Chemical Bonds Reading
Mendeleev’s organization of the elements in the periodic table was based entirely on their physical properties.
The modern periodic table is based on the structure of the atoms. Electrons and protons in the atom create
forces of attraction and repulsion (Coulomb’s Law). These forces influence the internal structure and reactivity
of the atom. Looking at these forces can help you understand the trends in the periodic table and the nature of
chemical bonding among atoms.
Chemical Reactivity
An element is considered
“reactive” when it goes
through a chemical reaction to
make a molecule or
compound. Some elements
are highly reactive, while
others are not. For example,
Figure 1 compares the
difference in reactivity
between oxygen and neon.
Notice that oxygen reacts
readily with magnesium, but
neon does not. Why is oxygen so reactive while neon is not? How much an element reacts depends on the
electron configuration of its atoms.
When two atoms get together, sometimes they will try to join with each
other to form a chemical bond. There are many things that control
chemical bonds. But the most important thing that controls chemical
bonds is the electrons around an atom – more importantly, the
electrons in the outermost layer of the atom.
You will remember that the organization of the periodic table shows us
lots of information about the properties of the elements in the table.
One of these properties is the atomic radius. As you go down the
periodic table from period to period, the atomic radius gets bigger. The
reason for this increase in size is because the electrons make a new orbital layer for each new period in the
periodic table. It is like people sitting in the bleachers at a basketball game. The first people come in and sit in
the first row. When the first row is filled up, they start to sit in the second row. When it is filled, they sit in the
third and so on until all the bleachers are filled. It works the same for an atom. As you add electrons
(increasing atomic number) they fill the atom’s “bleachers” called orbitals. When the first orbital is filled, they
start to fill the second, and so on. It is the last row of the “bleachers” being filled that is most important when
looking at chemical bonds.
Periodic Table Reveals an Atom’s Number of Valence Electrons
It is easy to find out how many valence electrons an atom has. All you have to do is check the periodic table.
The elements in the outermost orbitals (highest “bleachers”) are most often involved in chemical bonding. For
this reason, electrons in the outermost orbitals are called valence electrons. In period 1, there are 2 valence
electrons. In Periods 2 and 3, there are 8 valence electrons. Once you get down to periods 4 through 7, the
number of valence electrons
is a little more difficult to
determine, but that will be
covered in greater detail if
you take a regular chemistry
class.
To find the number of
valence electrons in an atom,
all you have to do is look at
the periodic table. There are
two ways to find the number of valence electrons. (For the purposes of this class, we will not be looking at the
elements in the Transition Metal group to make this section of the class easier.)
The first way to find the
number of valence of
electrons is to look at
the top of the periodic
table. On the periodic
table you will use for
the CST, there are two
numbers at the top of
each group. The first is
the group number
(going from 1-18). The
second is a number
and a letter (1A, 2A, up
to 8A – not looking at
the transition metals). This second number tells you the number of valence electrons in each group. For
example, Hydrogen is in Group 1 and has a second number of 1A. This means that Hydrogen has 1 valence
electron. Second example: Phosphorus is in Group 15 on the periodic table, but is has the number 5A above
it. That means that Phosphorus has 5 valence electrons.
The second way to find the number of valence electrons for an element is to
count across in a period (skipping over the transition metals) until you get to
the element you reach the element you want to know about. For example, if
you want to find the number of valence electrons in Oxygen, you start at the
beginning of Group 2 (Lithium) and count the elements in that group until you
reach oxygen (6th element over). Because Oxygen is the 6th element in Period
2, it has 6 valence electrons.
Chemical Bonding and Types of Chemical Bonds
Looking at the properties and trends in the periodic table can help us understand what is happening with each
element, but it is the combining of elements into molecules and compounds that is what is really fun about
chemistry. Understanding how and why elements get together can help explain why there are so many
different things in the world with less than 100 stable elements in the periodic table. It has to do with
electrostatic forces of attraction and repulsion.
When two atoms are far apart, the only force of attraction is between the protons of one atom and the electrons
of the same atom. As long as the atoms are far apart, there are no other forces large enough to have any
major effect. As the atoms approach each other, however, other forces become important. There are the
repulsive forces between the electrons of one atom and the electrons of the other atom. There are also the
repulsive forces between the protons of the two atoms. But there are also new attractive forces. The protons
of one atom can attract the electrons of the other atom, and vice versa. If these new attractive forces are
greater than the repulsive forces, the two atoms stick together to form a molecule or compound. The result is a
chemical bond. In a chemical bond, some of the valence electrons of the atoms move between the atoms
which “glue” the atoms together. How they move helps control the type of chemical bond and how strong the
atoms are bonded together.
The valence electrons of the two atoms can either be shared
between the atoms or the valence of one atom can move
over to the other atom making both into ions. The
electronegativity of each of the atoms is important in
determining whether the electrons will be shared or
completely moved. Remember that the electronegativity is a
measure of the ability of an atom of an element to attract
electrons toward it. The elements on the right hand side of
the periodic table have a larger electronegativity than those on the left hand side and the elements on the top
of table have a larger electronegativity than those on the bottom. The difference in the electronegativity of the
two atoms being bonded will control what type of chemical bond is made.
