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S AND P BLOCK ELEMENTS GUIDE
NOTES
Contents
ATOMIC ORBITALS
...
5
COVALENT BONDING - SINGLE BONDS
...
11
METALLIC BONDING
...
17
IONIC STRUCTURES
...
21
ELECTRON AFFINITY
...
28
SUCCESSIVE IONISATION ENERGIES
...
38
TRENDS IN MELTING AND BOILING POINTS
...
39
REACTIONS OF THE GROUP 1 ELEMENTS WITH OXYGEN AND CHLORINE
...
43
REACTIONS OF THE GROUP 1 ELEMENTS WITH WATER
...
44
THE SOLUBILITY OF GROUP 1 COMPOUNDS
...
48
REACTIONS OF THE GROUP 2 ELEMENTS WITH AIR OR OXYGEN
...
53
EXPLANATIONS FOR THE TRENDS IN SOLUBILITY OF SOME GROUP 2 COMPOUNDS
...
57
IN WATER
...
58
CHEMICAL REACTIONS OF THE PERIOD 3 ELEMENTS
...
67
ATOMIC AND PHYSICAL PROPERTIES OF THE PERIOD 3 ELEMENTS
...
76
PROPERTIES OF THE PERIOD 3 "HYDROXIDES"
...
82
OXIDATION STATE TRENDS IN GROUP 4
...
91
THE CHLORIDES OF CARBON, SILICON AND LEAD
...
100
1

THEORIES OF ACIDS AND BASES
...
A simple view of the
atom looks similar and you may have pictured the electrons as orbiting around the nucleus
...

The 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 says - loosely - that you can't know with certainty both where an electron is and where
it's going next
...
)
That makes it impossible to plot an orbit for an electron around a nucleus
...
If something is
impossible, you have to accept it and find a way around it
...
This is just
for clarity
...

Soon afterwards, you do the same thing, and find that it is in a new position
...

You keep on doing this over and over again, and gradually build up a sort of 3D map of the places that the
electron is likely to be found
...
The diagram
shows a cross-section through this spherical space
...
Such a region of space is called an orbital
...

Each orbital has a name
...
The "1" represents the fact
that the orbital is in the energy level closest to the nucleus
...
s orbitals are spherically symmetric around the nucleus - in each case, like a hollow
ball made of rather chunky material with the nucleus at its centre
...
This 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
...
("Electron density" is another way of talking about how likely you are to find an electron at a
particular place
...
The
effect of this is to slightly reduce the energy of electrons in s orbitals
...

3s, 4s (etc) orbitals get progressively further from the nucleus
...
At the first energy level, the
only orbital available to electrons is the 1s orbital, but at the second level, as well as a 2s orbital, there are
also orbitals called 2p orbitals
...
The diagram on the right is a cross-section through
that 3-dimensional region of space
...

Unlike an s orbital, a p orbital points in a particular direction - the one drawn points up and down the page
...
These are arbitrarily given the symbols px, py and pz
...

The p orbitals at the second energy level are called 2px, 2py and 2pz
...

All levels except for the first level have p orbitals
...


d and f orbitals

3

In addition to s and p orbitals, there are two other sets of orbitals which become available for electrons to inhabit at higher
energy levels
...
At the third level there are a total of nine orbitals altogether
...
s, p, d and f
orbitals are then available at all higher energy levels as well
...
On the first floor there is only 1 room (the 1s
orbital); on the second floor there are 4 rooms (the 2s, 2px, 2py and 2pz orbitals); on the third floor there are 9 rooms (one 3s
orbital, three 3p orbitals and five 3d orbitals); and so on
...
Each orbital can only hold 2
electrons
...
Often an up-arrow and a down-arrow are
used to show that the electrons are in some way different
...
If they live in
different orbitals, that's fine - but if they are both in the same orbital there has to be some subtle distinction between them
...

A 1s orbital holding 2 electrons would be drawn as shown on the right, but it can be written even more quickly as 1s2
...

You mustn't confuse the two numbers in this notation:

The order of filling orbitals
Electrons fill low energy orbitals (closer to the nucleus) before they fill higher energy ones
...

This filling of orbitals singly where possible is known as Hund's rule
...

The diagram (not to scale) summarises the energies of the orbitals up to the 4p level
...

The real oddity is the position of the 3d orbitals
...
Similar confusion occurs at higher levels, with so
much overlap between the energy levels that the 4f orbitals don't fill until after the 6s, for example
...
The particles may be arranged
regularly (in which case, the solid is crystalline), or at random (giving waxy solids like candles or some forms of polythene,
for example)
...
Liquids
In a liquid, the particles are mainly touching, but some gaps have appeared in the structure
...
Unless melting has broken up a substance consisting only of covalent
bonds (a giant covalent structure), the forces that held the solid particles together are also present in the liquid, but in a
somewhat loosened form
...
At ordinary pressures, the distance between individual particles is of the order
of ten times the diameter of the particles
...


5

Deducing the type of bonding from physical properties
The physical state and other properties
Melting point isn't always a good guide to the size of the attractions between particles, because the attractive forces have only
been loosened on melting - not broken entirely
...
The stronger the attractions, the higher the boiling point
...
Those oddities usually disappear if you consider boiling points instead
...
The boiling points reflect this: Al 2470°C, Mg
1110°C
...


So
...

The size of the melting point or boiling point gives a guide to the strength of the intermolecular forces
...

If it is a high melting point solid, it will be a giant structure - either ionic, metallic or giant covalent
...
If the substance also undergoes electrolysis when it is molten, that
would confirm that it was ionic
...
The clue
as to which you had would usually come from other data - appearance, malleability, etc
...
These noble gas structures are thought of as being in
some way a "desirable" thing for an atom to have
...

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
...


The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all
the electrons come from
...

The two chlorine atoms are said to be joined by a covalent bond
...
Hydrogen

Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium
...

Hydrogen chloride
The hydrogen has a helium structure, and the chlorine an argon structure
...
In this case, only
the outer electrons are shown for simplicity
...
Again, everything present has a noble gas structure
...
Is this a problem? No
...


7

Energy is released whenever a covalent bond is formed
...
It follows, therefore, that an atom will tend to make as many covalent bonds as
possible
...


Note: You might perhaps wonder why boron doesn't form ionic bonds with fluorine instead
...


Phosphorus(V) chloride, PCl5
In the case of phosphorus 5 covalent bonds are possible - as in PCl5
...
When phosphorus burns in chlorine both are
formed - the majority product depending on how much chlorine is available
...

