UNIT 4:COVALENT BOND AND MOLECULAR STRUCTURE
UNIT 4: COVALENT BOND AND MOLECULAR
STRUCTURE
Key Unit Competence
Demonstrate how the nature of the bonding is related to the properties of covalent
compounds and molecular structures.
Learning objectives
By the end of this unit, students should be able to:
• Define octet rule as applied to covalent compounds.
• Explain the formation of covalent bonds and
the properties of
covalent compounds.
• Describe how the properties of covalent compounds depend on their
bonding.
• Explain the rules of writing proper Lewis structures
• Draw different Lewis structures
• State the difference between Lewis structures from other structures.
• Apply octet rule to draw Lewis structures of different compounds.
• Make the structures of molecules using models.
• Write the structures of some compounds that do not obey octet rule.
• Explain the formation of dative covalent bonds in different molecules.
• Compare the formation of dative covalent to normal covalent bonding.
• Describe the concept of valence bond theory.
• Relate the shapes of molecules to the type of hybridization.
• Differentiate sigma from pi bonds in terms of orbital overlap and formation.
• Explain the VSEPR theory.
• Apply the VSEPR theory to predict the shapes of different molecules/ions.
• Predict whether the bonding between specified elements will be primarily
covalent or ionic.
• Relate the structure of simple and giant molecular covalent compounds to
their properties.
• Describe simple and giant covalent molecular structures.
• Describe the origin of inter-molecular forces.
• Describe the effect of inter and intra molecular forces on the physical
properties of certain molecules.
• Describe the effect of hydrogen bonding in the biological molecules.
• Relate the physical properties to type of inter and intra molecular forces in
molecules.• Compare inter and intra molecular forces of attraction in different molecules.
In UNIT 3, you have learnt that atoms have different ways of combination to achieve
the stable octet electronic structure; two of those ways of combination led to
the formation of ionic bond and metallic bond. But what happens where the two
combining atoms need electrons to complete the octet structure and no one is
willing to donate electrons? For example the combination of 2 hydrogen atoms or
the combination of 2 chlorine atoms?
When this happens, the combining atoms share a pair of electrons where each
atom brings or contributes one electron. In other words there is an overlapping of
two orbitals, one orbital from one atom, each orbital containing one electron (see
Fig.4.1): this bond is called “Covalent bond”. The attraction between the bonding
pair of electrons and the two nuclei holds the two atoms together.
The covalent bond is a bond formed when atoms share a pair of electrons to complete
the octet. Similarly, people need each other irrespective of their race, economic,
political and social status for the success of human race. Some compounds that
exist in nature such as hemoglobin in our blood, chlorophyll in plants, paracetamol,
responsible for transport of oxygen, green color in plants and as pain killer respectively
are made of the covalent bond. The covalent bonds mostly occur between nonmetals or between two of the same (or similar) elements.Two atoms with similar
electronegativity do not exchange an electron from their outermost shell; the atoms
instead share electrons so that their valence electron shell is filled.
In general, covalent bonding occurs when atoms share electrons (Lewis model),
concentrating electron density between nuclei. The build-up of electron density
between two nuclei occurs when a valence atomic orbital of one atom combines
with that of another atom (Valence bond theory).In Valence bond theory, the bonds
are considered to form from the overlap of two atomic orbitals on different atoms,
each orbital containing a single electron.
The orbitals share a region of space, i.e. they overlap. The overlap of orbitals allows
two electrons of opposite spin to share the common space between the nuclei,forming a covalent bond.
These two electrons are attracted to the positive charge of both the hydrogen
nuclei, with the result that they serve as a sort of ‘chemical glue’ holding the two
nuclei together.
The figure (Figure 4.1) shows the distance between the two nuclei. If the two nuclei
are far apart, their respective 1s-orbitals cannot overlap and no covalent bond is
formed. As they move closer each other, the orbital overlapping begins to occur, and
a bond starts to form.
The examples below represent different atoms overlapping in order to form covalentbonds.
4.1.1 Properties of covalent molecules
Covalent molecules are chemical compounds in which atoms are all bonded together
through covalent bonds. The covalent compounds possess different properties and
some are emphasized below.
• Covalent compounds exist as individual molecules, held together by weak
van der Waals forces.
• Due to the weak van der Waals forces that hold molecules together, covalent
compounds have low melting and boiling points; because the weak forces
between molecules can be broken easily to separate the molecules. That
is why covalent compounds can be solid, liquid and gaseous at room
temperature.
• Covalent compounds do not display the electrical conductivity either in
pure form or when dissolved in water. This can be explained by the fact
that the covalent compounds do not dissociate into ions when dissolves in
water.
