CHAPTER 4: MATTER (continued)
ATOMIC STRUCTURE
The electrons in each atom arrange themselves into
“energy levels” which are called shells, and shells follow
a pattern of “subshells” known as orbitals. The first, or
innermost, shell has a single, spherical orbital. The
second shell contains one spherical orbital and 3
“dumbbell”-shaped orbitals. The third shell includes
one spherical orbital, 3 dumbbell orbitals and 5 orbitals
of even more complex shape, with fourth and higher
order shells continuing this pattern. Furthermore, each
orbital can hold a maximum of 2 electrons. So, we can
summarize atomic shell structure as follows:
SHELL
#
1
2
3
4
…
# OF ORBITALS
1
1+3=4
1+3+5=9
1 + 3 + 5 + 7 = 16
…
MAX # OF
ELECTRONS
12=2
42=8
9  2 = 18
16  2 = 32
…
Up to argon (Ar, with Z = 18), the atomic structure
involves filling the orbitals of each shell in order (first
the single spherical, then the dumbbells) before moving
to the next shell. Beyond argon, this pattern changes
and becomes a bit more complicated.
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Argon has 18 protons and electrons, while the next two
elements, potassium (K, with Z = 19) and calcium (Ca,
with Z = 20), have 19 and 20 protons and electrons,
respectively.
Argon has a full first shell (2 electrons), a full second
shell (8 electrons) and the first four orbitals (spherical
and dumbbell) of its third shell completely full (8
electrons, for a total of 18). Even though argon’s outer
(third) shell still isn’t full, its first four orbitals are full,
which is why argon is “inert” and doesn’t react with
other elements.)
But, rather than beginning to fill the remaining five
orbitals of the third shell (which still have room for 10
more electrons), the 19th electron for potassium is
actually found in the first (spherical) orbital of the
fourth shell, while the 19th and 20th electrons for calcium
fill up that same spherical orbital. Only in the next ten
elements after calcium, do additional electrons “go
back” and, one-by-one, fill the remaining five orbitals of
shell number 3.
We won’t concern ourselves further with this
complication, but if you are interested in more detail on
the shells and orbitals, refer to Chapter 15 – The
Periodic Table.
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There is one final, but important point, to be made about
atomic structure. It has already been suggested that
atoms (like helium, neon and argon) whose outer shells
are full (or which have a full subset of orbitals) tend to
be “inert”. In contrast, atoms whose outer shells are
nearly empty (say 1 or 2 electrons only) will have a
tendency to shed those electrons, while those whose outer
shells are nearly full (say except for 1 or 2 electrons) will
have a tendency to fill those empty spaces. As we will
see, these two types of atoms are thus highly “reactive”
and readily combine with one another.
CHEMICAL BONDING (see Chapter 9 for more detail)
With a few exceptions, matter doesn’t normally exist as
single atoms. Rather, two or more atoms combine to
form molecules. (The exceptions are “inert” atoms like
those mentioned above which don’t react with anything
else, and hence their individual atoms are also individual
molecules!)
The mechanisms whereby atoms bind together into
molecules are called chemical bonds and we will discuss
three types: ionic bonds, covalent bonds and metallic
bonds. In ionic bonds, atoms give up electrons to become
positive ions or gain electrons to become negative ions,
and then these oppositely-charged ions attract one
another and form molecules. Covalent bonds are subtly
different, though the end result is the same. Rather than
giving up or gaining electrons, “reactive” atoms actually
come together and share pairs of electrons to “effectively
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fill” their outermost orbitals. In metallic bonds, electrons
in the outermost shells of the lattice structures found in
some common metals (e.g. copper) are essentially “set
free” to wander from atom to atom, giving these metals
an important characteristic known as high electrical
conductivity. Meanwhile, these negatively-charge free
electrons are still attracted to the positively-charged
nuclei within the lattice, binding the whole structure
together.
Ionic Bonds
Some atoms are readily able to “ionize” by giving up or
gaining electrons. Atoms which can lose electrons and
become positive ions are classified as metals, while those
that can gain electrons to become negative ions are called
nonmetals.
The metal sodium (Na, with Z = 11), for example, has full
first and second shells (2 + 8 electrons) and a single
electron in its outer (third) shell. So, it easily loses this
outermost electron to become a positive ion, Na+. On the
other hand, the non-metal chlorine (Cl, with Z = 17), also
has full first and second shells, but the first 4 orbitals of
its third shell are only one electron shy of being full.
Chlorine thus readily accepts an extra electron to
complete the filling of its “dumbbell” orbitals and
become an negative ion, Cl–. If these two ions find
themselves in close proximity, they will “attach” and be
held together by mutual electrostatic attraction to form
an atom everyday “salt” (sodium chloride or NaCl). We
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can show this ionic bonding reaction in a chemical
equation, .i.e.
