Bonding basics and Lewis dot structures:
Elements combine to form Compounds by “Bonding” to each other.
There are different kinds of Bonds: Metallic, Ionic and Covalent (some covalent bonds are called polar) the type of bond
determines many of the properties associated with the compound and you should know these trends.
The type of bond (and how many there are) determines the Shape of a molecule which is important in determining
other properties of the compound.
All bonding involves the Valence (outer shell electrons) and we will be using 2 standard methods of representing
bonding and shapes the “Lewis dot model” and the VSEPR model (valence shell electron pair repulsion model).
Lewis Dot Structures:
Octet Theory
(Chapter 8 of Glencoe’s “Chemistry Matter and Change, 2013”)
Compounds consisting of nontransition elements usually involve sharing of electrons between atoms so
that each atom has a completed outer shell. In the case of hydrogen the outer shell is the 1s orbital which
is complete with two electrons. For larger atoms it means that the s and p orbitals are filled with eight
electrons, two in the s orbital, six in the p orbital (s2p6). These are the electronic structures of the noble
gases. Thus the atoms, in bonding together to form a molecule, are approaching the stability of the noble
(or inert) gases. This applies to covalent bonding where the electrons are shared (such as in H 2) as well as
to highly ionic bonding, such as in NaCl, where the sodium ion has donated its outer electron to the
chloride ion.
Below are examples of atoms with the s and p outer electrons indicated with dots (each dot represents an
outer shell electron).
Two chlorine atoms can combine to form Cl2 and share two electrons as shown below.
The number of valence electrons can be found on your periodic table and follow a very predictable pattern for the
representative elements. The shape of a molecule can be based upon it’s central atom. The shapes we will need to know
are : Line, Linear, trigonal planer, tetrahedral, pyramidal, bent line and line (see page 263 of the text for good pictures
Orbital and dot diagrams and Molecular shape
Name: __________________ period: _____
Orbital diagrams and dot diagrams like those shown below can be used to show the electron configurations of atoms
and the shifts in electrons that occur during bonding. In the orbital diagram, each arrow represents an electron in a
particular orbital. In the dot diagram, each dot represents a valence electron
Carbon:
and it’s 4 valence electrons
Carbon 
This is the orbital diagram

s
 
p
p
p
or sp3
this is the dot diagram
Note that the orbital diagram represents the hybridization of the four orbitals of the valence shell: each orbital
(one s and 3 p) is half filled. All four orbitals are therefore available for bonding.
Such bonding is called sp3 bonding and allows for four bonds of the same character and generally results in a
tetrahedral shape for the molecule (or portion of a molecule) in which the carbon atom is bonded to four other atoms
(for example CH4). Bonding that involves three p orbitals (p3 bonding) generally results in pyraminal molecular shape.
Bonding involving two p prbitals (p2 bonding) generally forms a bent, non-linear molecule. Bonding involving one p
orbital (p bonding) results in a line molecule
Draw/fill in (valence) orbital and dot diagrams for the atoms below, and predict the type of bonding (sp3,
p3, p2, or p) characteristic of each and the approximate shapes of the molecules that would result from
such bonding.
1.
Oxygen
orbital
Type of bonding ________
2.
Fluorine
orbital
Type of bonding ________
dot structure
shape: _________
dot structure
shape: _________
3. Nitrogen
orbital
Type of bonding ________
4. Silicon
orbital
Type of bonding ________
5. Phosphorus
orbital
Type of bonding ________
dot structure
shape: _________
dot structure
shape: _________
dot structure
shape: _________
Now try it for a compound and make a lewis structure for the molecule:
6. Oxygen difluoride (OF2)
orbitals
Type of bonding ________
7.
Nitrogen trifluoride (NF3) orbitals
Type of bonding ________
dot structure
shape: _________
dot structure
shape: ________
Lewis Dot Structures
This is a method of trying to represent a three dimensional structure in two dimensions so it is not perfect.
General guidelines:
Represent only the valence electrons:
Atoms in column 1 have 1 dot, column 2 has 2, column 13 has 3, column 14
has 4, etc.
Draw the skeleton structure first try to be as symmetric as possible with the
most electronegative element in the center, and/or the element that forms the
most bonds in the center.
Figure out how many electron there are to share and position them so that
each non-metal atom gets 8 electrons around it. (elements like H, Li, Be, Na, K are
represented with only one or two dots). Remember for ions there will be more or
less electrons to share.
Try to do it with a single pair of electrons being shared first, if necessary
share 2 or 3 pairs when needed.
A Coordinate bond is when both electrons come from the same element.
If the bond is ionic draw brackets around the more electro-negative element. This
is to show “ownership” of the electron pair. And indicate which is positive and negative with the appropriate sign and
number of electrons gained/lost.
Put brackets completely around polyatomic ions and indicate the charge
Examples:
Ionic:
NaCl
Molecular
Cs2O
CH2Cl2
HNO3
H2 O
(has a coordinate bond and shows resonance structures)
Or it can be shown as
IONS: Put brackets around the entire structure to show the ionic nature of it.
OH-
PO4-3
Shapes of molecules: (due to S and P orbital hybridization) due to the 4 bonding areas around a representative
element.
Column 1
Line
Column 2
Linear
Column 13
Triangular planer
(two points determine a line)
(three in a row)
(three as far apart from the center as they can get)
Column 14
Tetrahedral (the new shape you need to know)
Column 15
Pyramidal
Column 16
Bent Line (can’t be straight and be filling two parts of a p orbital)
Column 17
Line (two points determine a line)
(three legged stool)
Column 18 These noble things don’t combine
ELECTRONEGATIVITY Table in your book is Page 194 (periodic table chapter)
The tendency for atoms of the element to attract electrons when they are chemically combined with another
element. The highest score is generally 4.0 for fluorine and there is no calculated value for the noble gases since they
generally don’t form compounds. This value will allow us to calculate bond types: Covalent (polar or non-polar) and
Ionic
Difference in electronegativity (aprox.)
Bond type
0.0 – 0.4
colvalent (nonpolar)
0.4 – 1.0
covalent (moderately polar)
1.0 – 2.0
Covalent (very polar)
greater than 2.0
IONIC
Most people say a difference greater than 1.66 is considered IONIC
The shape of a molecule has a lot to do with its polarity (lopsidedness)
Some shapes like pyramidal and bent line usually form polar molecules
Water is polar due to its shape:
Hydrogen a single electron in an “S” orbital and will form a single bond
Oxygen has 6 valence electrons (8 total) that fill the “S” orbital and
partially fill the “P”. The two available locations are perpendicular to each
other so a “bent” molecule is formed. Oxygen likes to form 2 bonds.
Element
1S
2S
2Px
2Py
2Pz
O





H





H















VSEPR: Valence Shell Electron Pair Repulsion theory
Is the fancy way of showing Hund’s rule (spread out if you can)