Structure and bonding
Atomic structure
Atomic orbitals
The electronic structures of atoms
Chemical bonding
- Ionic (electrovalent) bonding
- Covalent bonding
- Single bond
- Double bonds
- Co-ordinate(dative covalent) bonding
Electronegativity
Shape of molecules and ions
Metallic bonding
Intermolecular bonding
- Van der waals forces
- Hydrogen bonds
Atomic structure
The sub-atomic particles
Protons, neutrons and electrons.
Relative mass
1
1
1/1836
Proton
Neutron
Electron
Relative charge
+1
0
-1
Isotopes, the number of neutrons in an atom can vary within small
limits.
Proton
Neutron
Mass number
Carbon-12
6
6
12
Carbon-13
6
7
13
number of electrons = number of protons
Atomic orbitals
What is an atomic orbital?
Orbitals and orbits
Impossibility of drawing orbits for electrons.To plot a path for something
you need to know exactly where the object is and be able to work out
exactly where it's going to be an instant later. (The Heisenberg
Uncertainty Principle)
Hydrogen's electron - the 1s orbital
2s orbital is similar to a 1s orbital except that the region where
there is the greatest chance of finding the electron is further
from the nucleus this is an orbital at the second energy level.
there is another region of slightly higher electron density nearer
the nucleus. "Electron density" is another way of talking about
how likely you are to find an electron at a particular place.
A p orbital is rather like 2 identical balloons tied together at the
nucleus
Fitting electrons into orbitals
The order of filling orbitals
Electrons fill low energy orbitals (closer to the nucleus) before they
fill higher energy ones. Where there is a choice between orbitals
of equal energy, they fill the orbitals singly as far as possible.
The electronic structures of atoms
Relating orbital filling to the Periodic Table
d8 means
Electronic structures of Cl 1s22s22p63s23px23py23pz1
ions Cl- 1s22s22p63s23px23py23pz2
Electronic structures of Na 1s22s22p63s1
ions Na+ 1s22s22p6
Electronic structures of Fe 1s22s22p63s23p63d64s2
ions Fe3+ 1s22s22p63s23p63d5
The electronic structures of atoms
Ionisation energies are always concerned with the formation of
positive ions.
Electron affinities are the negative ion equivalent, and their use is
almost always confined to elements in groups 6 and 7 of the
Periodic Table.
Chemical bonding
IONIC (ELECTROVALENT) BONDING
COVALENT BONDING
- SINGLE BONDS
- DOUBLE BONDS
CO-ORDINATE (DATIVE COVALENT) BONDING
IONIC (ELECTROVALENT) BONDING
Ionic bonding in sodium chloride
Sodium (2,8,1) has 1 electron more than a stable noble gas
structure (2,8). If it gave away that electron it would become more
stable.
Chlorine (2,8,7) has 1 electron short of a stable noble gas structure
(2,8,8). If it could gain an electron from somewhere it too would
become more stable.
The answer is obvious. If a sodium atom gives an electron to a
chlorine atom, both become more stable.
COVALENT BONDING- SINGLE BONDS
As well as achieving noble gas structures by transferring electrons
from one atom to another as in ionic bonding, it is also possible for
atoms to reach these stable structures by sharing electrons to give
covalent bonds.
Methane
What is wrong with the dots-and-crosses picture of bonding in
methane?
We are starting with methane because it is the simplest case which
illustrates the sort of processes involved. You will remember that
the dots-and-crossed picture of methane looks like this. There is a
serious mis-match between this structure and the modern
electronic structure of carbon, 1s22s22px12py1. The modern
structure shows that there are only 2 unpaired electrons for
hydrogens to share with, instead of the 4 which the simple view
requires.
You can see this more readily using the electrons-in-boxes
notation. Only the 2-level electrons are shown. The 1s2 electrons
are too deep inside the atom to be involved in bonding. The only
electrons directly available for sharing are the 2p electrons. Why
then isn't methane CH2?
Promotion of an electron
When bonds are formed, energy is released and
the system becomes more stable. If carbon forms
4 bonds rather than 2, twice as much energy is
released and so the resulting molecule becomes
even more stable.
There is only a small energy gap between the 2s
and 2p orbitals, and so it pays the carbon to
provide a small amount of energy to promote an
electron from the 2s to the empty 2p to give 4
unpaired electrons. The extra energy released
when the bonds form more than compensates for
the initial input.
Now that we've got 4 unpaired electrons ready for
bonding, another problem arises. In methane all
the carbon-hydrogen bonds are identical, but our
electrons are in two different kinds of orbitals.
