Covalent Bonds Homework
[1] Drawing Lewis Structures
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Step 1: Find the total number of valence electrons.
Step 2: Choose a central atom.
o Always carbon, never hydrogen or a halogen
o Often the atom listed first or atom that is most electronegative.
o Multiple carbon atoms connect as chains
Step 3: Draw in single bonds
o Carbon always has 4 bonds except for CO (carbon monoxide).
o Oxygen usually has 2 bonds and hydrogen 1.
o Make symmetrical if possible.
Step 4: add in electron pairs and double/triple bonds as necessary to create octets and
use up all the electrons.
Step 5: Check formal charge to see if a better Lewis structure can be built
Step 6: put in brackets and charge is it’s a polyatomic ion, draw resonance structures if
necessary
Draw the following Lewis Structures
H2
NH3
CF4
CO2
CH4
O2
Cl2
N2
H2O
C2H4
Covalent Bonds Homework
HCl
C4H10
CO
C2H2
CF2I2
CH3OH
H2S
CH3CH2OH
O3
HCN
C2H4
N2H2
C3H8
SF6
Covalent Bonds Homework
[2] Identifying Bond Type with Electronegativity (𝝌) Differences
If the Electronegativity difference between two atoms in a non-metallic bond is greater
than 2 then the bond is considered to be ionic. This simply means that one atom is very
strongly pulling on the electrons, so much that we model the electron as being
completely transferred, even though this technically does not happen. The sliding scale
above shows how to determine bond type from electronegativity differences.
If the electronegativity difference between two bonded atoms is between 0 and 0.5 then
the bond is said to be non-polar as electrons will be shared roughly equally between the
two atoms. IF the value lies between 0.5 and 2.0, then the bond is considered to be
polar. The greater the difference, the stronger the polarity.
Water is a good example. Oxygen (3.44) is more electronegative than hydgrgen (2.20)
and when you get a difference between them (3.44-2.20) we end up with 1.22.
According to the sliding scale above, this creates a polar bond. The electrons are pulled
closer to oxygen, leading to oxygen having a partial negative charge and hydrogen
being partially positive. Its not a full positive or negative charge like we see in an ionic
bond so we use the Greek letter delta to indicate a partial charge.
Covalent Bonds Homework
Using the information on the previous page and the charts above, indicate why type of
bonds the following atoms would form:
a) Na and Cl
b) N and H
c) F and F
d) Ca and Cl
e) Al and Cl
f) H and Br
g) K and Cl
h) C and O
i) O and O
j) Cl and F
k) Li and O
l) H and O
m) What type of bonds are shown in the images:
Just because a BOND is polar, DOES NOT mean a molecule is polar. Only NH3 is a polar
molecule. Carbon dioxide has polar bonds but due to symmetry it is a non-polar molecule.
Covalent Bonds Homework
[3] Naming Covalent Molecules
A rose by any other name would smell just as sweet
The rules for naming covalent compounds are different than naming ionic compounds.
A system needs to be chosen that can easily distinguish between hundreds and even
thousands of molecules. When the metal sodium loses a valence electron to the
nonmetal chlorine and forms an ionic bond this happens in only one way. Sodium and
chlorine always join in a one to one ratio. But in a covalent molecule where two nonmetals like oxygen and nitrogen bond together, there are many possible combinations.
i.
N2O
ii.
NO
iii.
NO2
iv.
N2O4
v.
N2O5
If we gave each one a special name, then we would quickly have way too
many names to remember. So we use prefixes as a system to name
covalent compounds. The first one has two nitrogen so we call it dinitrogen
and there is only one oxygen so we use the prefix mono. The complete
name is dihydrogen monoxide, indicating that we add ide to the end similar
to what we do with ionic compounds. In the second case there is one atom
of nitrogen and one atom of oxygen but we don’t use the mono prefix for
the first element. So this compound would be called Nitrogen monoxide.