The diagram of PCl5 (like the previous diagram of PCl3) shows only the outer electrons
...

A more sophisticated view of covalent bonding
The bonding in methane, CH4
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
...

There is a serious mis-match between this structure and the modern electronic structure of
carbon, 1s22s22px12py1
...

You can see this more readily using the electrons-in-boxes notation
...
The 1s2 electrons are too deep inside the atom to be involved in bonding
...
Why then isn't methane CH2?

8

Promotion of an electron
When bonds are formed, energy is released and the system becomes 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
...


Now that we've got 4 unpaired electrons ready for bonding, another problem arises
...

You aren't going to get four identical bonds unless you start from four identical orbitals
...
This reorganises the
electrons into four identical hybrid orbitals called sp3 hybrids (because they are made from one
s orbital and three p orbitals)
...

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
...
For clarity,
the nucleus is drawn far larger than it really is
...
When a covalent bond is formed, the atomic orbitals (the
orbitals in the individual atoms) merge to produce a new molecular orbital which contains the electron pair which creates the
bond
...
Each orbital holds the 2 electrons that we've previously drawn as a dot and a cross
...
The bonding in the phosphorus chlorides, PCl3 and PCl5
What's wrong with the simple view of PCl3?
This diagram only shows the outer (bonding) electrons
...
) If you were going to take a more modern look at it, the argument would go like this:

Phosphorus has the electronic structure 1s22s22p63s23px13py13pz1
...
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
...

You might wonder whether all this is worth the bother! Probably not! It is worth it with
PCl5, though
...
This diagram also shows only the outer
electrons
...

If the phosphorus is going to form PCl5 it has first to generate 5 unpaired electrons
...

Which higher energy orbital? It uses one of the 3d orbitals
...
Not so! Apart from when you
are building the atoms in the first place, the 3d always counts as the lower
energy orbital
...
They would be called sp3d hybrids because that's what they are made
from
...

Why does phosphorus form these extra two bonds? It puts in an amount of energy to promote an electron, which is more than
paid back when the new bonds form
...

The advantage of thinking of it in this way is that it completely ignores the question of whether you've got a noble gas
structure, and so you don't worry about it
...
In fact, it doesn't
...

Nitrogen is 1s22s22px12py12pz1
...
The problem is that there aren't any 2d orbitals to promote an electron into - and the energy
gap to the next level (the 3s) is far too great
...

Atoms will form as many bonds as possible provided it is energetically profitable
...
The atoms are held together because the electron pair is
attracted by both of the nuclei
...

A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons
come from the same atom
...


11

Ammonium ions, NH4+, are formed by the transfer of a hydrogen ion from the hydrogen chloride to the lone pair of electrons
on the ammonia molecule
...
The hydrogen's electron is left behind on the chlorine to
form a negative chloride ion
...
Although the electrons are shown differently in the diagram, there is no difference between them in reality
...
The arrow points from the atom donating the lone pair to the
atom accepting it
...
A hydrogen ion (H+) is transferred from the chlorine to one of the lone pairs on the oxygen atom
...

The reaction between ammonia and boron trifluoride, BF3
If you have recently read the page on covalent bonding, you may remember boron trifluoride as a compound which doesn't
have a noble gas structure around the boron atom
...
BF3 is described as being electron deficient
...


Using lines to represent the bonds, this could be drawn more simply as:

The second diagram shows another way that you might find co-ordinate bonds drawn
...
We shan't use this method again - it's more confusing than just using an arrow
...
If it simply contained ions it would have a
very high melting and boiling point because of the strong attractions between the positive and negative ions
...
The dots-and-crosses diagram shows only
the outer electrons
...
There is likely to be a similarity, because aluminium and boron are in the same group of
the Periodic Table, as are fluorine and chlorine
...
It exists as a dimer (two molecules joined together)
...
Each chlorine atom has 3 lone pairs, but only the two
important ones are shown in the line diagram
...

The bonding in hydrated metal ions
Water molecules are strongly attracted to ions in solution - the water molecules clustering around the positive or negative
ions
...

Ions with water molecules attached are described as hydrated ions
...
Six water molecules bond to the
aluminium to give an ion with the formula Al(H2O)63+
...

The bonding in this (and the similar ions formed by the great majority of other metals) is co-ordinate (dative covalent) using
lone pairs on the water molecules
...
When it forms an Al3+ ion it loses the 3-level electrons to leave 1s22s22p6
...
The aluminium re-organises (hybridises) six of these (the 3s, three 3p,
and two 3d) to produce six new orbitals all with the same energy
...

You might wonder why it chooses to use six orbitals rather than four or eight or whatever
...
By making the maximum number
of bonds, it releases most energy and so becomes most energetically stable
...
The other lone pair is pointing away from the aluminium and so isn't
involved in the bonding
...

Two more molecules
Carbon monoxide, CO
Carbon monoxide can be thought of as having two ordinary covalent bonds between the carbon and
the oxygen plus a co-ordinate bond using a lone pair on the oxygen atom
...

In fact this structure is misleading because it suggests that the two oxygen atoms on the right-hand
side of the diagram are joined to the nitrogen in different ways
...
There is no way of
showing this using a dots-and-crosses picture
...


METALLIC BONDING
What is a metallic bond?
Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms
...
8°C) melts at a considerably higher temperature than the element (neon) which precedes it in the
Periodic Table
...
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
...
And each of these eight is in turn being touched
by eight sodium atoms, which in turn are touched by eight atoms - and so on and so on, until you have taken in all the atoms
in that lump of sodium
...
There have to be huge numbers of molecular orbitals, of course, because any orbital can only hold two
electrons
...

The electrons are said to be delocalised
...


This is sometimes described as "an array of positive ions in a sea of electrons"
...

Each positive centre in the diagram represents all the rest of the atom apart from the outer electron, but that electron hasn't
been lost - it may no longer have an attachment to a particular atom, but it's still there in the structure
...

Metallic bonding in magnesium
If you work through the same argument with magnesium, you end up with stronger bonds and so a higher melting point
...
Both of these electrons become delocalised, so the "sea" has twice the
electron density as it does in sodium
...

More realistically, each magnesium atom has one more proton in the nucleus than a sodium atom has, and so not only will
there be a greater number of delocalised electrons, but there will also be a greater attraction for them
...

