CHEMICAL BONDS
INTERMOLECULAR FORCES are attractive forces between molecules.
The attractions between
molecules are not nearly as
strong as the intramolecular
attractions
that
hold
compounds together.
They are, however, strong
enough to control physical
properties such as boiling
and melting points, vapor
pressures, and viscosities.
Generally, intermolecular forces are weaker than intramolecular forces.
While intramolecular forces stabilize individual molecules, intermolecular forces are
responsible for the bulk properties of matter e.g. melting and boiling points.
Examples:
41 kJ of energy is required to vapourise 1 mole of H 2 O at its boiling
point.
While 930 kJ of energy is required to break 2 x O-H bonds in H 2 O
These intermolecular forces as a group are referred to as van der Waals forces.
Examples of van der Waals forces are:
• Dipole-dipole interactions
• Hydrogen bonding
• London dispersion forces
Ion-Dipole Interactions
•
•
A fourth type of force, ion-dipole interactions are an important force in solutions
of ions.
The strength of these forces are what make it possible for ionic substances to
dissolve in polar solvents.
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Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
 The positive end of one is attracted to the negative end of the other and
vice-versa.
 These forces are only important when the molecules are close to each
other.
The more polar the molecule, the higher is its boiling point.
London Dispersion Forces
While the electrons in the 1s orbital of helium would repel each other (and, therefore,
tend to stay far away from each other), it does happen that they occasionally wind up on
the same side of the atom.
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At that instant, then, the helium atom is polar, with an excess of electrons on the left side
and a shortage on the right side
Another helium nearby, then, would have a dipole induced in it, as the electrons on the
left side of helium atom 2 repel the electrons in the cloud on helium atom 1.
London dispersion forces, or dispersion forces, are attractions between an instantaneous
dipole and an induced dipole.
•
•
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.
Factors Affecting London Forces
•
•
The shape of the molecule affects the strength of dispersion forces: long, skinny
molecules (like n-pentane tend to have stronger dispersion forces than short, fat
ones (like neopentane).
This is due to the increased surface area in n-pentane.
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•
•
The strength of dispersion forces tends to increase with increased molecular
weight.
Larger atoms have larger electron clouds, which are easier to polarize.
WHICH HAVE A GREATER EFFECT:
dipole-dipole interactions or dispersion forces
•
•
if two molecules are of comparable size and shape, dipole-dipole interactions will
likely be the dominating force.
if one molecule is much larger than another, dispersion forces will likely
determine its physical properties.
HYDROGEN BONDING
•
•
•
•
The dipole-dipole interactions experienced when H is bonded to N, O, or F are
unusually strong.
We call these interactions hydrogen bonds.
Hydrogen bonding arises in part from the high electronegativity of nitrogen,
oxygen, and fluorine.
Also, when hydrogen is bonded to one of those very electronegative elements, the
hydrogen nucleus is exposed.
•
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VISCOSITY
•
•
•
Resistance of a liquid to flow is called viscosity.
It is related to the ease with which molecules can move past each other.
Viscosity increases with stronger intermolecular forces and decreases with higher
temperature.
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SURFACE TENSION
Surface tension results from the net inward force experienced by the molecules on the
surface of a liquid.
INTRAMOLECULAR FORCES hold atoms together in a molecule.
While intramolecular forces stabilize individual molecules, intermolecular forces are
responsible for the bulk properties of matter e.g. melting and boiling points.
Examples:
41 kJ of energy is required to vapourise 1 mole of H 2 O at its boiling
point.
While 930 kJ of energy is required to break 2 x O-H bonds in H 2 O
Intramolecular forces include:
(A) Ionic bonding
(B) Covalent Bonding
(A) IONIC BONDING: is the electrostatic interaction of oppositely charged
ions in an ionic compound e.g. NaCl (m.pt. 104oC), LiF, CaO, MgO
(1100oC), KCl, KBr.
Takes place between cations and anions:
e.g.
Na+ + Cl- → NaCl
Li+ + Cl- → LiCl
Ca2+ + O2- → CaO
Elements most likely to form cations in ionic compounds are alkali metals (group I) and
alkaline earth metals (group II).
Elements most likely to form anions in ionic compounds are halogens (Group VII) and
oxygen.
Elements with low ionization energy (IE) – from cations
Elements with high electron affinity (EA) – from anions
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Recall:
Electronegativity: is the ability of an atom to attract electrons toward itself in a
chemical bond.
e.g. in H-F,
EN of F = 4.0
Generally the halogens have high EN values (F = 4.0, Cl = 3.0, Br = 2.8)
On the periodic chart, electronegativity increases as you go…
 …from left to right across a row.
 …from the bottom to the top of a column.
As a general rule an ionic bond forms when the EN difference between the 2
bonding atoms is 2 or more.
