions are often
thought of as the opposite of synthesis reactions since they involve a compound being broken
down into simpler compounds or even elements. In the case of Priestley’s oxygen, he had broken
down mercury (II) oxide (cinnabar) with heat into its individual elements. The reaction can be
summarized in the following equation.
Equation 2
2HgO(s) → 2Hg(l) + O2(g)
Single replacement reactions
The British chemist and meteorologist John Daniell, invented one of the very first practical
batteries in 1836 (Figure 6). In his cell, Daniell utilized a very common single
replacement reaction. His early cells were complicated affairs, with ungainly parts and
complicated constructs, but by contrast, the chemistry behind them was really quite simple.
Figure 6: Daniell cell batteries.
In certain chemical reactions, a single constituent can substitute for another one already joined in
a chemical compound. The Daniell cell works because zinc can substitute for copper in
a solution of copper sulfate, and in so doing exchange electrons that are used in the battery cell.
The reaction can be summarized as follows:
Equation 3
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
This particular single displacement is called a metal displacement since it involves one metal
replacing another metal, and many types of batteries are based on metal replacement reactions.
However, several other types of single replacement reactions exist, such as when a metal can
replace hydrogen from an acid or from water, or a halogen can replace another halogen in
certain salt compounds.
Combustion reactions
The controlled use of fire was a crucial development for early civilization. While it’s difficult to
pin down the exact time that humans first tamed the combustion reactions that produce fire,
recent research suggests it may have occurred at least a million years ago in a South African cave
(Berna et al. 2012).
Chemically, combustion is no more than the reaction of a fuel (wood, oil, gasoline, etc.) with
oxygen. For combustion to take place there must be a fuel and oxygen gas. However, these
reactions often require activation energy (discussed in more detail in the module Chemical
Bonding: The Nature of the Chemical Bond), which can be provided by a ‘spark’ or source
of energy for ignition. Fuel, oxygen, and energy are the three things make up what is known as
the fire triangle (Figure 7), and any one of them being absent means that combustion will not
take place.
Figure 7: The fire triangle is made up of three things - fuel, oxygen, and energy.image ©
Gustavb
In the modern world, many of the fuels that are typically burned for energy, are hydrocarbons –
substances that contain both hydrogen and carbon (as discussed in more detail in our Carbon
Chemistry module). Plants produce hydrocarbons when they grow, and thus make an excellent
fuel source, and other hydrocarbons are produced when plants or animals decay over time (such
as natural gas, oil, and other substances). When these fuels combust, the hydrogen and carbon
within them combine with oxygen to produce two very familiar compounds, water, and carbon
dioxide. One simple example is the combustion of natural gas, or methane, CH4:
Equation 4
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
As with the combustion of all fuels, heat and light are products, too, and it is these products that
are used to cook our food or to heat our homes.
Reduction-oxidation reactions
Each of the four types of reaction above are sub-categories of a single type of chemical
reaction known as redox reactions. A redox reaction is one where reduction and oxidation take
place together, hence the name. The individual processes of oxidation and reduction can be
defined in more than one way, but whatever the definition, the two processes are symbiotic, i.e.,
they must take place together.
In one definition, oxidation is described as the process in which a species loses electrons, and
reduction is a process where a species gains electrons. In this way, we can see how the pair must
take place together. If a chemical substance is to lose electrons (and therefore be oxidized), then
it must have another, interdependent chemical substance that it can give those electrons to. In the
process, the second substance (the one gaining electrons) is said to be reduced. Without such an
electron acceptor, the original species can never lose the electrons and no oxidation can take
place. When the electron acceptor is present, it gets reduced and the redox combination process
is complete. Redox reactions of this type can be summarized by a pair of equations – one to
show the loss of electrons (the oxidation), and the other to show the gain of electrons (the
reduction). Using the example of the Daniell cell above,
Equation 5
Oxidation: Zn → Zn2+ + 2eReduction: Cu2+ + 2e- → Cu
The electrons shown being lost by zinc in the first reaction, are the same electrons being
accepted by the copper ions in the second. Together, the reactions can be combined to cancel out
the electrons on either side of the reactions, into the overall redox reaction:
Equation 6
Zn + Cu2+ → Zn2+ + Cu
Other definitions of oxidation and reduction also exist, but in every case, the two halves of the
redox reaction remain symbiotic – one loses and the other gains. The loss from
one species cannot happen without the other species gaining.
Double displacement reactions
When soap won’t easily produce a lather in water, the water is said to be ‘hard’. Hard water
causes all kinds of problems that go beyond just making it difficult to form a lather. The buildup
of compounds in water pipes (known as ‘scale’), can block the flow of water and can cause
problems in industrial processes. Textile manufacturing and the beverage industry rely heavily
on water. In those situations, the quality of the water can make a difference to the end product, so
controlling the water composition is crucial.
Hard water contains magnesium or calcium ions in the form of a dissolved salt such as
magnesium chloride or calcium chloride. When soap (sodium stearate) comes into contact with
either of those salts, it enters into a double displacement reaction that forms
the insoluble precipitate known as ‘soap scum’.
A double displacement reaction (also known as a double replacement reaction) occurs when two
ionic substances come together and both substances swap partners. In general:
Equation 7
AB + CD → AD + CB
Where A and C are cations (positively charged ions), and B and D are anions (negatively
charged).
In the case of the reaction of soap with calcium chloride, the reaction is:
Equation 8
CaCl2(aq) + 2NA(C17H35COO)(aq) → 2NaCl(aq) + Ca(C17H35COO)2(s)
The solid calcium stearate is what we call soap scum, which is formed by the reaction of
the soluble sodium stearate salt (the soap) in a double replacement reaction with calcium
chloride.
Acid-Base reactions
Acid-base reactions happen around, and even inside of us, all the time. From the classic
elementary school baking soda volcano to the process of digestion, we
encounter acids and bases on a daily basis.
When a hydrogen atom loses its only electron, it forms a positive ion, H+. This hydrogen ion is
the essential component of all acids, and indeed one definition of an acid is that of a hydrogen
ion donor. Compounds such as the citric acid in lemon juice, the ethanoic acid in vinegar, or a
typical laboratory acid like hydrochloric acid, all give their hydrogen ions away in chemical
reactions known as acid-base reactions. The chemical opposites of acids are known as bases, and
bases can be defined as hydrogen ion acceptors. Whenever an acid donates a hydrogen ion to a
base, an acid-base reaction has taken place, for example, when hydrochloric acid donates a
hydrogen ion to a base such as sodium hydroxide:
Equation 9a
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
A closer look at this reaction reveals that in water the HCl gives away an H+ as shown below:
Equation 9b
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
The resulting species, H3O+ (the hydronium ion), can, in turn, act as an acid when it comes into
contact with any species that can accept a hydrogen ion, such as hydroxide ions from sodium
hydroxide:
Equation 9c
H3O+(aq) + NaOH(aq) → 2H2O(l) + Na+(aq)
Combining equations #9a and #9b gives us equation #9c.
Equation #9c can be re-written to show the individual ions that are found in solution, thus:
Equation 9d
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → H2O(l) + Na+(aq) + Cl-(aq)
Removing the spectator ions from the equation above, we get the net ionic equation:
Equation 9e
H+(aq) + OH-(aq) → H2O(l)
Any chemical reaction that forms water from the reaction between an acid and base as in
equation #9e is known as a neutralization reaction.
Comprehension Checkpoint
The type of chemical reaction where a single constituent can substitute for another one
already joined in a chemical compound is:

redox

single replacement
Anatomy of a chemical equation
Chemical equations are always linked to chemical reactions since they are the shorthand by
which chemical reactions are described. That fact alone makes equations incredibly important,
but equations also have a crucial role to play in describing the quantitative aspect of chemistry,
something that we formally call stoichiometry.
All chemical reactions take on the same, basic format. The starting substances, or reactants, are
listed using their chemical formula to the left-hand side of an arrow, with multiple reactants
separated with plus signs. In the case of a reaction between carbon and oxygen:
Equation 10a
C + O2 →
To the right hand of the arrow one finds the chemical formulas of the new substance or
substances (known as the products) that are produced by the chemical reaction. In this case, since
carbon dioxide is the result of burning carbon in the presence of oxygen:
Equation 10b
[Reactants] C + O2 → CO2 [Products]
Since reactions can result in both physical as well as chemical changes, each substance is given a
state symbol written as a subscript to the right of the formula, this describes the physical form of
the reactants and products. Common state abbreviations are (s) for solids, (l) for liquids, (g) for
gases and (aq) for any aqueous substances, i.e., those dissolved in water.
Equation 10c
C(s) + O2(g) → CO2(g)
Finally, in order to ensure that this representation abides by the law of conservation of mass, the
equation may need to be balanced by the addition of numbers in front of each species that create
equal numbers of atoms of each element on each side of the equation. In the case of the
formation of carbon dioxide from carbon and oxygen, there is no need for the addition of such
numbers (called the stoichiometric coefficients), since 1 carbon atom and 2 oxygen atoms appear
on each side of the equation.
Energy changes
In nature, chemical reactions are often driven by exchanges in energy. In this
respect, reactions are generally separated into two categories – those that release energy and
those that absorb energy.
Exothermic reactions are those that release energy to the surroundings (Figure 8,
right). Combustion reactions are an obvious example because the energy released by the reaction
is converted into the light and heat seen in the immediate surroundings.
By contrast, endothermic reactions are those that absorb energy from the surroundings (Figure
8, left). In this situation, one may have to heat up the reaction or add some other form of energy
to the system before seeing the reaction proceed.
Figure 8: On the left is an endothermic reaction, where energy is absorbed from the
surroundings. In contrast, on the right is an exothermic reaction, which releases energy into the
surroundings.
In both cases it is important to note that energy is neither created nor destroyed, rather it is
transferred from one type of energy to another, for example from chemical energy to that
of heat or light. The energy that goes into the formation of chemical bonds is exchanged for other
types of energy with the environment around that reaction. A classic example is
the photosynthesis reaction, in which plants absorb light energy from the sun in order to create
bonds between atoms that make up sugars, which are stored as chemical energy for later use by
the plant. The process of respiration is essentially the reverse of photosynthesis, where the bonds
in sugar molecules are broken and the released energy is then used by the plant.
Comprehension Checkpoint
_____ reactions are those that absorb energy from the surroundings.