Each magnesium atom also has twelve near neighbours rather than sodium's eight
...

The metallic bond in molten metals

16

In a molten metal, the metallic bond is still present, although the ordered structure has been broken down
...
That means that boiling point is actually a better guide to the strength of the metallic
bond than melting point is
...


ATOMIC AND IONIC RADIUS
This page explains the various measures of atomic radius, and then looks at the way it varies around the Periodic Table across periods and down groups
...

ATOMIC RADIUS
Measures of atomic radius
Unlike a ball, an atom doesn't have a fixed radius
...


As you can see from the diagrams, the same atom could be found to have a different radius depending on what was around it
...
The atoms are pulled closely together and so the measured radius is less than if
they are just touching
...
The type of atomic radius being measured here is called the metallic radius or the covalent radius depending
on the bonding
...
The attractive forces are much less, and the atoms
are essentially "unsquashed"
...

Trends in atomic radius in the Periodic Table
The exact pattern you get depends on which measure of atomic radius you use - but the trends are still valid
...

Trends in atomic radius in Periods 2 and 3

17

Trends in atomic radius down a group
It is fairly obvious that the atoms get bigger as you go down groups
...

Trends in atomic radius across periods
You have to ignore the noble gas at the end of each period
...
All the other atoms are being measured where
their atomic radius is being lessened by strong attractions
...

Leaving the noble gases out, atoms get smaller as you go across a period
...
(Look back to the left-hand side of the first diagram on this page if you aren't sure, and picture the
bonding electrons as being half way between the two nuclei
...
The increasing number of
protons in the nucleus as you go across the period pulls the electrons in more tightly
...

In the period from sodium to chlorine, the same thing happens
...

IONIC RADIUS
Ions aren't the same size as the atoms they come from
...


Positive ions
Positive ions are smaller than the atoms they come from
...
You've lost a whole layer of electrons,
and the remaining 10 electrons are being pulled in by the full force of 11 protons
...
Chlorine is 2,8,7; Cl- is 2,8,8
...
There are still only 17 protons,
but they are now having to hold 18 electrons
...
Compounds like this consist of a giant (endlessly repeating) lattice of
ions
...

You should be clear that giant in this context doesn't just mean very large
...

There could be billions of sodium ions and chloride ions packed together, or trillions, or whatever - it simply depends how
big the crystal is
...

A small representative bit of a sodium chloride lattice looks like this:

If you look at the diagram carefully, you will see that the sodium ions and chloride ions alternate with each other in each of
the three dimensions
...
We normally
draw an "exploded" version which looks like this:

Only those ions joined by lines are actually touching each other
...
By chance we might just as well have centred
the diagram around a chloride ion - that, of course, would be touched by 6 sodium ions
...

Why is sodium chloride 6:6-co-ordinated?
The more attraction there is between the positive and negative ions, the more energy is released
...

That means that to gain maximum stability, you need the maximum number of attractions
...
If they start touching, you introduce repulsions into the crystal which makes it less stable
...
Ionic substances all have high melting and boiling points
...

The 2+ and 2- ions attract each other more strongly than 1+ attracts 1-
...
Rubidium iodide, for
example, melts and boils at slightly lower temperatures than sodium chloride, because both rubidium and iodide ions
are bigger than sodium and chloride ions
...

Sodium chloride crystals are brittle
Brittleness is again typical of ionic substances
...


Ions of the same charge are brought side-by-side and so the crystal repels itself to pieces!
Sodium chloride is soluble in water
Many ionic solids are soluble in water - although not all
...
Positive ions are attracted to the lone
pairs on water molecules and co-ordinate (dative covalent) bonds may form
...

Sodium chloride is insoluble in organic solvents
This is also typical of ionic solids
...

The electrical behaviour of sodium chloride
Solid sodium chloride doesn't conduct electricity, because there are no electrons which are free to move
...
In the process, sodium and chlorine are produced
...

The positive sodium ions move towards the negatively charged electrode (the cathode)
...

Meanwhile, chloride ions are attracted to the positive electrode (the anode)
...
These then pair up to make chlorine molecules
...


20

The new electrons deposited on the anode are pumped off around the external circuit by the power source, eventually ending
up on the cathode where they will be transferred to sodium ions
...
Both of these have to happen if you are to get electrons flowing in the external circuit
...


IONISATION ENERGY
Defining first ionisation energy
The first ionisation energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to
produce 1 mole of gaseous ions each with a charge of 1+
...

Things to notice about the equation
The state symbols - (g) - are essential
...

Ionisation energies are measured in kJ mol-1 (kilojoules per mole)
...

All elements have a first ionisation energy - even atoms which don't form positive ions in test tubes
...
E
...

Patterns of first ionisation energies in the Periodic Table The first
20 elements

First ionisation energy shows periodicity
...
For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar
...

Factors affecting the size of ionisation energy
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus
...

The size of that attraction will be governed by:
The charge on the nucleus
...

The distance of the electron from the nucleus
...
An electron close to the nucleus will be much more strongly attracted than one
further away
...

Consider a sodium atom, with the electronic structure 2,8,1
...
Between it and the nucleus there are the
two layers of electrons in the first and second levels
...
The outer electron therefore only feels a net pull of approximately 1+ from the centre
...

Whether the electron is on its own in an orbital or paired with another electron
...
This offsets the attraction of the nucleus, so
that paired electrons are removed rather more easily than you might expect
...
It is a very small atom, and the single electron is close to the nucleus and
therefore strongly attracted
...

Helium has a structure 1s2
...
It is close to the
nucleus and unscreened
...

Lithium is 1s22s1
...
You might argue that
that would be offset by the additional proton in the nucleus, but the electron doesn't feel the full pull of the nucleus - it is
screened by the 1s2 electrons
...

If you compare lithium with hydrogen (instead of with helium), the hydrogen's electron also feels a 1+ pull from the nucleus,
but the distance is much greater with lithium
...

The patterns in periods 2 and 3
Talking through the next 17 atoms one at a time would take ages
...

The first thing to realise is that the patterns in the two periods are identical - the difference being that the ionisation energies
in period 3 are all lower than those in period 2
...

In the whole of period 2, the outer electrons are in 2-level orbitals - 2s or 2p
...

The major difference is the increasing number of protons in the nucleus as you go from lithium to neon
...
In fact the increasing nuclear charge
also drags the outer electrons in closer to the nucleus
...