• Generally non-polar covalent compounds do not dissolve in water; but
many polar covalent compounds are soluble in water( a polar solvent)
• Non-polar covalent compounds are soluble in organic solvents (themselves
non-polar covalent).
The two statements above are at the origin of the say by chemists: “Like dissolveslike”
bonding is usually placed at the center. The number of bonding sites is determined
by considering the number of valence electrons and the ability of an atom to expand
its octet. As you will progress in your study of chemistry, you will be able to recognisethat certain groups of atoms prefer to bond together in a certain way!
electrons in the molecule or ion. In the case of a neutral molecule, this is nothing
more than the sum of the valence electrons on each atom. If the molecule carries
an electric charge, we add one electron for each negative charge or subtract an
electron for each positive charge.
In Lewis structure, the least electronegative element is usually the central element,
except H that is never the central element, because it forms only one bond.
Another way of finding Lewis structure
1. Calculate n (the number of valence (outer) shell electrons needed by all atoms in
the molecule or ion to achieve noble gas configurations for instance,
NO3
-
, n=1× 8(for N atom) + 3×8 (for O atom) = 32 electrons.
2. Calculate A, number of electrons available in the valence (outer) shells of all the
atoms. For negatively charged ions, add to this number the number of electrons
equal to the charge of the anions. For cations you subtract the number of electrons
equal to the charge on the cation.
For instance: NO3
-
,
A= 1×5(for N) +3×6 (for O atom) +1(for -1 charge) = 5+18+1=24 electrons.
3. Calculate S, total number of electrons shared in the molecule or ion, using the
relationship
S = n-A
S= n-A= 32-24 =8 electrons shared (4pairs of electron shared)
4. Place S electrons into the skeleton as shared pairs. Use double and triple bonds
only when necessary. Lewis formulas may be shown as either dot formula or dashformulas.
There are three general ways in which the octet rule doesn’t work:
• Molecules with an odd number of electrons
• Molecules in which an atom has less than an octet
• Molecules in which an atom has more than an octet
a. Odd number of electrons
Consider the example of the Lewis structure for the molecule nitrous oxide (NO):
Total electrons: 6+5=11
Bonding structure:
Octet on “outer” element is realized and on central atom only 3 electrons remain free(11-8 = 3).
4.3 Coordinate or dative covalent bonding and properties
4.3.1. Co-ordinate or dative covalent bonding and properties
A dative covalent bond, or coordinate bond is another type of covalent bonding. In this
case, the shared electron pair(s) are completely provided by one of the participants
in the union, and not by contributions from the two of them. The contributors of the
shared electrons are either neutral molecules which contain lone pair(s) of electrons
on one of their atoms, or negatively charged groups with free pairs of electrons todonate for sharing
The solid copper (II) hydroxide which was initially formed reacts with the excess
ammonia (which acts as ligands) to form the water soluble tetra ammine copper (II)complex as shown below.
4.4 Valence bond theory (VBT)
As you notice, the density of bonding electrons is not on the inter-nuclei axis, it is
rather located outside the axis but surrounding it. This kind of covalent bond is called
“ Pi bond”, represented by the symbol “π”. Hence the double bond O=O is made of
two covelent bonds: a σ bond and a π bond.
Due to the position of their electrons density in relation with the two nuclei, σ bond
participates in maintaining the two nuclei together more strongly than the π bond;
that is why σ bond is stronger than π bond. In addition, π bond cannot exist alone, it
exists only where there is a double or triple bond. Hence, in a double or triple bond,
there is one σ bond and one or two π bonds respectively.
Checking Up 4.4
1. Describe the aspects and postulates of valence bond theory(VBT)
2. Use VBT to explain the formation of single(sigma) and double (pi)bonds
(a) Explanation of lateral overlap of atomic orbitals and(b) Explanation of head-to-head overlap of atomic orbitals
4.5 Valence Shell Electron Pair Repulsion Theory (VSEPR)
theory
nucleus. Hence they occupy more space. As a result, the lone pairs cause more
repulsion.
The order of repulsion between different types of electron pairs is as follows:
Lone pair - Lone pair > Lone Pair - Bond pair > Bond pair - Bond pair
The bond pairs are usually represented by a solid line, whereas the lone pairs are
represented by a lobe with two electrons.
3) In VSEPR theory, the multiple bonds are treated as if they were single bonds.
The electron pairs in multiple bonds are treated collectively as a single super pair.
The repulsion caused by bonds increases with increase in the number of
bonded pairs between two atoms i.e., a triple bond causes more repulsion
than a double bond which in turn causes more repulsion than a single bond.
