Ch. 12 Chemical Bonding
12.1 Ionic, covalent and polar
covalent bonds.
• A bond is a force that holds atoms
together.
• Ionic Bonding
a. When a metal reacts with a nonmetal,
electrons are transferred from a metal to a
nonmetal and an ionic compound is made.
Ionic Bonding
b. In ionic bonding, electrostatic attraction
holds atoms together.
+1
Na
-1
Cl
Covalent Bonding
a. Atoms make covalent bonds by sharing
electrons.
b. Electrons are attracted to the nucleus of
both atoms in the bond.
c. Nonmetals make covalent bonds.
Let’s
share!
H H
Electronegativity
•
Electronegativity is a measure of how
strongly an atom attracts electrons. Look
at the electronegativity chart on p403.
– Fluorine has an electronegativity value of
4.0.
– Hydrogen has an electronegativity value of
2.1.
– Difference in electronegativity =
4.0 – 2.1 = ________
Bond Types
• Determined by difference in electronegativity
values—absolute value of DEN
Pauling Electronegativity Values
Electronegativity
d. At left is a picture of hydrofluoric acid
(HF). At right is a picture of “The
Blob.”
e. F has a higher electronegativity than
H, so electrons are closer to F.
•
•
In a covalent bond, atoms have a
difference of electronegativity of 00.2. Electrons are shared equally.
In a polar covalent bond, atoms have
a difference in electronegativity of
0.3-1.7. Electrons are not shared
equally.
Electronegativity
• Exampe: Si—C
1.8
– Electronegativity of Si is ________.
Electronegativity of C is _______.
2.5
– Difference in electronegativity is
0.7
____________.
– Si – C bond is ( covalent /
polar covalent ) [circle one].
• Which attracts more electrons? ( Si /
C )
•
In the picture below,
– Label each molecule or compound as ionic
bonding, covalent bonding or polar covalent
bonding.
– For polar and ionic bonds, label the more
electronegative atom.
Covalent
Polar covalent
Ionic
Identify each of the following bonds as ionic,
covalent or polar covalent.
Bond
Is there a
metal and
nonmetal?
Electronegativity Bond Type
difference
H–H
No
2.1 – 2.1 = 0
Covalent
S–H
No
2.5 – 2.1 = 0.4
Polar
covalent
Na – I
Yes  ionic
Rb – N
B–P
H–P
ionic
Dipoles
•
•
In a polar molecule, one side has a partial
positive charge, and the other has a partial
negative charge.
A dipole moment is represented with an
arrow pointing to the negative side and the
Greek letter “delta” δ to show the partial
positive and negative charges:
Dipoles
•
Write the partial charges and draw the
dipole moment on Cl –I
12.2 Ionic Bonding
•
Ions
–
–
–
–
–
–
Metals ( lose / gain ) electrons.
Nonmetals ( lose / gain ) electrons.
Group 1 elements form ions with a charge of ____.
Group 2 elements form ions with a charge of ____.
Group 6 elements form ions with a charge of ____.
Group 7elements form ions with a charge of ____.
Ions
•
That’s interesting…but WHY???
– Atoms gain or lose electrons to get the
electron configuration of a noble gas.
– Noble gases have completely filled
energy levels, so they are very stable.
– He has a completely filled 2s sublevel.
– Other noble gases have filled s and p
sublevels.
Ions
•
Example: Li
– The electron configuration of Li is ______
– Li loses one 2s electron and becomes
Li+.
– The electron configuration of Li+ is _____
– Li + has the same electron configuration
as He.
Ions
•
The electron configuration of F– The electron configuration of F is
– F gains one 2p electron and becomes F– The electron configuration of F- is
– F- has the same electron configuration
as ______
Ions
•
What would happen if Li reacted with F?
–
–
–
•
Li gives an electron to F
Li + F  Li+ + FAnd they form _______ (write the formula of the
compound)
Ionic bonding and structures
–
–
LiF is packed together in a group in order to
maximize attractions of the cations and anions.
It makes a hard, tight crystal.
The size of ions
•
Which is larger, Na or Na+? Why?
–
Na. Na loses a 3s electron, and then only
has electrons in the n=2 level. n=2
orbitals are smaller than n=3 orbitals.
Na
Na+
F
F-
The size of ions
•
Which is larger, F or F-? Why?
–
•
•
•
F- because it gains electrons.
Which is larger F- or Na+? Why?
F-. They both have the electron
configuration of Ne, but Na+ has more
protons (a stronger + charge), which
pulls electrons closer.
Which is larger, Ca or Ca2+? Why?
Ionic Size
Taken from: http://www.chem.umass.edu/people/botch/Chem121F06/Chapters/Ch15/IonicRadii.jpg
12.3 Lewis Structures
•
The octet rule: Sharing of electrons
usually occurs so that atoms acquire
the electron configurations of noble
gases (1s2 or ns2np6)
Lewis dot structures
– The element symbol
represents the core electrons.
– Dots to show the
valence electrons.
Lewis dot structures
•
•
•
•
1. First, write the symbol for the
element.
2. Imagine the molecule has
four sides (but don’t draw the
“x”)
3. Draw one dot at a time in
each empty section.
4. You should only have pairs
of e- if there are no empty
sections.
Cl
Lewis dot structures
•
•
The “paired” electrons
cannot usually make bonds.
The three “unpaired” electrons
can make bonds.
Arsenic
• Dot Diagram?
• In order to get the configuration of a noble gas,
how many bonds will arsenic form? 3
• Dot Diagram for Hydrogen?
H
• In order to get the configuration of a noble gas,
how many bonds will H form? 1
Arsenic Trihydride, AsH3
Lewis dot structure
H
As
H
H
Structural formula
H
As
H
H
Drawing Lewis Dot Structures
1. Use the molecular formula to find the
total number of valence electrons
•
1+1+1+ 5=8
H H
H As
Drawing Lewis Dot Structures
2. Draw the symbols of each element.
Draw the backbone/skeleton
structure
– Terminal atoms go around
the central atoms.
– Least Electronegative is usually
central
– Hydrogen is always terminal.
– Carbon is usually central
– Oxygen and Halogens are usually
terminal
Drawing Lewis Dot Structures
3. Draw the valence
electrons
4. Make single bonds with
pairs
Drawing Lewis Dot Structures
5. Each pair represents a
bond.
6. Count electrons again
and make sure you get
the same number (8).
Single bond
Structural Formula
• The structural formula
is drawn with a “–“ line
to represent the bond.
• Double and triple bonds
–
–
Double bonds can be
formed by sharing two
pairs of electrons
Triple bonds involve
sharing three pairs of
electrons.
O
O
N N
Try the Dot Diagrams for each
of the following:
O2
O O
O
O
H2O
H O H
Each pair of
shared
electrons can
be
represented
O
Hby a line.H
BCl3
SiI4
Cl
Cl B Cl
I
I Si I
I
Cl
I
Cl B Cl
I
Si I
I
Structural formulas for
polyatomic ions
•
•
•
•
Example: NO3The negative charge means there is one
more electron
How many total electrons are there?
5 + 3(6) + 1 = 24 electrons.
•
•
NO3Draw:
Resonance
•
Resonance
– Example: ozone
– Ozone’s structure can be drawn like this:
O
O
O
O
O
O
– The actual bonding in ozone is not like
either of these structures. The actual
structure lies somewhere in between
these two.
– These drawings are resonance structures
of ozone