23

In period 3, the trend is exactly the same
...
They all have the same sort of environment, but there is an increasing nuclear charge
...
The outer electron is removed more easily from these atoms
than the general trend in their period would suggest
...
E
...
E
...
Offsetting that is the fact
that boron's outer electron is in a 2p orbital rather than a 2s
...
This has two effects
...

The 2p orbital is screened not only by the 1s2 electrons but, to some extent, by the 2s2 electrons as well
...


The explanation for the drop between magnesium and aluminium is the same, except that everything is happening at the
3level rather than the 2-level
...
E
...
E
...
Both of these factors offset the effect of the extra proton
...
What is offsetting it this time?
N

1s22s22px12py12pz1

1st I
...
= 1400 kJ mol-1

O

1s22s22px22py12pz1

1st I
...
= 1310 kJ mol-1

The screening is identical (from the 1s2 and, to some extent, from the 2s2 electrons), and the electron is being removed from
an identical orbital
...
The repulsion between the two
electrons in the same orbital means that the electron is easier to remove than it would otherwise be
...

Trends in ionisation energy down a group

24

As you go down a group in the Periodic Table ionisation energies generally fall
...


Taking Group 1 as a typical example:

Why is the sodium value less than that of lithium?
There are 11 protons in a sodium atom but only 3 in a lithium atom, so the nuclear charge is
much greater
...

Li
Na

1s22s1 1st I
...
= 519 kJ mol-1
1s22s22p63s1
1st I
...
= 494 kJ mol-1

Lithium's outer electron is in the second level, and only has the 1s2 electrons to screen it
...

The sodium's outer electron is in the third level, and is screened from the 11 protons in the nucleus by a total of 10 inner
electrons
...
In other words, the effect of the extra
protons is compensated for by the effect of the extra screening electrons
...
That lowers the ionisation energy
...

Ionisation energies and reactivity
The lower the ionisation energy, the more easily this change happens:

You can explain the increase in reactivity of the Group 1 metals (Li, Na, K, Rb, Cs) as you go down the group in terms of the
fall in ionisation energy
...

The danger with this approach is that the formation of the positive ion is only one stage in a multi-step process
...
The energy changes in
these processes also vary from
element to element
...


25

However, the ionisation energies of the elements are going to be major contributing factors towards the activation energy
of the reactions
...
The
lower the activation energy, the faster the reaction will be - irrespective of what the overall energy changes in the reaction
are
...


ELECTRON AFFINITY
First electron affinity
Ionisation energies are always concerned with the formation of positive ions
...

Defining first electron affinity
The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of
gaseous 1- ions
...


It is the energy released (per mole of X) when this change happens
...
For example, the first electron affinity of chlorine is -349 kJ mol-1
...


The first electron affinities of the group 7 elements
F
-328 kJ mol-1
Cl
-349 kJ mol-1
Br
-324 kJ mol-1
I
-295 kJ mol-1

Is there a pattern?
Yes - as you go down the group, first electron affinities become less (in the sense that less energy is evolved when the
negative ions are formed)
...

The electron affinity is a measure of the attraction between the incoming electron and the nucleus - the stronger the attraction,
the more energy is released
...


26

The increased nuclear charge as you go down the group is offset by extra screening electrons
...

For example, a fluorine atom has an electronic structure of 1s22s22px22py22pz1
...

The incoming electron enters the 2-level, and is screened from the nucleus by the two 1s2 electrons
...

By contrast, chlorine has the electronic structure 1s22s22p63s23px23py23pz1
...

But again the incoming electron feels a net attraction from the nucleus of 7+ (17 protons less the 10 screening electrons in the
first and second levels)
...
The greater the distance, the less the attraction and so the less energy is released as electron affinity
...

However, because fluorine is such a small atom, you are putting the new electron into a region of space already crowded with
electrons and there is a significant amount of repulsion
...

A similar reversal of the expected trend happens between oxygen and sulphur in Group 6
...

Comparing Group 6 and Group 7 values
As you might have noticed, the first electron affinity of oxygen (-142 kJ mol-1) is less than that of fluorine (-328 kJ mol-1)
...
Why?
It's simply that the Group 6 element has 1 less proton in the nucleus than its next door neighbour in Group 7
...

That means that the net pull from the nucleus is less in Group 6 than in Group 7, and so the electron affinities are less
...

Often in their reactions these elements form their negative ions
...
That explanation looks reasonable until you include fluorine!
An overall reaction will be made up of lots of different steps all involving energy changes, and you cannot safely try to
explain a trend in terms of just one of those steps
...
Second electron affinity
You are only ever likely to meet this with respect to the group 6 elements oxygen and sulphur which both form 2- ions
...

This is more easily seen in symbol terms
...

Why is energy needed to do this?
You are forcing an electron into an already negative ion
...
The second electron affinity of oxygen is
particularly high because the electron is being forced into a small, very electron-dense space
...

The Pauling scale is the most commonly used
...
0, and
values range down to caesium and francium which are the least electronegative at 0
...


What happens if two atoms of equal electronegativity bond together?
Consider a bond between two atoms, A and B
...


28

If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be
found on average half way between the two atoms
...

You will find this sort of bond in, for example, H2 or Cl2 molecules
...
The electrons are actually in a molecular orbital, and are moving
around all the time within that orbital
...

What happens if B is slightly more electronegative than A?
B will attract the electron pair rather more than A does
...
At the
same time, the A end (rather short of electrons) becomes slightly positive
...

Defining polar bonds
This is described as a polar bond
...
Examples include most
covalent bonds
...

What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged right over to B's end of the bond
...
Ions have been formed
...
In a pure covalent bond, the
electrons are held on average exactly half way between the atoms
...

How far does this dragging have to go before the bond counts as ionic? There is no real answer to that
...

Because of the properties of sodium chloride, however, we tend to count it as if it were purely ionic
...
In this case, the pair of
electrons hasn't moved entirely over to the iodine end of the bond
...

Summary
• No electronegativity difference between two atoms leads to a pure non-polar covalent bond
...


A large electronegativity

difference leads to an ionic bond
...
What about more complicated molecules?
In CCl4, each bond is polar
...
The whole of the outside of the
molecule is somewhat negative, but there is no overall separation of charge from top to bottom, or
from left to right
...


The hydrogen at the top of the molecule is less electronegative than carbon and so is slightly positive
...

