Download notes and handout

HONORS CHEMISTRY I
CHEMICAL BONDING
Define the following terms:
Chemical bond
Ionic bond
Covalent bond
Nonpolar covalent bond
Polar covalent bond
Bonding is not as black and white as we have described it. Bonds can have both ionic and covalent character.
How ionic or covalent a bond is can be determined by comparing the electronegativities of the atoms in the
bonds.

Example: IONIC BOND NaF
o Na – electronegativity = 0.9
o F – electronegativity = 4.0
o Difference = 3.1
o Therefore the bond is ionic

Example: NONPOLAR COVALENT BOND B – H (BH3)
o B – electronegativity = 2.0
o H – electronegativity = 2.1
o Difference = 0.1
o Therefore the bond is nonpolar

Example: POLAR COVALENT BOND H – O (H2O)
o O – electronegativity = 3.5
o H – electronegativity = 2.1
o Difference = 1.4
o Therefore the bond is polar
Bond type
Nonpolar covalent
Polar covalent
Ionic
Electronegativity difference
0.0 – 0.3
0.4 – 1.6
1.7 – 3.2
1
For the following compounds calculate the electronegativity difference between the atoms that are bonded
together and determine the type of bond: (note: do not involve subscripts, difference is just between two atoms)
Compound
NaCl
Electronegativity difference
Type of bond
FeCl3
NO2
PCl3
AgI
Notice that bonds that we called ionic are really polar covalent bonds. Naming rules still apply – metal
nonmetal combinations are still named as ionic compounds. However, when discussing chemical bonding,
different rules apply. When determining chemical formulas we still consider the formation of ions with metals
and nonmetals and these substances do behave ionically in solution.
COVALENT AND MOLECULAR COMPOUNDS:
Define the following:
Molecule:
Diatomic molecule:
Molecular compound:
Chemical formula:
Molecular formula:
Bond energy:
Bond length:
What makes the Noble gases different from the other elements on the periodic table? How many electrons do
they have in their outer shell (or how many valence electrons)?
*** In chemical bonding elements try to achieve the NOBLE GAS CONFIGURATION. Elements will gain,
lose or share electrons to achieve this configuration.***
2
EXCEPTIONS TO THE OCTET RULE!!
Hydrogen – stable configuration is obtained when hydrogen has two electrons in its valence when
forming a covalent bond
Boron – stable configuration is obtained when boron has 6 electrons in its valence when forming
covalent bonds
All other atoms (except Be) follow the octet rule.
LEWIS DOT STRUCTURES (ELECTRON DOT NOTATION)
Define the following:
Unshared pair:
Lewis structures:
Structural formula:
A Sure-Fire Way to Draw Lewis Structures!
Lewis structures are a way to write chemical compounds where all the atoms and electrons are shown.
Sometimes, people have a lot of trouble learning how to do this. However, using the methods on this page, you
should have very little trouble.
How to draw Lewis structures for molecules that contain no charged atoms
1) Count the total valence electrons for the molecule: To do this, find the number of valence electrons for
each atom in the molecule, and add them up.
2) Count 8 electrons for every atom in the molecule except hydrogen (two for H) or boron (6 for boron):
This satisfies the octet rule for the atoms so each atom can achieve the noble gas configuration.
3) Subtract the valence electrons from octet electrons: Or, in other words, subtract the number you found in
#1 above from the number you found in #2 above. The answer you get will be equal to the number of bonding
electrons in the molecule.
4) Divide the number of bonding electrons by two: This gives you the total number of bonds in the
molecule. (Remember, because every bond has two electrons, the number of bonds in the molecule will be
equal to the number of bonding electrons divided by two.)
3
5) Draw an arrangement of the atoms for the molecule that contains the number of bonds you found in
#4 above: Some handy rules to remember are these:
Hydrogen and the halogens bond once.
The family oxygen is in bonds twice.
The family nitrogen is in bonds three times. So does boron.
The family carbon is in bonds four times.
A good thing to do is to bond all the atoms together by single bonds, and then add the multiple bonds until the
rules above are followed.
6) Find the number of lone pair (nonbonding) electrons by subtracting the bonding electrons (#3 above)
from the valence electrons (#1 above). Arrange these around the atoms until all of them satisfy the octet
rule: Remember, ALL elements EXCEPT hydrogen want eight electrons around them, total. Hydrogen only
wants two electrons.
Let's do an example: CO2 Note: Each of the numbers below correspond to the same numbered step above.
1) The number of valence electrons is 16. (Carbon has four electrons, and each of the oxygens have six, for a
total of 4 + 12 = 16 electrons).
2) The number of octet electrons is equal to 24. (Carbon wants eight electrons, and each of the oxygens want
eight electrons, for a total of 8+16 = 24 electrons).
3) The number of bonding electrons is equal to the octet electrons minus the valence electrons, or 8.
4) The number of bonds is equal to the number of bonding electrons divided by two, because there are two
electrons per bond. As a result, in CO2, the number of bonds is equal to 4. (Because 8/2 is 4).
5) If we arrange the molecule so that the atoms are held together by four bonds, we find that the only way to do
it so that we get the following pattern: O=C=O, where carbon is double-bonded to both oxygen atoms.
