Lesson 8 - Covalent Bonding

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Unit:
Lesson 8:
Understanding Covalent
Bonding
Lesson Objectives:
1. strong electrostatic attraction between
two nuclei and the shared pair(s) of
electrons between them.
2. dot and cross diagrams to show electrons
in simple covalent molecules, including
those with multiple bonds and dative
covalent (coordinate) bonds.
3. the relationship between bond lengths
and bond strengths in covalent bonds.
4. tetrahedral basis of organic chemistry
Bonding Types
Theory:
Tutor Presentation
• The physical properties of a substance depend on its structure and type
of bonding present. Bonding determines the type of structure.
• CHEMICAL strong bonds
• ionic (or electrovalent)
• covalent
• dative covalent (or co-ordinate)
• metallic
• PHYSICAL weak bonds
• van der Waals‘ forces (weakest)
• dipole-dipole interaction
• hydrogen bonds (strongest)
Covalent bonding
If ionic bonds form between metals and
non-metals, between what class of
element do covalent bonds form?
Non-metals and non-metals
Theory:
Tutor Presentation
Nirvana
Why?
Electron loss
Electron gain
Covalent Bonding
Theory:
Tutor Presentation
• Covalent bonds are formed by atoms sharing electrons to form molecules.
• One single covalent bond is a sharing of 1 pair of electrons, two pairs of
shared electrons between the same two atoms gives a double bond and it
is possible for two atoms to share 3 pairs of electrons and give a triple
bond.
• This kind of bond or electronic linkage does act in a particular direction
i.e. along the 'line' between the two nuclei of the atoms bonded together,
this is why molecules have a particular shape. In the case of ionic or
metallic bonding, the electrical attractive forces act in all directions
around the particles involved.
Why do covalent bonds form?
Theory:
Tutor Presentation
Covalent bonds often form between atoms with too many
electrons in their valence shells to give away, but not
enough to easily fill.
Thus they share electrons with their
neighbours, in such a way that including the
shared electrons the shells are full
Delocalizing electrons over two atoms instead
of one lowers the energy of the system
Theory:
Tutor Presentation
Covalent Bonding
electrostatic attraction between nuclei and the shared electrons is localised and
in a specific direction
giant covalent structure is a lattice of many atoms bonded covalently
Covalent Bonding
Covalent bonding describes a chemical bond where a pair of electrons
are shared between 2 atoms.
Theory:
Tutor Presentation
But how does this hold 2 atoms together?
Attractive forces
Repulsive forces
The atomic separation of particles
in a nucleus is determined by the
balance of these forces
Atoms with a covalent bond are held together by an electrostatic
attraction between the nucleus of each atom and the shared electrons.
Theory:
Tutor Presentation
Covalent Bonding
A chlorine atom has an
electronic configuration
of [Ne]3s23p5.
By each sharing a 3p
electron, they can
achieve this.
Two chlorine atoms
would both want to
gain another electron.
Covalent Bonding – Dative
Theory:
Tutor Presentation
A co-ordinate bond (dative covalent bond) is a covalent bond where
both of the electrons in the bond come from just one of the atoms
involved.
• It is represented by an arrow coming from the atom donated the electrons.
• The atom that accepts does not have a full outer shell of electrons – it is electron
deficient.
• The donating atom, donates a non-bonding pair of electrons – this is known as a
lone pair.
• Co-ordinate bonds do not differ in strength from regular covalent bonds.
Theory:
Tutor Presentation
Covalent Bonding – Dative
Lone Pairs
Theory:
Tutor Presentation
Lone pairs occur in elements from group 5, 6 and 7
Lone
pair
Lone
pairs
Lone pairs affect the shape of the molecule
Covalent Bonding
Theory:
Tutor Presentation
• Example 1: Two hydrogen atoms form the molecule of the element
hydrogen (H2)
H
H
H
H
• Example 2: One atom of hydrogen combines with one atom of
Chlorine to give you hydrogen chloride (HCl)
Cl
H
Cl
H
Covalent Bonding
Theory:
Tutor Presentation

Draw the bonding in ammonia (NH3)
H
H
H
N
H
N
H
H
Covalent Bonding
Theory:
Tutor Presentation

Draw the bonding in methane (CH4)
H
H
H
H
C
H
H
C
H
H
Covalent Bonding
Theory:
Tutor Presentation

What about oxygen?