A polar molecule will need to be "lop-sided" in some way
...
If you remember that fact, everything becomes easy, because electronegativity
must always increase towards fluorine in the Periodic Table
...
Historically this is because they were believed not to form bonds - and if
they don't form bonds, they can't have an electronegativity value
...


30

Trends in electronegativity across a period
As you go across a period the electronegativity increases
...
It doesn't have an electronegativity, because it doesn't form bonds
...
(If it increases up to fluorine, it must decrease as you go down
...


Explaining the patterns in electronegativity
The attraction that a bonding pair of electrons feels for a particular nucleus depends on:
•
•
•

the number of protons in the nucleus;
the distance from the nucleus;
the amount of screening by inner electrons
...
Think of sodium
chloride as if it were covalently bonded
...
The electron pair is screened from both nuclei by the
1s, 2s and 2p electrons, but the chlorine nucleus has 6 more protons in it
...

Electronegativity increases across a period because the number of charges on the nucleus increases
...

Why does electronegativity fall as you go down a group?
Think of hydrogen fluoride and hydrogen chloride
...
In the chlorine case it is shielded by all the
1s22s22p6 electrons
...
But fluorine has the bonding pair in the 2level
rather than the 3-level as it is in chlorine
...

As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the
attraction of the nucleus
...
You can also think of it the
other way round
...
Imagine instead that it was ionic
...

The aluminium ion is very small and is packed with three positive charges - the "charge density" is therefore very high
...


32

In the case of aluminium chloride, the electron pairs are dragged back towards the aluminium to such an extent that the bonds
becomes covalent
...
The polarising ability depends on
the charge density in the positive ion
...

As a negative ion gets bigger, it becomes easier to polarise
...

A positive ion would be more effective in attracting a pair of electrons from an iodide ion than the corresponding electrons in,
say, a fluoride ion where they are much closer to the nucleus
...
On the other hand,
aluminium fluoride is ionic because the aluminium ion can't polarise the small fluoride ion sufficiently to form a covalent
bond
...


33

Explaining the increase in atomic radius
The radius of an atom is governed by
the number of layers of electrons around the nucleus
electrons feel from the nucleus
...
The positive charge on the nucleus is cut down by
the negativeness of the inner electrons
...
Obviously, the more layers of electrons you have, the more space they will take up - electrons
repel each other
...

Trends in First Ionisation Energy
First ionisation energy is the energy needed to remove the most loosely held electron from each of one mole of gaseous atoms
to make one mole of singly charged gaseous ions - in other words, for 1 mole of this process:

Notice that first ionisation energy falls as you go down the group
...


As you go down the Group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons
...

However, as you go down the Group, the distance between the nucleus and the outer electrons increases and so they become
easier to remove - the ionisation energy falls
...

Defining second ionisation energy
Second ionisation energy is defined by the equation:

It is the energy needed to remove a second electron from each ion in 1 mole of gaseous 1+ ions to give gaseous 2+ ions
...

The first four ionisation energies of aluminium, for example, are given by
1st I
...
= 577 kJ mol-1
2nd I
...
= 1820 kJ mol-1
3rd I
...
= 2740 kJ mol-1
4th I
...
= 11600 kJ mol-1
In order to form an Al3+(g) ion from Al(g) you would have to supply:
577 + 1820 + 2740 = 5137 kJ mol-1
That's a lot of energy
...

For example, if aluminium reacts with fluorine or oxygen, it can recover that energy in various changes involving the fluorine
or oxygen - and so aluminium fluoride or aluminium oxide contain Al3+ ions
...


35

Why doesn't aluminium form an Al4+ ion? The fourth ionisation energy is huge compared with the first three, and there is
nothing that aluminium can react with which would enable it to recover that amount of extra energy
...
Trying to remove a negative electron from a
positive ion is going to be more difficult than removing it from an atom
...

Why is the fourth ionisation energy of aluminium so large?
The electronic structure of aluminium is 1s22s22p63s23px1
...
Once they've gone, the fourth electron is removed from the 2p level - much closer to the nucleus, and only
screened by the 1s2 (and to some extent the 2s2) electrons
...
You can use
this to work out which group of the Periodic Table an element is in from its successive ionisation energies
...
It means that there are 2 electrons which are relatively easy to
remove (the 3s2 electrons), while the third one is much more difficult (because it comes from an inner level - closer to the
nucleus and with less screening)
...
The first 4 electrons are coming from the 3-level
orbitals; the fifth from the 2-level
...
That is the same as the group number
...
Then there is a huge jump of about
15000
...

Exploring the patterns in more detail
If you plot graphs of successive ionisation energies for a particular element, you can see the fluctuations in it caused by the
different electrons being removed
...

Chlorine has the electronic structure 1s22s22p63s23px23py23pz1
...
The green labels show which electron is being removed for
each of the ionisation energies
...
That is because the first two electrons are coming from pairs in the 3p levels and are therefore
rather easier to remove than if they were unpaired
...
That is because the 6th and 7th electrons are coming
from the 3s level - slightly closer to the nucleus and slightly less well screened
...
It is usually measured on the
Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4
...


All of these elements have a very low electronegativity
...
0
...
The atoms become less and less good
at attracting bonding pairs of electrons
...
Think of it to start with as a covalent bond - a pair of shared
electrons
...


The electron pair ends up so close to the chlorine that there is essentially a transfer of an electron to the chlorine - ions are
formed
...
Now compare this with the
lithium-chlorine bond
...
That means
that the electron pair is going to be closer to
the net 1+ charge from the lithium end, and
so more strongly attracted to it
...
Lithium
iodide, for example, will dissolve in organic solvents - a typical property of covalent compounds
...

Summarising the trend down the Group
As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly
attracted towards it
...

With the exception of some lithium compounds, these elements all form compounds which we consider as being fully ionic
...


TRENDS IN MELTING AND BOILING POINTS

You will see that both the melting points and boiling points fall as you go down the Group
...

The fall in melting and boiling points reflects the fall in the strength of the metallic bond
...
As the atoms get bigger, the
nuclei get further away from these delocalised electrons, and so the attractions fall
...

In the same way that we have already discussed, each of these atoms has a net pull from the nuclei of 1+
...
All that matters is the
distance between the nucleus and the bonding electrons
...
Almost everybody sees and describes colours differently, For example, the word
"red" has been used several times to describe colours which can be quite different from each other
...
Reactivity increases
as you go down the Group
...

Reaction with oxygen is just a more dramatic version of the reaction with air
...