You aren't going to get four identical bonds
unless you start from four identical orbitals.
sp3 Hybridisation
The electrons rearrange themselves again in a process called
hybridisation. This reorganises the electrons into four identical
hybrid orbitals called sp3 hybrids (because they are made from
one s orbital and three p orbitals). You should read "sp3" as "s p
three" - not as "s p cubed". sp3 hybrid orbitals look a bit like half a
p orbital, and they arrange themselves in space so that they are as
far apart as possible. You can picture the nucleus as being at the
centre of a tetrahedron (a triangularly based pyramid) with the
orbitals pointing to the corners. For clarity, the nucleus is drawn
far larger than it really is.
The bonding in the phosphorus chlorides, PCl3
Phosphorus has the electronic structure 1s22s22p63s23px13py13pz1.
There are 3 unpaired electrons that can be used to form bonds
with 3 chlorine atoms. The four 3-level orbitals hybridise to
produce 4 equivalent sp3 hybrids just like in carbon - except that
one of these hybrid orbitals contains a lone pair of electrons.
The bonding in the phosphorus chlorides, PCl5
Phosphorus
has
1s22s22p63s23px13py13pz1.
the
electronic
structure
COVALENT BONDING - DOUBLE BONDS
Hybridization sp2
Bonding in Benzene
The Kekulé structure for benzene, C6H6
Kekulé was the first to suggest a sensible structure for benzene.
The carbons are arranged in a hexagon, and he suggested
alternating double and single bonds between them. Each carbon
atom has a hydrogen attached to it.
Problems with the Kekulé structure
1. Problems with the chemistry, because of the three double bonds,
you might expect benzene to have reactions like ethene
2. Problems with the shape, benzene is a planar molecule (all the
atoms lie in one plane), and that would also be true of the Kekulé
structure. The problem is that C-C single and double bonds are
different lengths. C-C
0.154 nm, C=C, 0.134 nm
Bonding in Benzene
3. Problems with the stability of benzene, real benzene is a lot more
stable than the Kekulé structure would give it credit for.
This means that real benzene is about 150 kJ mol-1 more stable
than the Kekulé structure gives it credit for. This increase in
stability of benzene is known as the delocalisation energy or
resonance energy of benzene.
An orbital model for the benzene structure
Benzene is built from hydrogen atoms (1s1) and carbon atoms
(1s22s22px12py1).
Each carbon atom has to join to three other atoms (one hydrogen and
two
carbons) sp2 hybrids, because they are made by an s orbital and two p
orbitals
reorganising themselves. The three sp2 hybrid orbitals arrange
themselves as
far apart as possible - which is at 120° to each other in a plane. The
remaining
p orbital is at right angles to them
CO-ORDINATE (DATIVE COVALENT) BONDING
NH3 + HCl
NH4Cl
CO-ORDINATE (DATIVE COVALENT) BONDING
H2O+ HCl
H3O++Cl-
Measurements of the relative formula mass of aluminium chloride show
that its formula in the solid is not AlCl3, but Al2Cl6. It exists as a dimer
(two molecules joined together). The bonding between the two molecules
is co-ordinate, using lone pairs on the chlorine atoms
ELECTRONEGATIVITY
Definition
Electronegativity is a measure of the tendency of an atom to attract a
bonding pair
of electrons.The Pauling scale is the most commonly used. Fluorine (the
most
electronegative element) is assigned a value of 4.0, and values range down
to caesium
and francium which are the least electronegative at 0.7.
What happens if two atoms of equal electronegativity bond together?
What happens if B is slightly more electronegative than A?
What happens if B is a lot more electronegative than A?
Summary
 No electronegativity difference between two atoms leads to a pure non-polar
covalent bond.
 A small electronegativity difference leads to a polar covalent bond.
A large electronegativity difference leads to an ionic bond.
The SHAPES OF MOLECULES AND IONS
linear
trigonal planar
tetrahedral
trigonal bipyramid
pyramidal
octahedral.
bent or V-shaped.
square planar
The SHAPES OF MOLECULES AND IONS
linear
tetrahedral
trigonal planar
bent or V-shaped.
METALLIC BONDING
Metals tend to have high melting points and boiling points suggesting
strong bonds between the atoms. Even a metal like sodium (melting point
97.8°C) melts at a considerably higher temperature than the element
(neon) which precedes it in the Periodic Table.
Sodium has the electronic structure 1s22s22p63s1. When sodium atoms
come together, the electron in the 3s atomic orbital of one sodium atom
shares space with the corresponding electron on a neighbouring atom to
form a molecular orbital - in much the same sort of way that a covalent
bond is formed. The difference, however, is that each sodium atom is
being touched by eight other sodium atoms - and the sharing occurs
between the central atom and the 3s orbitals on all of the eight other
atoms. The electrons can move freely within these molecular orbitals, and
so each electron becomes detached from its parent atom. The electrons
are said to be delocalised. The metal is held together by the strong forces
of attraction between the positive nuclei and the delocalised electrons.