Na+ + Cl–  NaCl .
Note that individual ions can also consist of a cluster of
two or more atoms. One of the most familiar and
common of these is the negative hydroxyl ion OH–,
formed by an atom of oxygen (which has room for two
additional electrons in its second shell) and an atom of
hydrogen (which can provide the single electron in its
first shell. In fact, at any given moment, liquid water
always contains both positive and negative ions, as a
result of molecules of water spontaneously splitting
according to
H2O  H+ + OH– .
The presence of these ions in water explains why water is
able, at least to some degree, to conduct electricity (and
why, for example, it isn’t recommended that you swim
during a thunderstorm!)
Covalent Bonds
More common than ionic bonding is covalent bonding,
where pairs of electrons are “shared” between two atoms
rather than being given up altogether by one and gained
altogether by the other.
The bond by which two atoms of hydrogen and one atom
of oxygen form a molecule of water (appropriately
designated H2O) is perhaps the most familiar example of
a covalent bond. The atom of oxygen has room for two
more electrons in the “dumbbell” orbitals of its second
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shell and is therefore open to receiving two electrons to
fill that shell. Meanwhile, each of the hydrogen atoms
has only one electron in its first (spherical) shell and thus
has room for one additional electron. In the resulting
covalent bond, each of the two hydrogen atoms shares a
pair of electrons with oxygen (i.e. the sole hydrogen
electron and one oxygen electron are shared). The end
result is that the shells of all three atoms are effectively
“filled” on a “part-time” basis.
One way that chemists denote the sharing of a pair of
electrons by two atoms is to join their symbols with a
short line. Because oxygen does this with each of two
hydrogen atoms, there would be two such lines, i.e. the
covalent bond associated with H2O is represented by H –
O – H.
The way that these three atoms physically “attach” in
fact results in a triangular-shaped molecule which looks
like this:
The actual angle between the two “hydrogen arms” is
about 104. However, since the oxygen atom is larger
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than the hydrogen atoms, the overall shape of the
molecule is essentially that of an equilateral triangle, a
fact which will later turn out to be quite significant.
Other Examples of Covalent Bonds
There are many other examples of covalent bonds, of
course, including covalent bonds between atoms of the
same element.
1. A molecule of hydrogen consists of two hydrogen
atoms which share their respective two electrons in a
covalent bond, essentially “filling” the shells of both.
Again, using a line to denote this shared pair of
electrons, we can represent “molecular hydrogen”,
or H2, as H – H.
2. “Molecular oxygen” likewise consists of two oxygen
atoms which share two pairs of electrons, thereby
“filling” the second shells of both. Thus, O2 can be
shown as O = O. Similarly, molecular nitrogen (or
N2) shares three electron pairs and is represented by
N  N.
3. As a final example of covalent bonding, carbon
dioxide, or CO2, consists of one atom of carbon
(whose second shell, with 4 electrons, is only half
full) which shares two pairs of electrons with each of
two atoms of oxygen (whose second shells are 2
8
electrons short). In other words, carbon dioxide’s
covalent bond can be represented by O = C = O.
Metallic Bonds
Consider the example of copper (Cu, with Z = 29), whose
first 3 shells are completely filled, and whose single
remaining electron occupies the first orbital of shell
number 4. Normally, one would expect that this would
give copper a very high tendency to ionize by losing this
outer electron altogether, or to share it covalently. But
what happens in the case of pure metallic copper (as
opposed to molecular compounds which include copper)
is slightly different.
In general, when considering the atomic structure of any
atom, electrons found in “higher” shells have more
energy. At the same time, being slightly further from the
positively-charged nucleus, these same electrons are held
a little less strongly by electrostatic attraction. Putting
these two facts together, it makes sense that it is “easier”
(especially when an orbital isn’t full) for such outermost
electrons to “break free” temporarily and “jump” to an
adjacent atom.
Even a small sample of copper metal contains billions of
atoms within its lattice structure, each of which has one
4th shell electron with both sufficiently high energy and
sufficiently great distance from the nucleus to “escape”
on a regular basis. For this reason, we say that metallic
copper has one free electron for each atom in the sample.
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These outer electrons essentially wander randomly from
atom to atom, so that, statistically speaking, the atoms
near any one location remain on average electrically
neutral. However, if a voltage is applied across the ends
of a copper wire, these same electrons begin to move
steadily through the sample from one end (where
additional electrons are being injected) to the other end
(where they can leave). In other words, the free electrons
within metallic copper at any moment become the
carriers of electric current.