4) The shape of a molecule can be predicted from the number and type ofvalence shell electron pairs around the central atom.
The principle of the VSEPR is based on the idea that: the most stable structure of a
molecule is the one where the electron pairs are far away one from another in
order to minimize the repulsions between the pairs of electrons surrounding
the central atom.
The VSEPR theory assumes that each atom in a molecule will achieve a geometry
that minimizes the repulsion between electrons in the valence shell of that atom.
The use of VSEPR involves the following steps:
• Draw a Lewis structure for the ion or molecule in question.
• The shape is based on the location of the nuclei in a molecule, so double
and triple bonds count as one shared pair when determining the shape of
the molecule
• Locate the shared pairs and lone pairs on the central atom• Determine the shape based on the above considerations.
4.6. Hybridisation and types of Hybridisation
Activity 4.6
1. Write the electronic configuration of carbon and hydrogen using s,p, d.. Notation.
2. Use the electronic configurations above to identify the orbitals that contain
electrons used during the formation of methane.
3. Use the knowledge of overlap of atomic orbitals to indicate how orbitals overlap
in formation of hydrogen chloride, methane and beryllium chloride and predictthe shapes of the molecules.
4.7 Polar covalent bonds
Activity 4.7
1. Can you define the term electronegativity?
2. How is electronegativity related to polarity of the compound?
3. How does the polarity of a given molecule affect its physical properties?
4. Can you describe the general trends of electronegativity across and down
the groups in the periodic table?
5. What is meant by the term dipole and Net dipole
What happens if shairing of the bonding pair of electrons between the two atoms
forming the bond is not equal? For instance when two different non-metal elements
such as hydrogen and bromine combine?
In this case, there is unequal sharing where the more electronegative element takes
a bigger share of the bonding pair of electrons (Fig. 4.7
Figure 4.7: Polar covalent bond
(www.bbc.co.uk/bitesize/higher/chemistry/energy/bsp/revision/1/
In a polar covalent bond, binding pair of electrons is unequally shared between two
atoms. The power of an atom to attract the pair of electrons that constitutes the
bond in a molecule is called “electronegativity”.
The ‘electronegativity’ can be used to determine whether a given bond is nonpolar covalent, polar covalent or ionic bond. The electronegativity increases from left
to right across a period and decreases as you go down a group
The larger the electronegativity, the greater is the strength to attract a bonding pair
of electrons; and the larger the difference in electronegativites of the atoms, themore polar the covalent bond between the two atoms.
The following is the general thumb rule for predicting the type of bond based
upon the electronegativity differences:
• If the electronegativities are equal and the difference in electronegativity
difference is less than 0.5, the bond is non-polar covalent.
• If the difference in electronegativities between the two atoms is greater
than 0.4, but less than 2.0, the bond is polar covalent.
• If the difference in electronegativities between the two atoms is 2.0, or
greater, the bond is ionic
(ii) Poor electrical conductivity
There are no charged particles (ions or electrons) delocalized throughout the
molecular crystal lattice to conduct electricity. They cannot conduct electricity in
either the solid or molten state.
(iii) Solubility
Simple structures tend to be quite insoluble in water, but this depends on how the
polarized molecule is. The more polar the molecules, the more water molecules will
be attracted to them (some may dissolve in water as a result of forming hydrogen
bonds within it). Molecular crystals tend to dissolve in non-polar solvents such as
alcohol.
(iv) Soft and low density
Van der Waals forces are weak and non-directional. The lattice is readily destroyed
and the crystals are soft and have low density.
b. Giant covalent structures and their physical properties
Sometimes covalently bonded structures can form giant networks, known as Giant
Covalent Structures. In these structures, each network of bonds connects all the
atoms to each other.
These structures are usually very hard and have high melting and boiling points.
This is because of the strong covalent bonds holding each atom in place. In general,
Giant Covalent Structures cannot conduct electricity due to the fact that there are no
free charge carriers. One notable exception is Graphite. This is a structure composed
of ‘sheets’ of carbon atoms on top of each other. Electrons can move between the
sheets and carry the electricity. The main giant covalent molecular structures are the
two allotropes of carbon (diamond and graphite), and silica (silicon dioxide).
(i) Diamond structure and the physical properties
Diamond is a form of carbon in which each carbon atom is joined to four other
carbon atoms, forming a giant covalent structure with four single bonds. As a result,
diamond is very hard and has a high melting point. It does not conduct electricity.
Diamond is tetrahedral face-centered cubic as shown in the figure below
Diamond has a very high melting point (almost 4000°C): the carbon-carbon covalent
bonds are very strong and have to be broken throughout the structure before
melting occurs.