6) The number of nonbonding electrons is equal to the number of valence electrons (from #1) minus the number
of bonding electrons (from #3), which in our case equals 16 - 8, or 8. Looking at our structure, we see that
carbon already has eight electrons around it. Each oxygen, though, only has four electrons around it. To
complete the picture, each oxygen needs to have two sets of nonbonding electrons, as in this Lewis structure:
:O=C=O:
How to draw Lewis structures for molecules that contain one or more charged atoms
This method is basically the same one you learned above, except the way the number of valence electrons is
determined is different. That change is shown below.
1) Count the total valence electrons for the molecule: To do this, find the number of valence electrons for
each atom in the molecule, and add them up. For polyatomic anions, add the charge of the ion to the
number of valence electrons. For polyatomic cations, subtract the charge of the ion from the number of
valence electrons.
4
Draw the correct Lewis structure for these compounds or ions.
1) PCl3
2) CO
3) C2H2
4) H2CO
5) CCl4
6) BF3
7) N2F2
8) HCN
9) SO42-
10) NO31-
11) CO32-
12) PO43-
13) NH3
14) SO2
5
Formal charges:
To find the charge on each atom, compare the number of electrons that each atom has to the number of
valence electrons it usually has. For this purpose, each bond counts as one electron and each lone pair counts
as two electrons. For example, in CO2 above, carbon has four electrons (because it has four bonds) and oxygen
has six (two bonds + 4 lone pair electrons). If the number of electrons that the atom has is more than the
normal number of valence electrons, the atom has a negative charge. If the number is less than the normal
number of valence electrons, the atom has a positive charge. If it's the same, the atom has no charge at all.
Find the formal charges for each atom in the following compounds: (draw Lewis structures first)
1) SO3
2) CO
3) PO33-
4) O3
RESONANCE: Read the section on resonance found on pages 381-383 in your textbook.
Draw the Lewis structure and all of the resonance structures for these compounds:
1) SO3
2) SO2
3) NO31-
6
VSEPR (section 10.2)
(Valence Shell Electron Pair Repulsion)
VSEPR stands for Valence Shell Electron Pair Repulsion. It's a complicated acronym, but it means something
that's not difficult to understand. Basically, the idea is that covalent bonds and lone pair electrons like to stay as
far apart from each other as possible under all conditions. This is because covalent bonds consist of electrons,
and electrons don't like to hang around next to each other much because they both are negatively charged.
This VSEPR thing explains why molecules have their shapes. If carbon has four atoms stuck to it (as in CH4),
these four atoms want to get as far away from each other as they can. This isn't because the atoms necessarily
hate each other, it's because the electrons in the bonds hate each other. That's the idea behind VSEPR.
In order to find the shape of the molecule
1) Draw the Lewis structure for the molecule. This vital if you're going to get the answer right.
2) Count the number of "things" on the atom you're interested in (or central atom). Let's say that you're
looking at methane, CH4. Now, the vague term "things" refers to atoms and lone pairs. IT DOES NOT REFER
TO THE NUMBER OF BONDS! When you look at methane, there are four atoms stuck to it, if you were to
look at CO2 there would only be two “things” attached to it.
People get confused with multiple bonds. Take carbon dioxide, for example. There are four bonds (carbon is
double-bonded to each oxygen) but only two oxygen atoms bonded to carbon. In this case, we count two things
stuck to carbon, because we only count the atoms, NOT the number of bonds.
Likewise, with NH3 there are four things. Three of the things on nitrogen are hydrogen atoms and the fourth is a
lone pair. For the purposes of VSEPR, lone pairs count exactly the same as atoms, because they consist of
negative charge, too.
3) Count the number of lone pairs that are on the atom you're interested in. IMPORTANT: This does
NOT mean to count the number of lone pairs on all of the atoms in the molecule. Lone pairs on other atoms
aren't important - what's important is only what's directly stuck to the atom you're interested in.
Then look on the chart to determine the shape of your molecule.
MOLECULAR POLARITY:
Not only can bonds be polar and nonpolar but molecules can be polar or nonpolar as well. If a compound is
symmetrical, then even with polar bonds the molecule will be nonpolar.
Example: CBr4 the bond between C-Br is polar but the molecule is nonpolar because it is symmetrical.
Water: H2O the O-H bonds are polar and the molecule is polar because it has a bent shape and is not
symmetrical.
7
Form
AX2
Electron
Domains
2
Geometry
Linear
AX3
3
Trigonal
Planar
AX2E
3
Bent
Shape of
Molecule
Examples
CO2
BF3
SO2
AX4
4
Tetrahedral
CH4
AX3E
AX2E2
4
4
Trigonal
Pyramidal
NH3
Bent
H2O
8
Formula
Electron Dot Structure
Shape of Molecule
Molecular polarity (polar or
nonpolar)
H2
HBr
H2O
PH3
CH4
HClO
N2
SO3
H2CO
C2H2
CH3Br
HCOOH
HCN
H2O2
9
IONIC BONDS (section 6.3)
Define the following terms:
Ionic compound:
Formula unit:
Lattice energy:
Properties of ionic vs. covalent compounds
List 3 differences between ionic and covalent compounds in terms of their properties.
METALLIC BONDING
Define:
Metallic bond
Malleablility
ductility
10