O
O
each atom needs two electrons
to complete its outer shell
O
O
each oxygen shares 2 of its
electrons to form a
DOUBLE COVALENT BOND
Supervised Learning
Covalent Bond Diagrams
Draw dot cross diagrams for
the following:
Hydrogen
Oxygen
Carbon dioxide
Carbon tetrafluoride
Physical Properties (Covalent Bonding)
For simple covalent molecules (not giant covalent)
Theory:
Tutor Presentation
• Electrical
Don’t conduct electricity as they have no mobile ions or electrons
• Solubility
Tend to be more soluble in organic solvents than in water
• Boiling point
The forces between molecules (intermolecular forces) are weak and known as van der Waals
forces. Attractions between molecules increases as the molecules get more electrons.
e.g.
CH4 -161°C
C2H6 - 88°C
C3H8 -42°C
as the intermolecular forces are weak, little energy is required to separate molecules from each
other so boiling points are low
Bond Length in Covalent Molecules
Theory:
Tutor Presentation
• Bond length is defined as the distance between the centres of two
covalently bonded atoms.
• The length of the bond is determined by the number of bonded electrons
(the bond order).
• The higher the bond order, the stronger the pull between the two atoms
and the shorter the bond length.
• Generally, the length of the bond between two atoms is approximately
the sum of the covalent radii of the two atoms.
• Bond length is reported in picometers. Therefore, bond length increases in
the following order: triple bond < double bond < single bond.
Theory:
Tutor Presentation
Bond Length in Covalent Molecules
To find the bond length, follow these steps:
1.Draw the Lewis structure.
2.Look up the chart below for the radii for the corresponding bond.
3.Find the sum of the two radii.
EXAMPLE : CCL4
Determine the carbon-to-chlorine bond length in CCl4.
SOLUTION
Using Table A3, a C single bond has a length of 75 picometers and that a Cl single bond has a length of 99
picometers. When added together, the bond length of a C-Cl bond is approximately 174 picometers.
EXAMPLE : CO2
Determine the carbon-oxygen bond length in CO2.
SOLUTION
Using Table A3, we see that a C double bond has a length of 67 picometers and that an O double bond has a
length of 57 picometers. When added together, the bond length of a C=O bond is approximately 124
picometers.
Bond Strength in Covalent Molecules
Theory:
Tutor Presentation
Triple bonds between like atoms are shorter than double bonds, and because more energy is required to
completely break all three bonds than to completely break two, a triple bond is also stronger than a double
bond. Similarly, double bonds between like atoms are stronger and shorter than single bonds. Bonds of the
same order between different atoms show a wide range of bond energies.
Single Bonds
Multiple Bonds
H–H
432
C–C
346
N–N
≈167
O–O
≈142
F–F
155
C=C
602
H–C
411
C–Si
318
N–O
201
O–F
190
F–Cl
249
C≡C
835
H–Si
318
C–N
305
N–F
283
O–Cl
218
F–Br
249
C=N
615
H–N
386
C–O
358
N–Cl
313
O–Br
201
F–I
278
C≡N
887
H–P
≈322
C–S
272
N–Br
243
O–I
201
Cl–Cl
240
C=O
749
H–O
459
C–F
485
P–P
201
S–S
226
Cl–Br
216
C≡O
1072
H–S
363
C–Cl
327
S–F
284
Cl–I
208
N=N
418
H–F
565
C–Br
285
S–Cl
255
Br–Br
190
N≡N
942
H–Cl
428
C–I
213
S–Br
218
Br–I
175
N=O
607
H–Br
362
Si–Si
222
I–I
149
O=O
494
H–I
295
Si–O
452
S=O
532
Theory:
Tutor Presentation
Bond Strength in Covalent Molecules
1.