The compound is very hard due to the necessity to break very strong covalent bonds
operating in 3-dimensions.
Diamond does not conduct electricity: All the electrons are held tightly between the
atoms, and are not able to move freely.
The compound is insoluble in water and other organic solvents due to no possible
attractions which could occur between solvent molecules and carbon atoms which
could outweigh the attractions between the covalently bound carbon atoms.
(ii) Graphite and the physical properties
Graphite is another form of carbon in which the carbon atoms form layers. These
layers can slide over each other and graphite is much softer than diamond. Each
carbon atom in a layer is joined to only three other carbon atoms in hexagonal ringsas shown in the figure below
Silicon dioxide exhibits some physical properties such as:
• It has a high melting point (around 1700°C) which varies depending on
what the particular structure is (remember that the structure given is only
one of three possible structures).The silicon-oxygen covalent bonds are
very strong and have to be broken throughout the structure before the
melting occurs
• Silicon dioxide is hard due to the need to break the very strong covalent
bonds.
• Silicon dioxide is not displaying the property of electrical conductivity
because all the electrons are held tightly between the atoms, and are not
able to move freely. No any delocalized electrons are observed.
• It is insoluble in water and organic solvents because no possible attractions
occur between solvent molecules and the silicon or oxygen atoms which
could overcome the covalent bonds in the giant structure
4.
a) Draw a diagram to show the structure of silicon dioxide.
b) Explain why silicon dioxide
(i) is hard;
(ii) has a high melting point;
(iii) Doesn’t conduct electricity;(iv) is insoluble in water and other solvents.
4.9. Intermolecular Forces
Activity 4.9
1. Make a research and describe why:
i) Ice floats over water and the bottle full of water breaks on cooling(freezing)
ii) Water is a liquid at room temperature while Hydrogen sulfide is a gas
2. Trichloromethane (ii) ethanol (iii) aluminium fluoride. Arrange these compounds
in order of increasing boiling points.
Intermolecular forces are electrostatic forces which may arise from the interaction
between partial positively and negatively charged particles. Intermolecular forces
exist between two molecules while intramolecular forces hold atoms of a molecule
together in a molecule (Figure 4.11).
Intermolecular forces are much weaker than the intramolecular forces of attraction
but are important because they determine the physical properties of molecules
such as their boiling point, melting point, density, and enthalpies of fusion andvaporization.
Intramolecular forces hold the atoms in the molecule together; they are called
chemical bonds. Intermolecular forces hold covalent molecules together and are
responsible of a certain number of properties of the substance such as the melting
and boiling temperatures of covalent substances. They can be grouped in a category
of forces called van der Waals forces. There are three main kinds of intermolecular
interactions such as London dispersion forces, dipole-dipole interactions andhydrogen bonding later in the unit
dispersion forces than chlorine, contributing to increasing the boiling point of
bromine, 59 o
C, compared to chlorine, –35o
C. Those London forces are very weak
for non-polar covalent compounds; hence breaking them does not require much
energy, which explains why non-polar covalent compounds such as methane and
nitrogen which only have London dispersion forces of attraction between themolecules have very low melting and boiling points.
4.9.3. Hydrogen bond
For a hydrogen bond to be possible, there are necessary conditions:• The first condition is that the molecule contains one group where hydrogen is
Hydrogen bonds in DNA
(https://www.easynotescards.com/notecard_set/59549)
Hydrogen bonding in ice
Each water molecule is hydrogen-bonded to 4 others in a tetrahedral formation. The
ice molecule has a “diamond-like” structure. When liquid water freezes, the hydrogen
bonds become more rigid, and the volume becomes larger than the liquid because
of empty space generated by the rigidity of solid water. This explains why ice floatsover liquid water because its density is lower than the density of liquid water
4.10. End unit assessment
PART 1: MULTIPLE CHOICE QUESTIONS
1. In the periodic table, electronegativity generally decreases: A From right to left
in a period, B Upwards in a group; C From left to right in a period.
2. Which structure would sulphur, S8
, have?
A simple covalent molecules B simple covalent lattice C giant covalent lattice D
giant ionic lattice
3. Which statement(s) is/are true?
1) Water has hydrogen bonds which increase the boiling point.
2) Water as a solid is denser than as a liquid.
3) Water has bond angles of 180o
.
A 1, 2 and 3
B 1 and 2
C 2 and 3D only 1
a) The geometry of a molecule is determined by the number of electron groups
on the central atom.
b) The geometry of the electron groups is determined by minimizing repulsions
between them.
c) A lone pair, a single bond, a double bond, a triple bond and a single electron -
each of these is counted as a single electron group.
d) Bond angles may depart from the idealized angles because lone pairs of
electrons take up less space than bond pairs.
e) The number of electron groups can be determined from the Lewis structure ofthe molecule
PART 2: Filling in questions
16. Use the words listed below to fill in the correct appropriate word(s) in the
spaces below in the text.
