Wednesday, Feb. 19th: “A” Day
Thursday, Feb. 20th: “B” Day
Agenda
Collect homework: pg. 198: 1-14
Collect lab: “Nonmetal Reaction”
Sec. 6.1 quiz: “Covalent Bonds”
Begin Section 6.2: “Drawing/Naming Molecules”
Valence electron, Lewis structure, unshared pair, single
bond
In-Class: Practice pg. 202: #1,2
Practice pg. 203: #1,2
Homework: Concept Review
Homework
Pg. 198: #1-14
Lab: “Non-metal Reaction”
Questions, problems?
Turn in
Section 6.1: Quiz
“Covalent Bonds”
You may use your notes, your book, and your
10:00 lab partner on the quiz….
Question #2 IS NOT in
section 6.1. The answer is
“B”
When you’re done, pick up the notes for today…
Sec. 6.2:
“Drawing and Naming Molecules”
Both ionic and covalent bonds involve valence
electrons, the electrons found in the
outermost energy level of an atom that
determine the atom’s chemical properties.
Determining Number of
Valence Electrons
Review
How many valence electrons are in the following
elements?
Ca
P
I
C
Li
Lewis Electron-Dot Structures
In 1920, G.N. Lewis, an American chemist,
came up with a system to represent the valence
electrons of an atom.
This system, known as
electron-dot diagrams or
Lewis structures, uses dots to
represent valence electrons.
Lewis Structures Model Covalently
Bonded Molecules
Lewis structures: a structural formula in
which electrons are represented by dots. Dot
pairs or dashes between two atomic symbols
represent shared electrons in covalent bonds.
A Lewis structure shows only the valence
electrons, NOT all of the electrons in an atom.
Na•
•Ca•
Lewis Structures Show Valence Electrons
As you move from left to right across a period,
each element has an additional valence electron.
These valence electrons are represented by dots
on each side of the element’s symbol.
Lewis Structures Show Valence Electrons
You do not begin to pair dots (valence electrons)
until all four sides of the element’s symbol have a
dot.
Octet Rule
An element with an octet of valence electrons,
such as a noble gas, has a stable configuration.
The tendency of bonded atoms to have octets
of valence electrons is called the octet rule.
Lewis Structures Show Valence Electrons
When two chlorine atoms form a covalent bond,
each atom contributes one electron to form a
shared pair.
With this shared pair, both atoms can have a
stable octet.
Unshared pair: a non-bonding pair of electrons
in the valence shell of an atom; also called a lone
pair.
Lewis Structures Show Valence Electrons
Single bond: a covalent bond in which two
atoms share one pair of electrons.
The shared pair of electrons are replaced with
a dash, indicating a single bond.
(A dash counts for 2 electrons)
Rules for Drawing Lewis Structures
Pg. 201
There are 5 steps to follow to draw a Lewis
structure:
1. Determine the number of valence electrons in
each atom. Add them up to get the total
number of valence electrons in the molecule.
Rules for Drawing Lewis Structures
Pg. 201
2. Arrange the atoms
 Hydrogen atoms can only form 1 bond!
 Halogen atoms often bond to only one other
atom and usually at the end of the molecule.
 Carbon wants to be the center of attention and
is often placed in the center of the molecule.
 With the exception of carbon, the atom with the
lowest electronegativity is often the central
atom.
 Molecules also like to by symmetrical.
Rules for Drawing Lewis Structures
Pg 201.
3. Distribute the dots (valence electrons)
 Arrange the dots so that each atom (except
for hydrogen, beryllium, and boron) satisfies
the octet rule (8 dots).
4. Draw the bonds
 Change each pair of dots that represents a
shared pair of electrons to a long dash.
5. Verify the structure
 Count the number of electrons around each
atom. Make sure the octet rule is satisfied
and that you have the correct amount of dots
that you determined in Step 1.
Sample Problem A, Pg. 202
Draw a Lewis structure for CH3I.
1. Determine the total number of valence electrons
in each atom and in the molecule.
4 + 1 + 1 + 1 + 7 = total number of dots 14
2. Arrange the atoms. Put carbon in the center with
the other atoms around it.
Sample Problem A, Pg. 202
3. Distribute the dots.
 Put one shared pair of electrons between each
of the bonded atoms.
 Distribute the remaining electrons, in pairs,
around the remaining atoms to form an octet
for each atom.
Sample Problem A, Pg. 202
4. Change each pair of dots that represents a
shared pair of electrons to a long dash.
H
H C I
H
5. Verify the structure! Count the electrons and
make sure you have the correct amount.
(Each dash counts for 2 electrons.)
Additional Examples
Draw the Lewis structures for the following
molecules:
1. SCl2
2. AsF3
3. SiH4
4. CHF3
Lewis Structures for Polyatomic Ions
Lewis structures can be drawn for polyatomic
ions as well.
If the charge on the polyatomic ion is negative,
you must add electrons.
If the charge on the polyatomic ion is positive,
you must subtract electrons.
The entire Lewis structure is put in [brackets]
to show that the charge of the ion is distributed
over all of the atoms.
[ ]
Sample Problem B, Pg. 203
Draw a Lewis structure for the sulfate ion, SO42-
1. Determine the total number of valence
electrons in each atom and in the molecule.
Add two additional electrons to account for the
2− charge on the ion.
6 + 6 + 6 + 6 + 6 = 30 valence electrons
Plus 2 extra (for the 2- charge) = 32 total
Sample Problem B, Pg. 203
2. and 3. Arrange the atoms and distribute the
32 dots so that there are 8 dots around each
atom.
(Sulfur has the lower electronegativity, so it’s
placed in the center, surrounded by the 4
oxygen atoms. Plus, this is symmetrical.)
Sample Problem B, Pg. 203
4. Draw the bonds.
Change each shared pair of electrons to a long
dash. Place brackets around the whole
structure and a 2 charge outside the bracket
to show that the charge is spread out over
the entire ion.
Additional Examples
Draw the Lewis structures for the following
polyatomic ions:
1. OH-
2. NH4+
3. H3O+
4. ClO-
In-Class
Practice, pg. 202: #1,2
Practice, pg. 203: #1,2