Electron Dot Diagrams and
Bond Geometry
Lesson 6.1
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
I. Electron Dot Diagram
A. An electron dot diagram is a representation of an atom in which
only the valence electrons are illustrated.
1. Unpaired electrons are shown as single dots.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
2. There are many conventions as to the order of dot
placement around an element symbol.
a. For the purpose pf this class, our convention will be
to place the first valence electron at the top of the
symbol, and then fill in dots in a counter clockwise
sequence, until the top is reached again.
b. Subsequent electrons will then be filled in
counterclockwise order to complete the full octet.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
II. Lewis Dot Structures
A. Lewis structures (also known as Lewis dot diagrams, electron
dot diagrams, "Lewis dot formulas", Lewis dot structures, and
electron dot structures) are diagrams that show the bonding
between atoms of a molecule and the lone pairs of electrons that
may exist in the molecule.
Gilbert Lewis (1875-1946)
1. The Lewis structure was named after Gilbert N. Lewis, who
introduced it in his 1916 article The Atom and the Molecule.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
III. Drawing Lewis Dot Structures
A. Representation of Electrons
1. Non-bonding Electron Pairs
a. Non-bonding electrons are represented by dots.
b. Non-bonding electrons belong to one atom alone,
that is, they are not considered to be part of bond
connecting the atoms.
c. In the illustration above, the nitrogen atom has 1
pair of non-bonding electrons.
2. Bonding Electron Pairs
a. Bonding electrons pairs are represented by a
dash.
b. Bonding electron pairs are shared between the
connecting atoms.
c. In the example above, each hydrogen atom “sees”
two electrons (a full 1s2 orbital).
d. In the example above, the nitrogen atom “sees” 6
electrons from the 3 shared pairs, and 2 non-bonding
electrons for a full octet.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
B. Process for Drawing Lewis Dot Structures
1. Find the sum of the valence electrons of all of the atoms in
the polyatomic ion or molecule.
a. If the structure is a polyatomic ion, remove or add electrons
based on the final charge of the ion.
i. For cations, remove the number of electrons equal to the positive
charge (Na+ remove 1 electron from the total = 0 final valence
electrons).
ii. For anions, add the number of electrons equal to the negative
charge (O2- add two extra electrons to the total = 2 extra electrons
added to the 6 normal valence electrons for a total of 8 electrons)
2. Select the least electronegative atom and designate that
atom as the central atom in the molecule (not hydrogen).
a. Draw single electron pair bonds to the surrounding atoms from
the central atom.
b. Keep track of the number of electrons used in the bonded pairs.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
3. Draw non-bonding pairs of electrons around the peripheral
atoms.
a. Keep track of the number of electrons used in the non-bonding
pairs, and the bonded pairs.
4. Fill in the octet around the central atom using remaining
non-bonding pairs.
a. Keep track of the number of electrons used in the non-bonding
pairs, and the bonded pairs.
5. If there are not enough remaining electrons to complete the
central atom octet, form multiple bonds until all of the atoms,
including the central atom have complete valence shells.
a. The final number of electrons, involved in the unshared pairs and
bonded pairs must equal the initial number of summed valence
electrons.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
6. If multiple final Lewis Dot Diagrams are possible, assign
formal charges for each atom, and select the diagram that
results in the fewest formal charges.
a. For each atom, count the electrons in lone pairs and half the
electrons it shares with other atoms.
b. Subtract that from the number of valence electrons for that atom.
c. The difference is its formal charge.
i. Formal Charge = valence e- - (lone pair e- + 1/2 bond pair e-).
d. The best Lewis Dot Diagram has the lowest overall formal
charges, and places a negative formal charge on the most
electronegative atom.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
I. The Electron
A. J.J. Thomson (1856 –1940) used results from cathode ray tube
experiments to infer the presence of a particle smaller than the
atom – the electron.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
C. Example of the Process for Drawing a Lewis Dot Structure.
HCN
1. Find the sum of the valence electrons of all of the atoms in
the polyatomic ion or molecule.
HCN
H 1 e- Family IA
C 4 e- Family IVA
N 5 e- Family VA
10 e2. Select the least electronegative atom and designate that
atom as the central atom in the molecule (not hydrogen).
a. Draw single electron pair bonds to the surrounding atoms from
the central atom.
b. Keep track of the number of electrons used in the bonded pairs.
(4e-)
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
3. Draw non-bonding pairs of electrons around the peripheral
atoms.
(10 e- used)
4. Fill in the octet around the central atom using remaining
non-bonding pairs.
No additional e- are available to place around the central atom.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
5. If there are not enough remaining electrons to complete the
central atom octet, form multiple bonds until all of the atoms,
including the central atom have complete valence shells.
Not enough e- around central atom
Valence requirements satisfied.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
6. If multiple final Lewis Dot Diagrams are possible, assign
formal charges for each atom, and select the diagram that
results in the fewest formal charges.
Only one diagram is possible.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
D. Example of the Process for Drawing a Lewis Dot Structure.
NCS1. Find the sum of the valence electrons of all of the atoms in the
polyatomic ion or molecule.