Electron Affinity: is the energy change that occurs when an electron is accepted by an
atom in the gaseous state.
e.g. X (g) + e- → X- (g)
Cl 2 (g) + 2e- → 2Cl-
Elements most likely to form cations in ionic compounds are alkali metals (group I) and
alkaline earth metals (group II).
Elements most likely to form anions in ionic compounds are halogens (Group VII) and
oxygen.
Elements with low ionization energy (IE) – from cations
Elements with high electron affinity (EA) – from anions
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(B) Covalent Bonding
Covalent bond: a bond in which two electrons are shared by two atoms
Covalent compounds are compounds that contain only covalent bonds e.g. HCl, NH 3 ,
CH 4 , CCl 4 , H 2 O.
There are several electrostatic interactions in these bonds:
 Attractions between electrons and nuclei
 Repulsions between electrons
 Repulsions between nuclei
Polar Covalent Bonds
•
•
Fluorine pulls harder on
the electrons it shares
with hydrogen than
hydrogen does.
Therefore, the fluorine
end of the molecule has
more electron density
than the hydrogen end.
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Although atoms often form compounds by sharing electrons, the electrons are not always
shared equally
•
•
•
When two atoms share electrons unequally, a bond dipole results.
•
The dipole moment, µ, produced by two equal but opposite charges separated by a
distance, r, is calculated:
µ = Qr
It is measured in debyes (D).
Covalent bonding between many-electron atoms involves only the valence electrons.
Let us consider the F 2 molecule:
It involves the covalent bonding of 2 F atoms.
2 2
5
The electronic configuration of F is 1s 2s 2p
The 1s electrons are low in energy and stay near the nucleus, hence do not participate in
2
5
bonding. Each F has 7 valence electrons: 2s 2p = 2 + 5= 7
F
+
F F
F
A LEWIS structure is a representation of covalent bonding where shared electrons are
shown as lines or dots as shown in the F 2 molecule above.
Example: CO 2
O C O
O C O
Other non-bonding electrons are called lone pairs
H O H
Lone pair of electrons
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The Octet Rule: An atom other that hydrogen tends to form bonds until it is
surrounded by 8 valence electrons.
The Octet rule works mainly for elements in period 2 . Why?
The following molecules illustrates the Octet Rule
The F 2 molecule
F F
The N 2 molecule
N N
Each atom is surrounded by 8 electrons in each case: Octet Rule
Exercise 1.:
Draw the Lewis Dot Structure for the CCl 4 molecule
Solution:
Step 1. Count up all the valence electrons first
= 4 eC = group 4 = 4 eCl = group 7 = 4 x 7e- = 28 eTotal electrons
32 eStep 2. Insert the bonding electrons
Cl
Cl
C Cl
Cl
A total of 10 electrons are used in the 5 bonds
Cl
Step 3. Distribute the remaining electrons around substituent atoms i.e. complete
the Octet around substituent groups
Cl
Cl
Cl
C Cl
Cl
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Step 4. Place any remaing electrons around the central atom
Exercise 2. Draw the Lewis Dot Structure for the H 2 SO 4 molecule (S is the central atom)
Atoms can form different types of covalent bonds
(a) A Single bond: 2 atoms held together by one electron pair e.g C:C or C-C.
(b) A Double bond: two atoms are held together or share 2 pairs of electrons e.g. C=C.
(c) Triple bond: arises when two atoms share 3 pairs of electrons e.g.
C
C
Double and triple bonds are Multiple bonds which are shorter than single bonds.
Bond length: is the distance between the nuclei of two covalently bonded atoms in a
molecule.
e.g.
C-H 107 pm (picometers)
C-O 143 pm
C=O 121 pm
C-C 154 pm
C≡C 120 pm
Comparison between Ionic and Covalent Compounds
They differ markedly in physical properties because of differences in the nature of their
bonds.
There are two types of attractive forces in covalent compounds i.e. intermolecular and
intramolecular forces.
Covalent compounds are usually gases, liquids or low melting solids at room temp.
Ionic compounds are solids at room with high melting points.
Most ionic compounds are soluble in water, conduct electricity hence are strong
electrolytes.
Most covalent compounds are insoluble in water or if they do dissolve, they do not
conduct electricity.
Liquid or molten covalent compounds do not conduct electricity because no ions are
present.
Molten ionic compounds conduct electricity because they contain mobile cations and
anions.
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Example.
Bond type
Compound
Appearance
Melting point (M.pt.)
Boiling point (B.pt.)
Solubility in H 2 O
Ionic
NaCl
White solid
801oC
1413oC
High
Covalent
CCl 4
Colorless liquid
-23oC
76 oC
Very low
Poor
Good
Poor
Poor
Electrical conductivity of:
(a) Solid
(b) Liquid
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