Intermolecular Bonds and Relative Strengths
Types of intermolecular
bonds
Key Learning
•
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-
dipole attraction and hydrogen bonding) and evidence and factors that
determine bond strength including explanations for the floating of ice and
expansion of water at higher temperatures.
Types of intermolecular force
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There are three main types of intermolecular force:
The molecules in simple covalent substances are not entirely isolated from one another.
There are forces of attraction between them. These are called
intermolecular forces
.
Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the molecules may align so there
is
electrostatic attraction
between the opposite charges on neighbouring molecules.
Permanent dipole–dipole
forces
(dotted lines) occur
in hydrogen chloride (HCl)
gas.
The permanent dipole–dipole forces are approximately one hundredth the strength
of a covalent bond.
•
the relative strengths of bonds (
covalent bonding, dispersion forces,
dipole-dipole attraction
and hydrogen
bonding)
and evidence and factors that determine bond strength
including explanations for the floating of ice
and expansion of water at higher temperatures.
Permanent dipole–dipole or not?
Hydrogen bonding
Notice that in each of these molecules:
•
The hydrogen is attached directly to one of the most electronegative elements, causing
the hydrogen to acquire a significant amount of positive charge.
•
Each of the elements to which the hydrogen is attached is not only significantly
negative, but also has at least one lone pair.
•
Must be a Nitrogen, Oxygen or Fluorine bonded to Hydrogen to mean that Hydrogen
bonds become possible
•
the relative strengths of bonds
(covalent bonding, dispersion forces, dipole-dipole attraction and
hydrogen
bonding) and evidence and factors that determine bond strength
including explanations for the floating of ice
and expansion of water at higher temperatures.
What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than
in other polar bonds.
This is because these atoms are
highly electronegative
due to their high core charge and small size. When
these atoms bond to hydrogen, electrons are
withdrawn from the H atom, making it slightly positive.
Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole
forces.
The H atom is very small so the positive charge is more concentrated, making it easier to
link with other molecules.
Hydrogen bonding
•
The + hydrogen is so strongly attracted to the lone pair that
it is almost as if you were beginning to form a covalent bond.
It doesn't go that far, but the attraction is significantly
stronger than an ordinary dipole-dipole interaction.
•
Hydrogen bonds have about a tenth of the strength of an
average covalent bond and are about 10 times stronger
than a dipole-dipole bond,
and are being constantly broken
and reformed in liquid water.
•
Each water molecule can potentially form four hydrogen
bonds with surrounding water molecules, every one of them
can be involved in hydrogen bonding.
•
This is why the boiling point of water is higher than that of
ammonia or hydrogen fluoride.
•
the relative strengths of bonds
(covalent bonding, dispersion forces, dipole-dipole attraction and
hydrogen
bonding) and evidence and factors that determine bond strength
including explanations for the floating of ice and
expansion of water at higher temperatures.
Hydrogen bonding
In molecules with OH or NH
groups, a lone pair of electrons on
nitrogen or oxygen is attracted to
the slight positive charge on the
hydrogen on a neighbouring
molecule.
Hydrogen bonding makes the melting and boiling points of water higher than might be
expected. It also means that alcohols have much higher boiling points than alkanes of
a similar size.
•
the relative strengths of bonds
(covalent bonding, dispersion forces, dipole-dipole attraction and
hydrogen bonding) and evidence and factors that determine bond strength
including explanations for
the floating of ice and expansion of water at higher temperatures.
Hydrogen bonding and boiling points
Boiling points of the hydrogen halides
The boiling point of
hydrogen fluoride is much
higher than that of other
hydrogen halides, due to
fluorine’s high
electronegativity.
The means that
hydrogen bonding between molecules of hydrogen fluoride is much stronger
than the permanent dipole–dipole forces between molecules of other hydrogen halides
. More
energy is therefore required to separate the molecules of hydrogen fluoride.
Permanent dipole–dipole forces
How do we explain the existence of intermolecular forces in non polar
substances?
We know that intermolecular forces are present in non-polar substance
because we see non-polar substances form liquids and solids. Without
intermolecular forces there would be nothing to hold the molecules
together and non-polar substances would only exist as gases. Lots of non-
polar substances are liquids at room temperature (vegetable oil) and even
non-polar solids (candle wax). And all non polar substances can form
liquids or solids if cooled to low enough temperatures.
So how to we explain these forces??
Dispersion forces
Dispersion forces
•
Attractions are electrical in nature. In a symmetrical
molecule like hydrogen, however, there doesn't seem
to be any electrical distortion to produce positive or
negative parts. But that's only true on average.
•
The diagram represents a small symmetrical molecule
- H
2
, perhaps, or Br
2
. The even shading shows that on
average there is no electrical distortion.
•
But the electrons are mobile, and at any one instant
they might find themselves towards one end of the
molecule, making that end -. The other end will be
temporarily short of electrons and so becomes +.
•
An instant later the electrons may well have moved up
to the other end, reversing the polarity of the
molecule.
•
This constant moving around of the electrons in the
molecule causes rapidly fluctuating dipoles even in the
most symmetrical molecule. It even happens in
monatomic molecules - molecules of noble gases, like
helium, which consist of a single atom.