Fill in the names of the Nitric Oxides above and name the molecules below:
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CO =
S02 =
NO3 =
SF6 =
CO2 =
CF4
P4S5
Write the Chemical Formula for the following molecules:
• Dinitrogen trioxide =
• Carbon Tetrafluoride =
• Dinitrogen pentoxide =
• Silicon TetraChloride =
• Disilicon Hexachloride
• Hexaboron silicide
• Antimony tribromide
Which of the following is named incorrectly?
a) SbCl3, antimony trichloride
b) C2O5, dicarbon pentoxide
c) CF4, carbon tetrafluoride
d) H3As, hydrogen arsenide
Covalent Bonds Homework
[4] Calculating Formal Charge
Formal charge is a simple book-keeping method for determining which Lewis structure
is the best when multiple Lewis Structures are possible. Recall that atoms beyond row
three of the periodic table can violate the octet rule.
Task: Draw 2 Lewis Structures for sulfate, one where the sulfur has all single bonds
and one where the sulfur atom has 2 single bonds and 2 double bonds:
[SO4]2-
[SO4]2-
all single bonds
2 double & 2 single bonds
How do we know which structure is preferable? We use formal charge and the rule is
we want a formal charge as close to zero as possible.
F.C. = #VEs – (Bonds + Non-bonded Electrons)
In the Lewis structure to the left, Carbon has a formal
charge of zero. F.C. = 4 – (4 + 0) as Carbon has 4 VEs,
4 Bonds and 0 non-bonded electrons. Each Oxygen has the
same formal charge of zero as both of them have 2 bonds,
4 non-bonded electrons and 6VEs.
F.C. = 6 – (2 + 4) = 0
The formal charge on all the individual atoms in a molecule must add up to the total
charge of the molecule itself. Thus, sulfate has a 2- charge and this means when you
consider all the atoms within it, their formal charge should total up to negative two. Use
the formula above (show your work) and calculate the formal charge on every atom in
both molecules. Which one has a preferable formal charge?
Sulfate with all Single Bonds
Sulfur:
Hydrogen-L
Hydrogen-R
Oxygen-L
Oxygen-R
Oxygen-T
Oxygen-B
Sulfate with 2 Single & 2 Double Bonds
Sulfur:
Hydrogen-L
Hydrogen-R
Oxygen-L
Oxygen-R
Oxygen-T
Oxygen-B
Covalent Bonds Homework
[5] Drawing Lewis Structures of Polyatomic Ions
[NO3]- The Nitrate ion has 24 Valence electrons
N = 5 VE O = 6VE x 3 and the Negative sign means one additional electron.
1) Down below, draw an N in the middle and connect three oxygens to it with single
bonds and then fill in 6 additional dots around each oxygen.
2) We have used all 24 VEs but there is a problem here. The oxygens all have octets
but the nitrogen atom does not. To fix this, pick any oxygen atom and cross off one of its
pairs of electrons and make a double bond between it and nitrogen. Now all the atoms
have octets. Draw the image below.
3) An additional problem surfaces. We could have chosen any of the atoms for a double
bond so we have to show this in our Lewis structure. Experimental measurements
actually show the bond lengths are all the same in the nitrate ion so one bond isn’t really
a double bond while two are single bonds. The actual bond type lies somewhere
between these and all three bonds are equal. So we must draw two more structures
with double sided arrows in between them, changing the location of the double bond in
each case. The double sided arrows indicate these are resonance structures as there is
no difference between the bonds in the nitrate ion. The final model and correct Lewis
Structure looks like what we see in the image below.
4) Notice the brackets. Anytime we draw a Lewis structure for a polyatomic ion we must
include the brackets and the charge.
Covalent Bonds Homework
Draw the Lewis Structure for the following Polyatomic Ions
Include Resonance Structures if Necessary, Brackets, Charge and formal charge on each atom
Ammonium
Phosphate
Carbonate
Hydroxide
Chlorate
Nitrite
Chlorite
Sulfite
Cyanide
Sulfate
Covalent Bonds Homework
[6] Polar Bonds vs Polar Molecules : the shape of molecules is extremely important in the
world around. The properties of molecules and how they react with each other and substances
around them is predicated on their bond type and shape. A polar molecule is a molecule that has
two different ends with opposite partial charges. Recall that water has a partial positive end and a
partial negative end. This is due to the fact that oxygen is more electronegative than hydrogen
and pulls the electrons closer when sharing them—resulting in a negative charge around the
oxygen atom.