Endothermic

Exothermic
The context of chemical reactions
Chemical reactions happen all around us every day. Whether it is a single replacement reaction
in the battery of our flashlight, a synthesis reaction that occurs when iron rusts in the presence of
water and oxygen, or an acid-base reaction that happens when we eat – we experience chemical
reactions in almost everything we do. Understanding these reactions is not an abstract concept
for a chemist in a far off laboratory, rather it is critical to understanding life and the world around
us. To truly master chemical reactions, we need to understand the quantitative aspect of these
reactions, something referred to as stoichiometry, and a concept we will discuss in another
module.
Summary
This modules explores the variety of chemical reactions by grouping them into general types. We
look at synthesis, decomposition, single replacement, double replacement, REDOX (including
combustion), and acid-base reactions, with examples of each.
Key Concepts






The steps from a qualitative science to quantitative one, were crucial in understanding chemistry
and chemical reactions more completely.
When a substance or substances (the reactants), undergo a change that results in the formation of
a new substance or substances (the products), then a chemical reaction is said to have taken
place.
Mass and energy are conserved in chemical reactions. Matter is neither created or destroyed,
rather it is conserved but rearranged to create new substances. No energy is created or destroyed,
it is conserved but often converted to a different form.
Chemical reactions can be classified into different types depending on their nature. Each type has
its own defining characteristics in terms of reactants and products.
Chemical reactions are often accompanied by observable changes such as energy changes, color
changes, the release of gas or the formation of a solid.
Energy plays a crucial role in chemical reactions. When energy is released into the surroundings
the reaction is said to be exothermic; when energy is absorbed from the surroundings the reaction
is said to be endothermic
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NGSS
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Further Reading
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
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

HS-C5.4, HS-PS1.A2, HS-PS1.A3, HS-PS1.B3
Early Ideas about Matter
Atomic Theory I
Atomic Theory II
Chemical Bonding
Acids and Bases
Carbon Chemistry
References

Berna, F., Goldberg, P., Horwitz, L. K., Brink, J., Holt, S., Bamford, M., & Chazan, M.
(2012). Microstratigraphic evidence of in situ fire in the Acheulean strata of Wonderwerk Cave,
Northern Cape province, South Africa. Proceedings of the National Academy of
Sciences, 109(20), E1215-E1220.
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Dalton, John (1808). A New System of Chemical Philosophy.
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Lavoisier, Antoine (1789). Traité Élémentaire de Chimie, présenté dans un ordre
nouveau, et d'après des découvertes modernes.

Priestley, Joseph (1775). "An Account of Further Discoveries in Air". Philosophical
Transactions. 65: 384–94.

Proust Joseph Louis (1804). “Sur les Oxydations Métalliques.” J Phys. 59: 321-343.
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Chemical Equations
References

Fleming, D.G., Manz, J., Sato, K., and Takayanagi, T. (2014). Fundamental change in the
nature of chemical bonding by isotopic substitution. Angewandte Chemie International Edition,
53(50): 13706–13709.
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Frankland, E. (1852). On a new series of organic bodies containing metals. Philosophical
Transactions, 417: 417-444. Retrieved
from http://rstl.royalsocietypublishing.org/content/142/417.full.pdf+html
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Langmuir, I. (1919). The arrangement of electrons in atoms and molecules. Journal of the
American Chemical Society, 41(6): 868-934.

Lewis, G.N. (1916). The atom and the molecule. Journal of the American Chemical Society,
38(4): 762-786.

Newton, I. (1704). Opticks: or, a treatise of the reflexions, refractions, inflexions and colours of
light.
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Pauling, L. (1931). The nature of the chemical bond. Application of results obtained from the
quantum mechanics and from a theory of paramagnetic susceptibility to the structure of
molecules. Journal of the American Chemical Society, 53(4): 1367-1400.