Section 3.1 (pg. 78-84)
Homework: Lewis Symbols
Pg. 82 #2 – 4 Pg. 84 # 2, 4, 5, 7-10
Objectives:
1) Define valence electron, electronegativity, and ionic bond
2) Use the Periodic Table and Lewis structures to support and explain ionic bonding
3) Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions.
Bonding Theory: Valence Electrons & Orbitals
 To describe where electrons exist in the atom, chemists
created the concept of an orbital.
 Orbital – region of space around an atom’s nucleus where an electron
may exist
 An “orbital” is not a definite race track, it is a 3-D space that defines
where an electron may be (like a rain drop in a cloud)
 For bonding study we are only concerned with an atom’s valence
orbitals (the volume of space that can be occupied by electrons in an
atom’s highest energy level)
 WHY? Bonding only involves valence e-’s because lower energy levels are held so
strongly by their positively charged nucleus
FYI Read pg. 78-79 for the history on Bonding Theory
Bonding Theory: Valence Electrons & Orbitals
 According to bonding theory, valence electrons are
classified in terms of orbital occupancy.
(0 = empty, 1 = half filled , 2 = full)
 An atom with a valence orbital that has a single electron
can theoretically share that electron with another atom

Such an electron is called a BONDING ELECTRON
 An atom with a full valence orbital (2 e-’s), repels nearby
orbitals and wants to be alone

Such a pairing is called a LONE PAIR
The Four Rules of Bonding Theory
1.
The first energy level has room for only one orbit
- can only hold 2 e-’s max
2.
Energy levels above the first have room for
four orbitals = 8 electrons max
2e2 p+
He
- Noble gases have this structure; their lack or reactivity indicates
that eight electrons filling a valence orbital is very stable
(Remember the OCTET RULE)
FYI: Only C, N, O,
and F atoms
always obey the
octet rule when
bonding
8e8e2e18 p+
Ar
2e-
2e-
2e-
2eEXCEPTIONS:
B = stable with 6 valence e- (3 orbitals)
P = stable with 10 valence e- (5 orbitals)
S = stable with 12 valence e- (6 orbitals)
2e-
The Four Rules of Bonding Theory
3. An orbital can be unoccupied, or it may contain one or two
electrons – but never more than two (Pauli Exclusion Principle)
4. Electrons “spread out” to occupy any empty valence orbitals
before forming electron pairs
“Aluminum has three half-filled
orbitals and one vacant orbital.”
How would you describe Sulfur?
Never more than
2e- in an orbital
Atomic Models: LEWIS SYMBOLS
(aka Lewis Dot Diagrams, Electron Dot Diagrams, LDD, Lewis Models
• Named after Gilbert Lewis who is responsible for the Octet Rule. He
reasoned that all atoms strive to be like the nearest noble gas.
• Used dots or ‘x’ to represent the valence electrons
• The inner electrons and the nucleus are represented by the element symbol
How to draw Lewis Symbols:
1.
2.
3.
4.
Write the element symbol
Add a dot to represent each valence eStart by placing valence e-’s singly into each of the four valence orbitals (4 sides)
If additional e-’s need to be placed, start filling each of the orbitals with a second e- up to 8
Q: Which element has 4 bonding e-’s?
Which has 3 lone pairs and 1 bonding e-?
Practice
 Draw the Lewis Symbols for the elements in Period 3
 For each one indicate how many lone pairs or bonding
electrons are present
 It is important to remember that the Lewis symbols do not mean that
electrons are dots or that they are stationary.
 The four sides represent the four orbitals that may be occupied by electrons;
it is a simplistic 2-D diagram of a complex 3-D structure
Electronegativity
-A measure of the force that an atom exerts on electrons of
other atoms; (the “pull” on bonding electrons)
-Each atom is assigned a value between 0.0 – 4.0; the larger the number
the greater the “pulling” force
-Example: Fluorine has an EN = 4.0 and francium has an EN = 0.7
-This means fluorine wants to pull on other electrons very strongly
-This means francium doesn’t want to pull on other electrons
Q: Does lithium (EN = 1.0) want to lose or gain an electron to be stable?
Q: Does fluoride (EN = 4.0) want to lose or gain an electron to be stable?
Do you see any relation to their electronegativity numbers?
So how do we assign each atom an electronegativity number?
a) The farther away from the nucleus that electrons are, the weaker
their attraction to the nucleus
.
EN = 0.8
Cesium's valence electrons are not
held as tightly by its nucleus
because the atom is larger
EN = 2.6
b) Inner electrons shield valence electrons from the attraction of
the positive nucleus
1 e8e8e2e19p+
K
EN = 0.8
Potassium’s valence electrons are not
attracted to its nucleus as much as
Nitrogen’s valence electrons because their
are more inner electrons present in K
EN = 3.0
5e2e7p+
N
c) The greater the number of protons in the nucleus, the greater the
attraction for more electrons
14p+
Si
EN =1.9
Bromine has more protons (+ charge)
which attracts the negative charge of
electrons more so than silicon’s 14 protons
EN = 3.0
35p+
Br
Electronegativity
In this 3-D image, the
electronegativity scale is
vertical.
Q: What is the EN trend
within a period and a
group?
Q: Which element has the highest
EN? Give three reasons why?
Why do we care about electronegativity??
 Imagine that two atoms, each with an orbital containing one bonding
electron, collide in such a way that these half-filled orbitals overlap.
 As the two atoms collide, the nucleus of each atom attracts and
attempts to “capture” the bonding electrons of the other atoms
 A “Tug of War” over the bonding electrons occurs
 Which atom wins?
 By comparing the electronegativities of the two atoms we can
predict the result of the contest = 3 different types of bonds result
Covalent Bonding
 Both atoms have a high EN so neither atom “wins”
 The simultaneous attraction of two nuclei for a shared pair of
bonding electrons = covalent bond
Cl2 = diatomic
 EN difference can be zero = Cl – Cl
EN = 3.2 EN = 3.2
 EN difference can be small = H - Cl
EN = 2.2 EN = 3.2