Details for the individual metals Lithium
Lithium burns with a strongly red-tinged flame if heated in air
...

With pure oxygen, the flame would simply be more intense
...
Lithium is the only element in this Group to
form a nitride in this way
...
Using larger amounts of sodium or burning it
in oxygen gives a strong orange flame
...

The equation for the formation of the simple oxide is just like the lithium one
...
Larger pieces of potassium burn with a lilac flame
...
and for the superoxide:

Rubidium and caesium
Both metals catch fire in air and produce superoxides, RbO2 and CsO2
...

Different oxides formed as you go down the Group?
•
•
•

Lithium (and to some extent sodium) form simple oxides, X2O, which contain the common O2- ion
...

Potassium, rubidium and caesium form superoxides, XO2
...

BUT
...

At the top of the Group, the small ions with a higher charge density tend to polarise the more complicated oxide ions to the
point of destruction
...

For example, lithium oxide reacts with water to give a colourless solution of lithium hydroxide
...
For example, sodium oxide will react with dilute
hydrochloric acid to give colourless sodium chloride solution and water
...


If the temperature increases (as it inevitably will unless the peroxide is added to water very, very, very slowly!), the hydrogen
peroxide produced decomposes into water and oxygen
...

Reaction with dilute acids
These reactions are even more exothermic than the ones with water
...
The hydrogen peroxide will decompose to give water and oxygen if the temperature rises - again, it is almost
impossible to avoid this
...
Once
again, these are strongly exothermic reactions and the heat produced will inevitably decompose the hydrogen peroxide to
water and more oxygen
...
A solution containing a salt and hydrogen
peroxide is formed together with oxygen gas
...
Violent!

42

THE REACTIONS OF THE ELEMENTS WITH CHLORINE
This is included on this page because of the similarity in appearance between the reactions of the Group 1 metals with
chlorine and with oxygen
...
The
rest also behave the same in both gases
...
There is nothing in any way complicated about
these reactions!

REACTIONS OF THE GROUP 1 ELEMENTS WITH WATER
General
All of these metals react vigorously or even explosively with cold water
...


This equation applies to any of these metals and water - just replace the X by the symbol you want
...

Details for the individual metals
Lithium
Lithium's density is only about half that of water so it floats on the surface, gently fizzing and giving off hydrogen
...
The reaction generates heat too slowly
and lithium's melting point is too high for it to melt (see sodium below)
...
A white trail of sodium hydroxide is seen in the water under the sodium, but this soon dissolves to give a colourless
solution of sodium hydroxide
...
If the sodium becomes
trapped on the side of the container, the hydrogen may catch fire to burn with an orange flame
...


43

Potassium
Potassium behaves rather like sodium except that the reaction is faster and enough heat is given off to set light to the
hydrogen
...

Rubidium
Rubidium is denser than water and so sinks
...
Rubidium hydroxide solution and hydrogen are formed
...
Caesium hydroxide and hydrogen are
formed
Summary of the trend in reactivity
The Group 1 metals become more reactive towards water as you go down the group
...
This is in part due to a decrease in ionisation
energy as you go down the Group, and in part to a fall in atomisation energy reflecting weaker metallic bonds as you go from
lithium to caesium
...


SOME COMPOUNDS OF THE GROUP 1 ELEMENTS
The effect of heat on Group 1 compounds
Group 1 compounds are more stable to heat than the corresponding compounds in Group 2
...

Heating the nitrates
Most nitrates tend to decompose on heating to give the metal oxide, brown fumes of nitrogen dioxide, and oxygen
...


The rest of the Group, however, don't decompose so completely (at least not at Bunsen temperatures) - producing the metal
nitrite and oxygen, but no nitrogen dioxide
...
As you go down the Group, the decomposition gets more difficult, and you have to use higher
temperatures
...

For example, a typical Group 2 carbonate like calcium carbonate decomposes like this:

In Group 1, lithium carbonate behaves in the same way - producing lithium oxide and carbon dioxide
...
The
decomposition temperatures again increase as you go down the Group
...
Any
attempt to get them out of solution causes them to decompose to give the carbonate, carbon dioxide and water
...
For example, for sodium hydrogencarbonate:

Explanations for the trends in thermal stability
Detailed explanations are given for the carbonates because the diagrams are easier to draw
...

There are two ways of explaining the increase in thermal stability as you go down the Group
...

Explaining the trend in terms of the polarising ability of the positive ion
A small positive ion has a lot of charge packed into a small volume of space - especially if it has more than one positive
charge
...

A bigger positive ion has the same charge spread over a larger volume of space
...

The structure of the carbonate ion

45

If you worked out the structure of a carbonate ion using "dots-and-crosses" or some similar method, you would probably
come up with:

This shows two single carbon-oxygen bonds and one double one, with two of the oxygens each carrying a negative charge
...
We say that the charges are delocalised
...
The shading is intended to show that there is a
greater chance of finding them around the oxygen atoms than near the carbon
...
The positive ion attracts the delocalised electrons in
the carbonate ion towards itself
...
The diagram shows what happens with an ion from
Group 2, carrying two positive charges

If this is heated, the carbon dioxide breaks free to leave the metal oxide
...
If it is highly polarised,
you need less heat than if it is only slightly polarised
...
That is why the Group 1 compounds are
more thermally stable than those in Group 2
...


46

The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion
...
To compensate for
that, you have to heat the compound more in order to persuade the carbon dioxide to break free and leave the metal oxide
...

What about the nitrates and hydrogencarbonates?
The argument is exactly the same here
...

And, again, the Group 1 compounds will need to be heated more strongly than those in Group 2 because the Group 1 ions are
less polarising
...
Magnesium carbonate is soluble to the extent of about 0
...

By contrast, the least soluble Group 1 carbonate is lithium carbonate
...
3 g per 100 g of water at 20°C
...
5 g per 100 g of water at this temperature for caesium carbonate
...

The hydroxides
The least soluble hydroxide in Group 1 is lithium hydroxide - but it is still possible to make a solution with a concentration of
12
...
The other hydroxides in the Group are even more soluble
...

In Group 2, the most soluble one is barium hydroxide - and it is only possible to make a solution of concentration around 3
...

The Group 1 hydrides
Saline (salt-like) hydrides
The hydrides of Group 1 metals are white crystalline solids which contain the metal ions and hydride ions, H-
...