NCS1e- extra for charge
N 5 e- Family VA
C 4 e- Family IVA
S 6 e- Family VIA
16 e2. Select the least electronegative atom and designate that atom as
the central atom in the molecule (not hydrogen).
a. Draw single electron pair bonds to the surrounding atoms from the
central atom.
b. Keep track of the number of electrons used in the bonded pairs.
(4e- used)
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
3. Draw non-bonding pairs of electrons around the peripheral
atoms.
(16 e- used)
4. Fill in the octet around the central atom using remaining
non-bonding pairs.
No additional e- are available to place around the central atom.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
5. If there are not enough remaining electrons to complete the central atom
octet, form multiple bonds until all of the atoms, including the central atom
have complete valence shells.
Not enough e- around central atom.
Valence requirements satisfied for all three options
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
6. If multiple final Lewis Dot Diagrams are possible, assign formal charges for
each atom, and select the diagram that results in the fewest formal charges.
Formal Charge = valence e- - (lone pair e- + 1/2 bond pair e-).
The best option is
Because it has the lowest overall formal charge and the negative
formal charge is assigned to the most electronegative atom.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
III. VSEPR Theory and Bond Geometry
A. Most molecules are three dimensional structures.
1. Lewis Dot structures can be used to represent fairly complex molecular
shapes.
2. The shape depends on the numbers of bonded pair (BP) and lone pairs
(LP) around the central atom.
3. The VSEPR (pronounced “vesper”) or Valence Shell Electron Repulsion
Theory provides a system for predicting the shape of molecules.
B. The following VSPER rules are used to predict the shapes of molecules:
1. The shape of a molecule depends on the repulsion between electron
pairs in the valence shell around the central atom of the molecule.
2. Electron pairs can either be bonded pairs (BP: single, double or triple
bonds are treated as “one” bonded pair), or lone pairs (LP) around the
central atom.
3. The electron pairs will repel each other so they are as far apart as
possible.
4. Lone pairs spread out and take up more room than bonded pairs so they
will push the bonded atoms closer together.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
C. VSEPR or AXnEm Notation
1. The number of bonded pairs
is determined from the Lewis
Dot Diagram for the molecule
and could be written using the
general VSEPR notation:
AXnEm.
a. A= central atom
b. Xe = # of bonded pairs (BP)
attached to the central atom
c. Em = # of lone pairs (LP) on
the central atom
D. The AXnEm Notation values
correspond to specific
geometries.
1. These geometries are based
on the interaction of electron
clouds and potential energy
reduction.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
IV. Process for Designating Molecular Shape
1. Draw the Lewis structure for the molecule.
2. Count the number of bond pairs (BP) and lone pairs (LP) around
the central atom.
3. Decide on the total number of electron groups (treat multiple
bonds as single electron groups).
4. Consider the locations of lone pairs and any distortions from
"regular" shapes.
5. Name the shape based on the arrangement of the bonding
atoms as outlined on the following slides.
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
a. If the total number of electron groups (bond pairs + lone pairs) is
TWO:
i. two bond pairs and no lone pairs THEN the molecule is linear
AX2E0
Ex. CO2
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
b. If the total number of electron groups (bond pairs + lone pairs) is
THREE:
i. three bond pairs and no lone pairs THEN the molecule is trigonal
planar
AX3E0
Ex.SO3
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
ii. two bond pairs and one lone pair THEN the molecule is bent or
V-shaped
AX2E1
ex. CCl2
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
c. If the total number of electron groups (bond pairs + lone pairs) is
FOUR:
i. four bond pairs and no lone pairs THEN the molecule is
tetrahedral
AX4E0
Ex. CH4
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
ii. three bond pairs and one lone pair THEN the molecule is trigonal
pyramidal
AX3E1
ex. NH3
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
iii. two bond pairs and two lone pairs THEN the molecule is bent or
V-shaped
AX2E2
Ex. H2O
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
d. If the total number of electron groups (bond pairs + lone pairs) is
FIVE:
i. five bond pairs and no lone pairs THEN the molecule is trigonal
bipyramidal
AX5E0
ex. PCl5
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
ii. four bond pairs and one lone pair THEN the molecule is a “seesaw” shape
AX4E1
ex.SF4
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
iii. three bond pairs and two lone pairs THEN the molecule is “Tshaped”
AX3E2
ex.CF3
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
iv. two bond pairs and three lone pairs THEN the molecule will be
linear
AX2E3
Ex. XeF2
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
e. If the total number of electron groups (bond pairs + lone pairs) is
SIX:
i. six bond pairs and no lone pairs THEN the molecule is octahedral
AX6E0
SF6
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
ii. five bond pairs and one lone pair THEN the molecule is a
square-based pyramid
AX5E1
Ex. ClF5
Lesson 6.1 Electron Dot Diagrams and Bond Geometry
iii. four bond pairs and two lone pairs THEN the molecule is square
planar
AX4E2
ex. XeF4