When the electronegativity of the two atoms is about equal, the electrons will be
shared between the two atoms. When the electrons are shared it is called a
Covalent Bond. Depending on how close the electronegativity of one atom is to
the other will determine whether the electrons are shared equally between the two
atoms or if one atom will have the electrons for a longer period of time. If the
electronegativities are about equal, the sharing will be equal and the bond will be
purely covalent. If the electronegativities are close but one atom is slightly
stronger, the sharing will not be equal. This unequal sharing is called a Polar
Covalent Bond.
When the electronegativity of the two atoms are very different (one
large and the other smaller), the electrons will not be shared between
the two atoms. In this case, the atom with the larger electronegativity
will take the electron from the atom with the lower electronegativity
which will make ions out of both atoms. The atom that loses the
electron will become a positively charged ion called a Cation. The
atom that gains the electron will become a negatively charged ion
called an Anion. Once both atoms are ions, they are attracted to each
other by electrostatic attraction force which makes the chemical bond
called an Ionic Bond. When many Ionic bonded compounds are near
each other, the positive charges from one compound can attract the
negative charges of another compound. The attractive force between
the two compounds will cause them to organize themselves into
repeating patterns of positive and negative ions held together by
electrostatic forces. This repeating pattern is called a crystal lattice.
The energy holding the crystal lattice together is called the crystal
lattice energy.
Valence Electrons and the Octet Rule
Just because two atoms could bond together, do all combinations of
atoms actually bond? How is it possible to predict whether or not two or
more atoms will bond? Earlier we learned that the noble gases do not
form many compounds. Since the electrons in the outermost orbital
control the reactivity of an atom, it would appear that the way the
electrons are arranged (also called the electron configuration) in the
noble gases is very stable (won’t easily react). If other atoms could either
add or subtract electrons so that their outer orbital could look like that of a
noble gas, it would also be more stable. This thought led scientists to
come up with the “Rule of Eight” or the “Octet Rule.” This rule helps predict which atoms whether or not an
atom will bond with one or more other atoms.
The assumption of the Octet Rule is that the elements in groups 1, 2, 16 and 17 and other non-metallic
elements in the periodic table usually gain or lose electrons through the formation of either ionic or covalent
bonds resulting in eight electrons in the outermost shell. The rule holds true for a large number of molecules,
although there are many exceptions.
For example, if an element in group 1 attempts to bond with an element in group 17, it might go something like
this: The element in group 1 will lose an electron to the element in group 17 (ionic bond). When this happens,
the element in group 1 will now have 8 electrons in its outermost shell (the old outer shell is now empty – no
electrons). Additionally, the element in group 17 will add an electron to its outermost shell bringing the total
number of “valence” electrons to 8. Both atoms now meet the octet rule. (example – Sodium and Chlorine)
Another example: If an element in group 16 wants to
bond with an atom from group 1, it won’t work with just
one atom from each group. Here’s why. The atom
from group 1 will lose its electron and satisfy the octet
rule, but the atom from group 16 will only have 7
electrons – which doesn’t satisfy the octet rule.
However, if one atom from group 16 meets up with 2
atoms from group 1, the octet rule will be satisfied for
all three atoms. The two atoms from group 1 will lose
their electrons – satisfying the rule. The atom from
group 16 will gain 2 electrons – also satisfying the rule.
The 3 atoms will bond. (example – 2 hydrogen and 1
oxygen)
Remember, the octet rule is a general guide that does
not explain all the different combinations that actually
happen in real life. Molecules do not “obey” the octet
rule or any other rule. The concept is used as a
scientific tool or “rule of thumb.” There are some
atoms, especially phosphorus and sulfur that do not always follow the rule, but for the most part, atoms will
bond with other atoms so that they have 8 electrons in their outermost valence electron shell – either by
gaining or losing electrons during the bonding process.
Covalent Bonds and Biological Molecules/Compounds
Organic and biological molecules are made mostly of the four “organic nonmetal” elements – Nitrogen, Carbon, Hydrogen and Oxygen (NCHO – or
“Nacho”). These elements share valence electrons to form bonds to that
the outer electron shells are filled and are like those of the nearest noble
gas element (octet rule). For example nitrogen has 5 electrons so it can
bond with three hydrogen atoms. Carbon has four valence electrons and
combines with hydrogen, nitrogen and oxygen in various combinations to
make up a majority of the biological compounds we see today.
The bonds created by these elements are primarily covalent and polar
covalent (shared electrons). This is seen by comparing the
electronegativity values for each of these elements. The great variety of
combinations of carbon, nitrogen, oxygen and hydrogen make it possible,
through covalent bond formation, to have many hundreds (even thousands)
of different compounds from these few elements.