Bonds between hydrogen and atoms in the same column of the periodic table decrease in
strength as we go down the column. Thus an H–F bond is stronger than an H–I bond, H–C is
stronger than H–Si, H–N is stronger than H–P, H–O is stronger than H–S, and so forth. The
reason for this is that the region of space in which electrons are shared between two atoms
becomes proportionally smaller as one of the atoms becomes larger.
2.
Bonds between like atoms usually become weaker as we go down a column (important
exceptions are noted later). For example, the C–C single bond is stronger than the Si–Si single
bond, which is stronger than the Ge–Ge bond, and so forth. As two bonded atoms become
larger, the region between them occupied by bonding electrons becomes proportionally
smaller.
3.
Noteworthy exceptions are single bonds between the period 2 atoms of groups 15, 16, and 17
(i.e., N, O, F), which are unusually weak compared with single bonds between their larger
congeners. It is likely that the N–N, O–O, and F–F single bonds are weaker than might be
expected due to strong repulsive interactions between lone pairs of electrons on adjacent
atoms. The trend in bond energies for the halogens is therefore
4.
Cl–Cl>Br–Br>F–F>I–I
Theory:
Tutor Presentation
Bond Strength in Covalent Molecules
1. Elements in periods 3 and 4 rarely form multiple bonds with themselves, their multiple bond energies are
not accurately known. Nonetheless, they are presumed to be significantly weaker than multiple bonds
between lighter atoms of the same families. Compounds containing an Si=Si double bond, for example,
have only recently been prepared, whereas compounds containing C=C double bonds are one of the beststudied and most important classes of organic compounds.
2. Multiple bonds between carbon, oxygen, or nitrogen and a period 3 element such as phosphorus or sulfur
tend to be unusually strong. In fact, multiple bonds of this type dominate the chemistry of the period 3
elements of groups 15 and 16. Multiple bonds to phosphorus or sulfur occur as a result of d-orbital
interactions.
3. In contrast, silicon in group 14 has little tendency to form discrete silicon–oxygen double bonds.
Consequently, SiO2 has a three-dimensional network structure in which each silicon atom forms four Si–O
single bonds, which makes the physical and chemical properties of SiO2 very different from those of CO2.
Bond Strength in Covalent Molecules
Theory:
Tutor Presentation
Bond strengths increase as bond order increases, while bond
distances decrease.
The Strength of Covalent Bonds Depends on the Overlap between the Valence Orbitals of the Bonded
Atoms. The relative sizes of the region of space in which electrons are shared between (a) a hydrogen
atom and lighter (smaller) vs. heavier (larger) atoms in the same periodic group; and (b) two lighter
versus two heavier atoms in the same group. Although the absolute amount of shared space increases
in both cases on going from a light to a heavy atom, the amount of space relative to the size of the
bonded atom decreases; that is, the percentage of total orbital volume decreases with increasing size.
Hence the strength of the bond decreases.
Theory:
Tutor Presentation
Shapes of molecules
• The shapes of molecules and ions can be accurately predicted using
the Sidgwick-Powell Rules.
• These rely on the fact that the vast majority of molecules contain an
even number of electrons arranged in pairs - bond pairs and lone
pairs
Shapes of molecules
Theory:
Tutor Presentation
We often draw molecules as 2-dimensional structures, take methane
for example.
However, realistically very few shapes are ‘flat’. Methane actually
looks like this:
We can show that molecules
are 3D like so:
Shapes of molecules
Theory:
Tutor Presentation
So how do we predict the shape that a molecule will be?
We use electron pair repulsion theory!
This is based on the concept that:
1. Each pair of electrons around an atom repels all of the
other pairs.
2. Pairs of electrons take up positions as far away from each
other as possible in order to minimize this repulsion.
Shapes of molecules
Theory:
Tutor Presentation
In order to work out the shape of a molecule, you need to
work out the number of electron pairs around the central
atom.
In order to do this you must draw a dot and cross structure for
the molecule.