Bigger, covalent bond, diamond, free electrons, halogens, hard crystals,
high electrical conductivity, high melting points, increase, intermolecular
forces, low electrical conductivity, low melting points, non-metals,
sharing, soft crystals, strong, strong bond, weak, weak force.
A ………………… is formed by two atoms…………….. one or more pairs of
electrons to make a …………………..between the two atoms in a molecule.
However, between small molecules, only a ………………holds them
together in the bulk liquid or solid. This results in small covalent molecules
having ……………………… and ………………………………..if solid. Small
covalent molecules have no …………………. and so have a
……………………………………
The Group 7 ………………………. collectively known as the…………………..
form diatomic molecules of two atoms. The …………………………..between
the molecules are ……………. giving them relatively low melting points and
boiling points. This also explains why they are gases, liquids or solids with
……………………. As you go down Group 7 the melting boiling points and
boiling points …………………..because the molecules get ………….. and the
intermolecular forces ………………….
In giant covalent structures the forces between all the atoms are
……………..forming ………………… like diamond or silica . In the
atomic giant structure metals there are free electrons which allow
…………………………………………….
PART3:
17. Fill in the table by putting a check mark in the compare-and-contrast matrixunder the column(s) that each physical attribute describes.
a) A 3-D, repeating pattern of + and – ions, formed by ionic compound
b)Tendency for an atom to attract the bonding pair electrons when chemically
bonded to another atom
c) A sharing of a pair of electrons
d) Atoms will gain or lose enough electrons in order to become isoelectronic with
a noble gas
d) A transfer of electrons from one atom to another
e) A chemical formula that is arranged in the smallest whole number ratio
f) The term that means dissolved in water
g) A chemical formula that describes the makeup of a single molecule
h) The shape (geometry) that is an exception to the octet rule
i) A bond where electrons are shared unequally between atoms
j) One of the shapes (geometries) that is polar
k) A bond where electrons are shared equally between atoms
Across
11. A covalent bond between atoms in which the electrons are shared unequally
12. A covalent bond in which the electrons are shared equally by the two atoms
14. Intermolecular forces resulting from the attraction of oppositely charged
regions of polar molecules
15. A bond formed when two atoms share a pair of electrons
16. The two weakest intermolecular attractions - dispersion interactions and
dipole forces
18. A covalent bond in which one atom contributes both bonding electrons
20. A chemical formula that shows the arrangement of atoms in a molecule or
polyatomic ion
21. A covalent bond in which three pairs of electrons are shared by two atoms
22. A bond in which two atoms share two pairs of electrons
23. A compound that is composed of molecules
26. A molecule consisting of two atoms
28. One of the two or more equally valid electron dot structures of a molecule or
polyatomic ion
31. valence-shell electron-pair repulsion theory; because electron pairs repel,
molecules adjust their shapes so that valence electron pairs are as far apart as
possible
Down
1. An orbital that applies to the entire molecule
2. A bond angle of 109.5 degrees that results when a central atom forms four
bonds directed toward the center of a regular tetrahedron
3. The mixing of several atomic orbitals to form the same total number of
equivalent hybrid orbitals
4. A tightly bound group of atoms that behaves as a unit and has a positive or
negative charge
5. A pair of valence electrons that is not shared between atoms
6. A molecule in which one side of the molecule is slightly negative and the
opposite side is slightly positive.
7.a covalent bond in which the bonding electrons are most likely to be found in
sausage-shaped regions above and below the bond axis of the bonded atoms
8. A covalent bond between atoms in which the electrons are shared unequally
9. A neutral group of atoms joined together by covalent bonds
10. A molecule that has two poles, or regions, with opposite charges
12. The energy required to break the bond between two covalently bonded atoms
17. A chemical formula of a molecular compound that shows the kinds and
numbers of atoms present in a molecule of a compound
19. Attractions between molecules caused by the electron motion on one
molecule affecting the electron motion on the other through electrical forces
24. Attractive forces in which hydrogen covalently bonded to a very electronegative
atom is also weakly bonded to an unshared electron pair of another electronegative
atom
25. A solid in which all of the atoms are covalently bonded to each other
27. A molecular orbital that can be occupied by two electrons of a covalent bond
29. A bond formed by the sharing of electrons between atoms
30. A bond formed when two atomic orbitals combine and form a molecular
orbital
that is symmetrical around the axis connecting the two atomic nuclei