INTERMOLECULAR BONDING - VAN DER WAALS FORCES
Intermolecular versus intramolecular bonds
Intermolecular attractions are attractions between one molecule and a
neighbouring molecule. The forces of attraction which hold an individual
molecule together (for example, the covalent bonds) are known as
intramolecular attractions.
The origin of van der Waals dispersion forces
Electrons are mobile, and at any one instant they might find themselves
towards one end of the molecule, making that end d-. The other end will be
temporarily short of electrons and so becomes d +.Imagine a molecule which
has a temporary polarity being approached by one which happens to be
entirely non-polar just at that moment. 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
How molecular shape affects the strength of the dispersion
forces
The shapes of the molecules also matter. Long thin molecules can develop
bigger temporary dipoles due to electron movement than short fat ones
containing the same numbers of electrons.
Long thin molecules can also lie closer together - these attractions are at
their most effective if the molecules are really close.
For example, the hydrocarbon molecules butane and 2-methylpropane both
have a molecular formula C4H10, but the atoms are arranged differently. In
butane the carbon atoms are arranged in a single chain, but 2-methylpropane
is a shorter chain with a branch. Butane has a higher boiling point because
the dispersion forces are greater. The molecules are longer (and so set up
bigger temporary dipoles) and can lie closer together than the shorter, fatter
2-methylpropane molecules.
van der Waals forces: dipole-dipole interactions
A molecule like HCl has a permanent dipole because chlorine is more
electronegative than hydrogen. These permanent, in-built dipoles will cause
the molecules to attract each other rather more than they otherwise would if
they had to rely only on dispersion forces.
It's important to realise that all molecules experience dispersion forces.
Dipole-dipole interactions are not an alternative to dispersion forces - they
occur in addition to them. Molecules which have permanent dipoles will
therefore have boiling points rather higher than molecules which only have
temporary fluctuating dipoles.
INTERMOLECULAR BONDING - HYDROGEN BONDS
The evidence for hydrogen bonding
Many elements form compounds with hydrogen - referred to as
"hydrides". If you plot the boiling points of the hydrides of the Group 4
elements, you find that the boiling points increase as you go down the
group.
The increase in boiling point happens because the molecules are getting
larger with more electrons, and so van der Waals dispersion forces
become greater.
INTERMOLECULAR BONDING - HYDROGEN BONDS
If you repeat this exercise with the hydrides of elements in Groups 5, 6
and 7, something odd happens.
Although for the most part the trend is exactly the same as in group 4
(for exactly the same reasons), the boiling point of the hydride of the first
element in each group is abnormally high.
In the cases of NH3, H2O and HF there must be some additional
intermolecular forces of attraction, requiring significantly more heat
energy to break. These relatively powerful intermolecular forces are
described as hydrogen bonds.
The origin of hydrogen bonding
The molecules which have this extra bonding are:


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 "active" lone pair.
Lone pairs at the 2-level have the electrons contained in a relatively
small volume of space which therefore has a high density of negative
charge. Lone pairs at higher levels are more diffuse and not so attractive
to positive things.
The origin of hydrogen bonding
Consider two water molecules coming close together.
The d+ hydrogen is so strongly attracted to the lone pair that it is almost
as if you were beginning to form a co-ordinate (dative 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 being constantly broken and reformed in liquid
water
Hydrogen bond effect
The boiling points of ethanol and methoxymethane show the dramatic
effect that the hydrogen bonding has on the stickiness of the ethanol
molecules:
ethanol (with hydrogen bonding) 78.5°C
methoxymethane (without hydrogen bonding)-24.8°C
It is important to realise that hydrogen bonding exists in addition to van
der Waals attractions. Comparing the two alcohols (containing -OH
groups), both boiling points are high because of the additional hydrogen
bonding due to the hydrogen attached directly to the oxygen - but they
aren't the same.
The boiling point of the 2-methylpropan-1-ol isn't as high as the butan-1ol because the branching in the molecule makes the van der Waals
attractions less effective than in the longer butan-1-ol.
Hydrogen bond effect
Hydrogen bonding in organic molecules containing nitrogen
Hydrogen bonding also occurs in organic molecules containing N-H
groups - in the same sort of way that it occurs in ammonia. Examples
range from simple molecules like CH3NH2 (methylamine) to large
molecules like proteins and DNA. The two strands of the famous alphahelix in DNA are held together by hydrogen bonds involving N-H groups.