Those metals (like copper and aluminum) with the
highest proportions of free electrons are thus the best
electrical conductors. In particular, not only does
copper have the highest electrical conductivity of any
metal, but because it is both relatively plentiful and
relatively cheap here on Earth, it is by far the most
widely employed metal in electrical wiring and circuitry.
That said, there have been some concerns expressed in
recent years that the world’s copper supply will someday
be exhausted, and this could cause serious problems.
WATER (A MIRACULOUS COMPOUND)
Earth is often called the “water planet” not only because
of the abundance of water (found mainly in the oceans)
but because water was essential to the initial
development and survival of life on planet Earth and it
remains essential today to sustaining all terrestrial
lifeforms.
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We have seen already that, within a very narrow range
of temperatures, water can readily change among its
three states: solid (ice); liquid (water); and gas (water
vapour). This fact is critical to the familiar, so-called
water cycle, the global mechanism whereby the Earth’s
precious water is continually being re-circulated through
evaporation, condensation and precipitation, with a lot of
help from gravity. (Note that since water molecules can
go directly from ice to water vapour, evaporation can
continue in severe winter climates, though most
evaporation of course continues to occur over the
oceans.)
There is another amazing property possessed by water
which deserves mention here. Virtually all substances in
their solid or liquid state expand when heated and
contract when cooled. Water is “normal” in that, as it is
cooled, liquid water contracts until it reaches 4 C, and,
once frozen (at 0 C), ice continues to contract as it cools
further. However, between 4 C and 0 C, liquid water
actually expands. To begin to understand how this
might be the case, we must introduce what is called the
hydrogen bond.
Consider the following diagram showing a sample of
water molecules in their liquid state.
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We have seen earlier that individual molecules of H2O
are shaped like equilateral triangles. Within liquid
water, these molecules move around quite freely, and are
distributed and oriented more-or-less randomly. The socalled hydrogen bond is not a true chemical bond at all.
Rather, it refers to a much weaker electrostatic
attraction (perhaps 10% as strong as the covalent bond
within each molecule) which can exist between a
hydrogen atom from one molecule and an oxygen atom
from an adjacent molecule. (These bonds are shown as
dark zig-zags above.) It is thought that these weak
“inter-molecular” attractions help explain why liquid
water “sticks together”.
Note in the lower left that, depending on their
orientations, water molecules can even get very close
together and have two hydrogen bonds between them.
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When liquid water freezes and turns to ice, we also know
that its structure (like most solids) changes to a crystal
lattice. There are in fact many different such lattices
that can form when water freezes. Sometimes individual
crystals combine in long, thin slivers, while at other
times they form larger, flat shapes (like those found in
“snowflakes”). The diagram below represents a typical
(but by no means the only possible) alignment of a
collection of water molecules as they crystallize during
ice formation.
The thing to notice is that, because individual water
molecules are triangular, these crystal lattice formations
will tend to end up with a characteristic 6-sided, or
hexagonal, shape. (In fact, both “slivers” of ice and
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“snowflakes”, when viewed microscopically, exhibit
hexagonal crystal structures.)
Note that in a hexagonal lattice, hydrogen bonds continue
to be present between individual molecules. But note
also that these bonds (as well as “hydrogen repulsion”)
seem to be contributing to the creation of the hexagonal
symmetry itself, as well as preventing the molecules from
getting as close together as they can in liquid water (e.g.
in the case of the two molecules sharing a double
hydrogen bond in the previous diagram).
In other words, the water molecules within an ice crystal
(more specifically, the center points of the molecules) are
“on average” slightly farther apart than they would be in
liquid water. This explains why water “expands” a little
as it freezes, why ice is therefore slightly less dense than
water, and hence why ice cubes float!
It is less clear why this process of “expansion of water in
preparation for crystallization” actually begins at 4 C,
but it is miraculous that it does so. Think about what
happens in the winter in a cold, shallow freshwater lake.
If water continued to contract from 4 C until it froze at
0 C, the water at the surface (where cooling first occurs)
would continue to contract and sink to the bottom,
eventually freezing the lake solid and killing all the fish
in the process. Instead, once the water temperature at
the surface drops below 4 C, the water actually
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expands, keeping it less dense than (and hence floating
above) the 4 C water immediately beneath it. A layer of
(denser) liquid water whose temperature is 4 C thus
continues to exist at the bottom of the lake even when
thick ice has formed at the surface, and, as a result,
aquatic life can survive the winter!
More will be discussed later on about water’s role in
fostering the earliest forms of life on the planet (and, for
that matter, about where all of Earth’s water came from
in the first place).