Don’t take short-cuts, it often leads to silly mistakes!
You also need to learn the bond angles and shapes that can
occur!
Theory:
Tutor Presentation
Sidgwick Powell Rules
1. The electron pairs in a molecule arrange themselves to be as far
apart as possible so as to minimise repulsion.
Theory:
Tutor Presentation
Sidgwick Powell Rules
1. The electron pairs in a molecule arrange themselves to be as far
apart as possible so as to minimise repulsion.
2. Lone pairs are more repelling than bond pairs
Theory:
Tutor Presentation
Sidgwick Powell Rules
1. The electron pairs in a molecule arrange themselves to be as far
apart as possible so as to minimise repulsion.
2. Lone pairs are more repelling than bond pairs
3. A double or triple bond behaves just as a single pair for these
purposes
Shapes of molecules
Theory:
Tutor Presentation
2 pairs of electrons eg BeCl2
• To be as far apart as possible, the electron pairs are arranged at
180°- the molecule is linear
Shapes of molecules
Theory:
Tutor Presentation
3 pairs of electrons eg BF3
Three pairs as far apart as possible is achieved by arranging the pairs
at 120° - trigonal planar
Theory:
Tutor Presentation
Shapes of molecules
Four pairs of electrons
In each of the 3 molecules below there are four pairs. Methane is
tetrahedral; ammonia is pyramidal and water is V -shaped. Note the
lone pairs are more repelling
Theory:
Tutor Presentation
Shapes of molecules
Five bonding pairs of electrons
PCl5 has five pairs of bonding electrons - these are arranged in a trigonal
bipyramid
Theory:
Tutor Presentation
Shapes of molecules
Six bonding pairs of electrons
SF6 has six pairs of bonding electrons; these are arranged symmetrically
as shown - octahedral
Shapes of molecules
Worked example: What is the H-C-H bond angle in methane?
Theory:
Tutor Presentation
1. Draw a dot and cross
diagram for the molecule.
It is okay to show outer electrons only.
2. Count the number of areas
of electron density around the
central atom.
In this case there are 4.
3. Use this to chose the shape
that the molecules is based
on.
4. Use the shape to assign an
appropriate bond angle.
4 areas of electron density is
tetrahedral.
109.5o
This method works when there
are no non-bonding pairs of
electrons.
Shapes of molecules
Theory:
Tutor Presentation
Worked example, dealing with lone (non-bonding) pairs: What is the HN-H bond angle in ammonia?
1. Draw a dot and cross
diagram for the molecule.
In this case there are 4.
2. Count the number of areas
of electron density around the
central atom.
4 areas of electron density is
tetrahedral.
This shape is trigonal pyramidal.
3. Use this to chose the base
geometry of the molecule.
Note: Lone pairs repel more than
bonding pairs, so the angle changes.
4. Identify that one of the areas of
electron density is a lone pair and
assign the actual shape.
5. Use the new shape to predict
the bond angle.
107o
Shapes of molecules: Tips and Tricks
Theory:
Tutor Presentation
We use the term electron density rather than pairs of electrons
to avoid confusion. In CO2 there are four pairs of electrons, but
only 2 areas of electron density due to the double bond.
Sometimes central atoms can bond
happily without having a full pair of
electrons – don’t get caught out!
Repulsion between pairs of electrons:
Sometimes you have to predict
shapes of ions!
bonding pair – bonding pair
lone pair – bonding pair
lone pair – lone pair
Increasing
repulsion
Remember to include this in your dot
and cross!
Try and predict the shape of the NH4+
ion and the ClF4- ion!.
Theory:
Tutor Presentation
Theory:
Tutor Presentation
Required Reading
Websites
• http://chemguide.co.uk/atoms/bonding/covalent.html#top
• http://chemguide.co.uk/atoms/bonding/doublebonds.html#top
• http://chemguide.co.uk/atoms/bonding/dative.html#top
• http://chemguide.co.uk/atoms/bonding/shapes.html#top
• http://chemguide.co.uk/atoms/bonding/shapesdouble.html#top
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