How do Dispersion Forces Work?
•
Imagine a molecule that has a temporary dipole alreday
set up, is being approached by a molecule with no
polarity.
•
As the right hand molecule approaches, its electrons will
tend to be attracted by the slightly positive end of the
left hand one.
•
This sets up an
induced dipole
in the approaching
molecule, which is orientated in such a way that the +
end of one is attracted to the - end of the other.
•
An instant later the electrons in the left hand molecule
may well have moved up the other end. In doing so, they
will repel the electrons in the right hand one.
•
The polarity of both molecules reverses, but you still
have + attracting -. As long as the molecules stay close to
each other the polarities will continue to fluctuate in
synchronisation so that the attraction is always
maintained.
•
This diagram shows how a whole lattice of molecules
could be held together in a solid using van der Waals
forces in one instant.
Strength of dispersion forces
•
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen
bonding) and
evidence and factors that determine bond strength
including explanations for the floating of ice
and expansion of water at higher temperatures.
Strength of dispersion forces
Straight chain alkanes can pack closer together than branched alkanes, creating more
points of contact between molecules. This results in stronger van der Waals forces.
The points of contact between molecules also affects the strength of dispersion
forces.
•
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen
bonding) and
evidence and factors that determine bond strength
including explanations for the floating of ice
and expansion of water at higher temperatures.
Boiling points of alkanes
Understanding the types of intermolecular bonds - covalent bonding, dispersion forces, dipole-dipole attraction, and hydrogen bonding, along with their relative strengths and factors determining bond strength. Learn about permanent dipole-dipole forces and hydrogen bonding, crucial for phenomena like ice floating and water expansion at higher temperatures.
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Presentation Transcript
Types of intermolecular bonds Key Learning the relative strengths of bonds (covalent bonding, dispersion forces, dipole- dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures.
Types of intermolecular force The molecules in simple covalent substances are not entirely isolated from one another. There are forces of attraction between them. These are called intermolecular forces. There are three main types of intermolecular force: dispersion forces also known as van der Waals forces for example, found between I2 molecules in iodine crystals. permanent dipole dipole forces for example, found between HCl molecules in hydrogen chloride. hydrogen bonds for example, found between H2O molecules in water.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Permanent dipole dipole forces If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules. Permanent dipole dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas. The permanent dipole dipole forces are approximately one hundredth the strength of a covalent bond.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Hydrogen bonding Notice that in each of these molecules: The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one lone pair. Must be a Nitrogen, Oxygen or Fluorine bonded to Hydrogen to mean that Hydrogen bonds become possible
What is hydrogen bonding? When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds. This is because these atoms are highly electronegative due to their high core charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive. The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules. Hydrogen bonds are therefore particularly strong examples of permanent dipole dipole forces.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Hydrogen bonding The + hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a covalent bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction. Hydrogen bonds have about a tenth of the strength of an average covalent bond and are about 10 times stronger than a dipole-dipole bond, and are being constantly broken and reformed in liquid water. Each water molecule can potentially form four hydrogen bonds with surrounding water molecules, every one of them can be involved in hydrogen bonding. This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Hydrogen bonding In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule. hydrogen bond lone pair Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.
Boiling points of the hydrogen halides 40 20 boiling point ( C) 0 -20 The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine s high electronegativity. -40 -60 -80 -100 HF HCl HBr HI The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.
How do we explain the existence of intermolecular forces in non polar substances? We know that intermolecular forces are present in non-polar substance because we see non-polar substances form liquids and solids. Without intermolecular forces there would be nothing to hold the molecules together and non-polar substances would only exist as gases. Lots of non- polar substances are liquids at room temperature (vegetable oil) and even non-polar solids (candle wax). And all non polar substances can form liquids or solids if cooled to low enough temperatures. So how to we explain these forces??
Dispersion forces Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average. The diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion. But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end -. The other end will be temporarily short of electrons and so becomes +. An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule. This constant moving around of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in monatomic molecules - molecules of noble gases, like helium, which consist of a single atom.
How do Dispersion Forces Work? Imagine a molecule that has a temporary dipole alreday set up, is being approached by a molecule with no polarity. As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one. This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the + end of one is attracted to the - end of the other. An instant later the electrons in the left hand molecule may well have moved up the other end. In doing so, they will repel the electrons in the right hand one. The polarity of both molecules reverses, but you still have + attracting -. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the attraction is always maintained. This diagram shows how a whole lattice of molecules could be held together in a solid using van der Waals forces in one instant.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Strength of dispersion forces 200 The strength of dispersion forces increases as molecular size increases. 150 boiling point ( C) 100 50 0 -50 -100 -150 -200 This is illustrated by the boiling points of group 7 elements. F2 Cl2 Br2 I2 element Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce.
the relative strengths of bonds (covalent bonding, dispersion forces, dipole-dipole attraction and hydrogen bonding) and evidence and factors that determine bond strength including explanations for the floating of ice and expansion of water at higher temperatures. Strength of dispersion forces The points of contact between molecules also affects the strength of dispersion forces. butane (C4H10) boiling point = 272K 2-methylpropane (C4H10) boiling point = 261K Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.