Carbon Dioxide has Polar Bonds as Oxygen is far more polar than Carbon but Carbon dioxide is
a liner molecule so opposite ends do not have opposite partial charges.
CO2 has polar bonds bot
opposite ends are both
negative. It is NOT polar.
Water on the other hand, is a bent molecule so opposite ends have opposite partial charges.
If water was linear like CO2, it would not be polar, and there wouldn’t be hydrogen bonds
leading to many of water unique properties such as its high specific heat capacity, high surface
tension, high boiling point and the fact that ice is denser than its liquid equivalent. If not for this
polarity, water would not be a liquid at room temperatures and we would all be dead.
DNA sequencing is another application of molecular gemoegry and (Coulomb’s law). The shape
of DNA molecules it what allows them to fit together in a specific sequence. Figuring out why
water, which has two oxygens bonded to a carbon, has a different molecular geometry than
carbon dioxide, which also has two atoms bonded to a central one, is our current task. We are
going to go beyond this and figure out how to determine the modeled 3D geometry of any simple
Lewis structure.
SUMMARY: in order for a MOLECULE to be polar you need polar bonds
AND a non-symmetrical shape.
Question: the molecule on the right is CF4. Fluorine is the most
electronegative atom on the periodic table so these bonds are definitely polar.
Is the molecule itself polar or non-polar?
Covalent Bonds Homework
[7] Molecular Geometry: VSEPR Theory The image bellows shows some of the shapes
molecules fall into. For example, if you have a central atom surrounded by 3 atoms and no lone
pairs, it would have a trigonal planar geometry. The rational here is that electron pairs want to
be as far apart from one another as possible (like charges repel).
Questions In the trigonal planar shape, what is the angle between the outer atoms? To answer
this, think about how many degrees are in a circle and that it is being slices into three equal
angles.
Question: How many degrees are there between atoms in a linear molecule?
Notice the little green balls on the bent and trigonal pyrimadal shapes. These represent lone
electron pairs, not bonds. Water has two hydrogen atoms
connected to an oxygen atom and two lone pairs. There are four
things around the oxygen atom in a water molecule so it looks
tetrahedral with two atoms missing. Water is bent because of
the lone electron pairs. When we look at a molecule like Carbon
Dioxide we don’t have any
lone electron pairs so the
two oxygens surrounding
the carbon lead to a linear
shape. When determine what shape a molecule is you must first
1) Draw its Lewis Structures
2) Ask how many things are around the central atoms (lone pairs and any bond counts as 1
thing)
3) Match it to the chart based on its structure.
Fill in the Chart on the next two pages using VSEPR theory
Covalent Bonds Homework
MOLECULE
Lewis Structure
Just draw double arrows to show resonance
Electron
Geometry
Final
Molecular
Geometry
Polar or
Non Polar
Bond
Polar or
nonpolar
Molecule?
NA: Ion
NA: Ion
Carbon
Monoxide
Dihydrogen
Sulfide
Nitrate Ion
Carbon
Tetrachloride
Ammonia
Covalent Bonds Homework
Ozone
NA: Ion
NA: Ion
NA: Ion
NA ION
NA: Ion
NA ION
Phosphate
Nitrate
Sulfate
Sulfur
Hexafluoride
Question: Why do Nitrate [NO3]- and Ammonia (NH3) have different molecular geometries
when each case is clearly a Nitrogen atom surrounded by three other atoms?
Covalent Bonds Homework
[8] Hybridization and Sigma (𝝈) and Pi (𝝅 )bonds
Carbon has 6 electrons and an electron configuration of 1s22s22p2. Carbon
is in Column 4A and has 4 electrons available for bonding but according to
the Aufbau Principle and Hund’s Rule, the electron diagram below only
shows two unpaired electrons available for bonding.
Carbon is extremely important and forms the basis of life and all organic molecules.
How can our model of where electrons live in the atom be inadequate at describing
such a fundamental atom and the bonds it forms? In order to remedy this situation,
Chemists came up with the idea of hybridized orbitals. Up above they combine the 2s
and 2p orbitals into one sp3 hybridized orbital. The s indicates one s box (suborbitals)
that holds two electrons and the p3 indicate 3p boxes (suborbitals) that each hold two
electrons.