Because they can react violently with water or moist air, they are normally supplied as suspensions in mineral oil
...
For example, for lithium hydride:

47

Reactions of the Group 1 hydrides
Electrolysis
On heating, most of these hydrides decompose back into the metal and hydrogen before they melt
...

The metal is released at the cathode as you would expect
...

The anode equation is:

The other Group 1 hydrides can be electrolysed in solution in various molten mixtures such as a mixture of lithium chloride
and potassium chloride
...

Reaction with water
These hydrides react violently with water releasing hydrogen gas and producing the metal hydroxide
...


ATOMIC AND PHYSICAL PROPERTIES OF THE GROUP 2 ELEMENTS

Trends in Atomic Radius
You can see that the atomic radius increases as you go down the Group
...

Explaining the increase in atomic radius
The radius of an atom is governed by
the number of layers of electrons around the nucleus
the pull the outer electrons feel from the nucleus
...
The
positive charge on the nucleus is cut down by the negativeness of the inner
electrons
...
Work it out for calcium if
you aren't convinced
...
Obviously, the more layers of electrons you have, the more space they will take up - electrons
repel each other
...

Trends in First Ionisation Energy
First ionisation energy is the energy needed to remove the most loosely
held electron from each of one mole of gaseous atoms to make one mole
of singly charged gaseous ions - in other words, for 1 mole of this
process:

Notice that first ionisation energy falls as you go down the group
...


As you go down the Group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons
...


49

However, as you go down the Group, the distance between the nucleus and the outer electrons increases and so they become
easier to remove - the ionisation energy falls
...
It is usually measured on the Pauling scale, on which the most
electronegative element (fluorine) is given an electronegativity of 4
...

All of these elements have a low electronegativity
...
0
...
The atoms become less and
less good at attracting bonding pairs of electrons
...
Think of it to start with as a covalent bond - a pair of shared
electrons
...


The electron pair ends up so close to the chlorine that there is essentially a transfer of an electron to the chlorine - ions are
formed
...

Now compare this with the beryllium-chlorine bond
...
That
means that the electron pair is going to be closer to the net 2+ charge from the
beryllium end, and so more strongly attracted to it
...
Because of its small size, beryllium forms covalent bonds, not ionic ones
...

Summarising the trend down the Group
As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly
attracted towards it
...

As you go down the Group, the bonds formed between these elements and other things such as chlorine become more and
more ionic
...


50

REACTIONS OF THE GROUP 2 ELEMENTS WITH AIR OR OXYGEN
The reactions with oxygen
Formation of simple oxides
On the whole, the metals burn in oxygen to form a simple metal oxide
...
Beryllium has a very strong (but very thin) layer of
beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it
...

It is almost impossible to find any trend in the way the metals react with oxygen
...

To be able to make any sensible comparison, you would have to have pieces of metal which were all equally free of oxide
coating, with exactly the same surface area and shape, exactly the same flow of oxygen around them, and heated to exactly
the same extent to get them started
...

Calcium is quite reluctant to start burning, but then bursts dramatically into flame, burning with an intense white
flame with a tinge of red at the end
...

Barium: The flame appeares to be white with some pale green tinges
...

Strontium forms this if it is heated in oxygen under high pressures, but barium forms barium peroxide just on normal heating
in oxygen
...


The strontium equation would look just the same
...
In each case, you will get a mixture of the metal oxide and the metal nitride
...

Why do some metals form peroxides on heating in oxygen?
Beryllium, magnesium and calcium don't form peroxides when heated in oxygen, but strontium and barium do
...

The peroxide ion, O22- looks like this:

The covalent bond between the two oxygen atoms is relatively weak
...
Electrons in the peroxide ion will be strongly attracted
towards the positive ion
...


We say that the positive ion polarises the negative ion
...

Ions of the metals at the top of the Group have such a high charge
density (because they are so small) that any peroxide ion near them falls to pieces to give an oxide and oxygen
...

Barium peroxide can form because the barium ion is so large that it doesn't have such a devastating effect on the peroxide
ions as the metals further up the Group
...

Nitrogen is fairly unreactive because of the very large amount of energy needed to break the triple bond joining the two atoms
in the nitrogen molecule, N2
...
All of these processes absorb energy
...

Energy is evolved when the ions come together to produce the crystal lattice
...

The size of the lattice energy depends on the attractions between the ions
...
In the whole of Group 2, the attractions
between the 2+ metal ions and the 3- nitride ions are big enough to produce very high lattice energies
...
The excess energy evolved makes the overall process exothermic
...

Their ions only carry one positive charge, and so the lattice energies of their nitrides will be much less
...
Lithium has by far the smallest ion in the Group, and so lithium nitride
has the largest lattice energy of any possible Group 1 nitride
...

In all the other cases in Group 1, the overall reaction would be endothermic
...


REACTIONS OF THE GROUP 2 ELEMENTS WITH WATER
Beryllium
Beryllium has no reaction with water or steam even at red heat
...


Very clean magnesium has a very slight reaction with cold water
...


Calcium, strontium and barium
These all react with cold water with increasing vigour to give the metal hydroxide and hydrogen
...

The equation for the reactions of any of these metals would be:

53

The hydroxides aren't very soluble, but they get more soluble as you go down the Group
...
You get less precipitate as you go down the Group
because more of the hydroxide dissolves in the water
...

Explaining the trend in reactivity
Beryllium as a special case
There is an additional reason for the lack of reactivity of beryllium compared with the rest of the Group
...
(This is just like the aluminium case that you are
probably familiar with
...

Looking at the enthalpy changes for the reactions
The enthalpy change of a reaction is a measure of the amount of heat absorbed or evolved when the reaction takes place
...
That's really all you need to know for this
section!
If you calculate the enthalpy change for the possible reactions between beryllium or magnesium and steam, you come up with
these answers:

Notice that both possible reactions are strongly exothermic, giving out almost identical amounts of heat
...
The explanation for the different reactivities must lie somewhere else
...

Looking at the activation energies for the reactions
The activation energy for a reaction is the minimum amount of energy which is needed in order for the reaction to take place
...

When Group 2 metals react to form oxides or hydroxides, metal ions are formed
...
These stages involve the input of:
•

the atomisation energy of the metal
...


54

•

the first + second ionisation energies
...


After this, there will be a number of steps which give out heat again - leading to the formation of the products, and overall
exothermic reactions
...


Notice that the ionisation energies dominate this - particularly the second ionisation energies
...
Because it gets easier to form the ions, the reactions will happen more quickly
...
This is mainly due to a decrease in ionisation
energy as you go down the Group
...