This is called a polar covalent bond – because
one side pulls on the electrons more but we
will learn more about this in Section 3.3
Ionic Bonding
 The EN of the two atoms are quite different
 The atom with the higher EN will remove the bonding
e- from the other atom
electron transfer occurs
 Positive and negative ions are formed which electrically
attract each other
EN = 0.9
EN = 3.2
Metallic Bonding
 Both atoms have a relatively low EN so atoms share valence
electrons, but no actual chemical reaction takes place
 In metallic bonding:
a) e-’s are not held very strongly by their atoms
b) the atoms have vacant valence orbitals
- This means the electrons are free to move around between
the atoms and the (+) nuclei on either side will attract them
Analogy: The positive nuclei are held
together by a glue of negative e-’s
Metallic bonding visual
This diagram
represents Mg atoms
that have released
their electrons and are
embedded in a sea (or
glue) of electrons.
Note: These metal atoms don`t have to be in a particular arrangement
to attract each other therefore they are flexible, malleable and
ductile = useful alloys (Brass, Stainless Steel, etc.)
Summary of Bonding Theory:
Chemical Bond = competition for bonding electrons
1) Atoms with equal EN = electrons shared equally
If both have high EN = covalent bond (equal = non-polar)
If both have a low EN = metallic bond
2) Atoms with unequal EN = covalent bond (unequal = polar)
3) Atoms with unequal EN = ionic bond
metallic
The nature of chemical bonds changes
in a continuous way, creating a broad
range of characteristics.
PRACTICE
Copy pg. 84 – Bonding Theory Summary into your Notes
Pg. 82 #2 - 4
Pg. 84 # 2, 4, 5, 7-10