We combine the 2s
and 2p orbitals into
one hybridized orbital
Now there are four electron available for bonding in the carbon atoms
and it can form something like methane (CH4) or Carbon Tetrachloride.
Because the carbon atom on the right has four bonding regions around
it, we say that it is sp3 hybridized.
Electron pairs count as one electron cloud and any bonding regions
count as an. Electron cloud A double bond only counts as one.
Hybridization of Central atom:
2 electron domains = sp hybridization
3 electron domains = sp2 hydridication
4 electron domains = sp3 hybridization
Example: the Cl on the right has one bonding region and three electron pairs which is four
electron clouds or domains. Ergo, the Cl is sp3 hybridized.
Covalent Bonds Homework
Sigma (𝝈) and Pi (𝝅 )bonds
Lone nitrogen gas atoms are unstable and will pair off with one
another forming (diatomic) nitrogen gas in order to achieve an
octet. As we saw previously, Nitrogen forms a triple bond with
itself but the Lewis tructure we draw leaves a lot to be desired. It
gives the impression that there are 6 electrons shared right between the two nitrogen atoms.
Electron are negativelycharged and want nothing to do with one another. The image shown
bleow actually shows a molecule with a triple bond.
Two electrons overlap between the molecules forming
a sigma bond (𝝈) and the other four electrons overlap
in p orbitals,far apart from one another.
The image to the
right shows the
hybridization of
diatomic nitrogen.
Nitrogen has an electron configuration of 1s22s22p3.Nitrogen
is in column 5A and has 5 VEs available for bonding. Two
are paired off and 3 form bonds in an N2 molecule.
Technically, we don’t have to hydridize nitrogen like the
Carbon atom as the old system still works, but we do so
anyways. Each nitrogen has two electron clouds around it
(one pair of electrons and one bonding region) which means
sp hybridization and two lone p orbitals left over. It is the
two lone p orbitals which form the double and triple bond.
Hence the name pi orbirats.
The first bond between atoms is always a sigma bond occuring
directly between their internuclear axis. The second (and third if
necessary) is always a pi bond formed away from the axis of the
nuclues in unhybridized p orbitals.
Each Oxygen atom in Oxygen (O2) gas one bonding region (a
double bond) and two electron pairs around it. This leads to sp2
hybridization and leaves one lone p orbital. It is this lone p orbital
which forms the second bond (pi bond!)
 The first bond is a sigma bond.
 The second bond is a pi bond (sp2)
 The third bond is also a pi bond (sp)
Hybridized orbitals allow us to model carbon based chemistry
and explain how double and triple bonds can exist when
electrons want nothing to do with one another.
Covalent Bonds Homework
Bond Strength Sigma bonds are stronger than pi bonds but a double bond will be stronger than a
single bond as there are two points of bonding in the molecule. A triple bond wil be stronger than
a double bond. Think of hanging a heavy mass from a single rubber band. It may easily break but
if you hung the mass from two rubber bands it would be less likely to break and even less likely
still if you used three rubber bands (bonds!).



H-H bond requires 432 KJ/mol to break
The bonds between oxygen in O2 require 494KJ/mol
Nitrogen bonds in N2 require 942kJ/mol to break
Explain the above three bullet points in terms of what you know about diatomic hydrogen,
nitrogen and oxygen gas.
Usse the pictures below to fill in the table:
Molecule
BeCl2
BH3
CH4
NH3
H2O
HF
CH2O
CO2
C2H2
C2H4
Geometry
Hybridization
Sigma Bonds (𝜎)
Pi bonds (𝜋)
Covalent Bonds Homework
Δ Electronegativity
Bond Type
0.0-0.5
Non-Polar Covalent
≥0.5 to >1
Polar Covalent
≥1 to <2
Very Polar Covalent
≥2
Ionic Bond (Extremely
Polar)
Number
Prefix
1
mono
2
di-
3
tri
4
tetra
5
penta
6
hexa
7
hepta
8
octa
9
nona
10
deca