Lesson 8.1 Review of Lewis Structures
Suggested Reading

Zumdahl Chapter 8 Section 8.10
Essential Question

How are Lewis structures used to show bonding in chemical
compounds?
Learning Objective



Differentiate between lone pairs and bonding pairs.
Deduce the Lewis structures of molecules and ions for up to four
electron pairs on each atom.
Draw Lewis structures for compounds whose central atom does not
obey the octet rule.
Lewis Dot Symbols
Recall that a Lewis electron-dot symbol is a symbol in which the electrons in
the valence shell of an atom or ion are represented by dots placed around the
letter symbol of the element. The table below shows the Lewis symbols for
several main-group elements. Note that the dots are placed one to each side
of a letter symbol until all fours sides are occupied. Note how Boron shows
that the pairing of dots does not necessarily correspond to the pairing of
electrons in orbitals in the ground state.
Watch the following tutorial to review Lewis dot symbols:
YouTube Video
https://www.youtube.com/watch?v=EOkzt42DvE4
Lewis Dot Symbols & Ionic Bonds
In an ionic bond, one atom looses all its outer electrons (leaving behind a filled
inner shell) while another atom gains electron(s) to fill its valence shell. The
two ions attract each other according to Coulombic interactions. Lewis dot
symbols can be used to illustrate what occurs during the formation of an ionic
bond.
Coulombic Interactions: Attractions between opposite charges or repulsions
between like charges that grow stronger as the charges become closer to
each other.
For example, consider sodium chloride.
The Lewis structure for the salt NaCl, shows two ions which have their (Now)
outer shells of electrons filled with a complete octet. In the case of the sodium
cation, the filled shell is the outermost of the 'core' electron shells. In the
chloride ion, the outer shell of valence electrons is complete with 8 electrons.
The two ions are bonded by an ionic bond.
Consider the compound magnesium oxide: MgO.
Note that in both cases, the positive ions are metals and the negative ions are
non-metals. This is generally true for ionic compounds Recall that in general
metals have lower ionization energies and form cations while non metals have
high ionization energies (and high electron affinities) and form anions.
1
2
Li+
Ba2+
Na+
Mg2+
K+
Ca2+
Rb+
Sr2+
Cs+
Ba2+
3
4
5
6
7
Al3+
Lewis Dot Symbols & Covalent Bonds
Covalent bonds involve the sharing of electrons to form 'bonding pairs' of
electrons between atoms. Usually, when this occurs each atom ends up
having eight electrons around it. This is the so-called octet rule.
Take Cl2. Each Cl atom has one unpaired electron
. It's unlikely that an
ionic type bond will form since the two atoms are identical and such a transfer
would leave one chlorine with only 6 electrons. More likely, they will share a
pair of electrons so that each can 'think' it has 8 electrons in its valence shell.
The other electrons around the chlorine atoms are referred to as lone pairs.
Thus, each chorine has one bonding pair of electrons and three lone pairs of
electrons. Mostly, this type of bonding occurs between non-metal
atoms. Consider PCl3:
Note: the lone pairs on the chlorines are hidden to emphasize the electron
configuration around the phosphorous. Phosphorous normally has five
electrons in its valance shell. The P shares one of its unpaired electrons
with each of the three Cl atoms to form three covalent bonds. All four atoms
now have a completed octet.
It is not always necessary to use two dots to represent the pair of electrons. A
dash can also represent the pair of electrons, especially in bonds. Consider
the following examples
or using a — to represent a bonding pair
Watch the video to review the procedure for drawing Lewis structures for
covalent compounds.
YouTube Video
https://www.youtube.com/watch?v=nw3xVVmEAU8
Exceptions to the Octet Rule
There are exceptions to the octet rule. For example, H has only a 1s orbital
occupied. The other orbitals are all higher in energy and not involved in the
chemistry of hydrogen.
Consider water H2O.
There are only two electrons in the valence shell of hydrogen but that's all it
takes to fill the first shell. Hence, the octet rule doesn't apply.
Other compounds where the octet rule is not obeyed include some boron
compounds. Boron has only three unpaired electrons to
share.
Here, Boron has only 6 valence electrons. It cannot complete its octet and
remain a neutral compound. BF3 can combine with an F- ion to complete its
octet as follows.
One of the lone pairs on the fluoride is shared with the boron and forms a
bond.
Expanded Octets (Hypervalence)
Elements in row 3,4,... have access to d orbitals (unlike row 2 elements that
have only s and p orbitals) and can therefore have more than 8 electrons
around them in certain compounds. Recall that P has a normal valence that
includes 3 unpaired electrons and one lone pair. If that lone pair is divided into
two extra unpaired electrons to be used in creating bonds, we then have a
hypervalence of 5.
By expanding the valence of P, we change it from Valence=3 with one lone
pair to Valence=5 with no lone pairs. PCl5 is one such compound that has P
with a hypervalence of 5 and no lone pairs.
Other elements, like S can have more than one hypervalence. S has a
normal valence of 2 unpaired electrons with two lone pairs. If we split the lone
pairs one at a time to create extra unpaired electrons for bonding, we can
create hypervalences of 4 with one lone pair and 6 with no lone pairs. The
compounds are SF2, SF4, and SF6. Try drawing these Lewis structures out.
Note that the number of lone pairs on the S decreases by one as the number
of unpaired electrons goes up by 2.
At one time, it was believed that the octet rule was the overriding factor in
determining bonding. Using this logic, it was once thought that the Noble
gases (He, Ne, ...) were totally unreactive since they already have a filled
octet. Using the concept of hypervalences, we see that we can expect some
compounds of noble-gases. We have in recent years produced compounds of
Xe as follows:
XeF2, XeF4, and XeF6, XeO4. All four of these compounds are using a
hypervalent state of the Xenon.
HOMEWORK: Book questions pg. 385 questions 67 & 68
Practice exercises 8.6-8.9, 8.11