EXPLANATIONS FOR THE TRENDS IN SOLUBILITY OF SOME GROUP 2 COMPOUNDS
The usual explanations
Enthalpy changes during the process
The usual explanation is in terms of the enthalpy changes which occur when an ionic compound dissolves in water
...

As you go down a Group, the energy needed to break up the lattice falls as the positive ions get bigger
...

Again as the positive ions get bigger, the energy released as the ions bond to water molecules (their hydration enthalpies) falls
as well
...


55

Since both of these important enthalpy terms fall as you go down the Group, what matters in deciding whether the change
becomes more endothermic or more exothermic overall is how fast they fall relative to each other
...

In this case, we are defining lattice enthalpy as the heat needed to convert 1 mole of crystal in its standard state into separate
gaseous ions - an endothermic change
...

•

•

If the hydration enthalpy falls faster than the lattice enthalpy (as in this case), the net effect is that the overall change
becomes more endothermic (or less exothermic in other possible cases where the total enthalpy change turns out to be
negative)
...


What controls the relative rate of fall of the two terms?
It turns out that the main factor is the size of the negative ion
...
That means that the enthalpy of solution will become less positive (or more negative)
...
In this case, the enthalpy of solution will become more positive (or less negative)
...
Where you have a big negative ion, this inter-ionic distance is largely controlled by the size of that negative ion
...

With sulphates, for example, the percentage increase in the inter-ionic distance as you go from magnesium to calcium
sulphate isn't as great as it would be with a smaller negative ion like hydroxide
...

The relationship between enthalpy of solution and solubility
The assumption is made that the more endothermic (or less exothermic) the enthalpy of solution is, the less soluble the
compound
...


SOLUBILITY OF THE HYDROXIDES, SULPHATES AND CARBONATES OF THE GROUP 2
ELEMENTS IN WATER
Solubility of the hydroxides
•

The hydroxides become more soluble as you go down the Group
...

Some examples may help you to remember the trend:
Magnesium hydroxide appears to be insoluble in water
...
This shows that there are more hydroxide ions in the solution than there were in
the original water
...

Calcium hydroxide solution is used as "lime water"
...

Barium hydroxide is soluble enough to be able to produce a solution with a concentration of around 0
...


Solubility of the sulphates
•

The sulphates become less soluble as you go down the Group
...

The Nuffield Data Book quotes anyhydrous beryllium sulphate, BeSO4, as insoluble , whereas the hydrated form,
BeSO4
...
(The Data Books agree on this - giving a figure of about 39 g dissolving in 100 g of water at room
temperature
...

Two common examples may help you to remember the trend:

57

You are probably familiar with the reaction between magnesium and dilute sulphuric acid to give lots of hydrogen and a
colourless solution of magnesium sulphate
...
The magnesium sulphate is
obviously soluble
...

The ready formation of a precipitate shows that the barium sulphate must be pretty insoluble
...

Solubility of halides
The halides of the alkaline earth metals, MX2, are easily isolated and the anhydrous salts can be obtained by heating the
hydrated salts and are essentially ionic
...
The solubility of
the halide decreases on descending the group because the hydration enthalpies decrease faster than the lattice enthalpies do
...
This is because the very small size of the
fluoride ions means that in the solid state the much larger M2+ ions are more and more in contact with each other as their size
increases causing a faster lowering of the lattice energies

Solubility of the carbonates
The carbonates tend to become less soluble as you go down the Group
...
Magnesium carbonate is soluble to the extent of about
0
...

The trend to lower solubility is, however, broken at the bottom of the Group
...

There are no simple examples which might help you to remember the carbonate trend
...
Thermal
decomposition is the term given to splitting up a compound by heating it
...

•

The carbonates become more stable to heat as you go down the Group
...

Again, if "X" represents any one of the elements:

As you go down the Group, the nitrates also have to be heated more strongly before they will decompose
...


Summary
Both carbonates and nitrates become more thermally stable as you go down the Group
...

Explaining the trend in terms of the polarising ability of the positive ion
A small 2+ ion has a lot of charge packed into a small volume of space
...

A bigger 2+ ion has the same charge spread over a larger volume of space
...


59

The structure of the carbonate ion
If you worked out the structure of a carbonate ion using "dots-and-crosses" or some similar method, you
would probably come up with:

This shows two single carbon-oxygen bonds and one double one, with two of the oxygens each carrying a negative charge
...
We say that the charges are delocalised
...

For the purposes of this topic, you don't need to understand how this bonding has come about
...
The shading is intended to show that there is a
greater chance of finding them around the oxygen atoms than near the carbon
...
The positive ion attracts the delocalised electrons in
the carbonate ion towards itself
...


If this is heated, the carbon dioxide breaks free to leave the metal oxide
...
if it is highly polarised,
you need less heat than if it is only slightly polarised
...
As the
positive ions get bigger as you go down the Group, they have less effect on the carbonate ions near them
...

In other words, as you go down the Group, the carbonates become more thermally stable
...
The small positive ions at the top of the Group polarise the nitrate ions more than the
larger positive ions at the bottom
...


60

Explaining the trend in terms of the energetics of the process
Looking at the enthalpy changes
If you calculate the enthalpy changes for the decompostion of the various carbonates, you find that all the changes are quite
strongly endothermic
...

The calculated enthalpy changes (in kJ mol-1) are given in the table
...
Remember that the reaction we are talking about is:

MgCO3
CaCO3
SrCO3

+117
+178
+235

+267
BaCO3
You can see that the reactions become more endothermic as you go down the Group
...
You have to supply increasing amounts of heat energy to make them
decompose
...
Don't waste your time looking at it
...

Lattice enthalpy is the heat needed to split one mole of crystal in its standard state into its separate gaseous ions
...


For reasons we will look at shortly, the lattice enthalpies of both the oxides and carbonates fall as you go down the Group
...

The oxide lattice enthalpy falls faster than the carbonate one
...


62

Explaining the relative falls in lattice enthalpy
The size of the lattice enthalpy is governed by several factors, one of which is the distance between the centres of the positive
and negative ions in the lattice
...
If the attractions
are large, then a lot of energy will have to be used to separate the ions - the lattice enthalpy will be large
...
The inter-ionic distances are increasing and so the attractions become weaker
...
065; Ca2+

0
...
140; CO32-

?

The lattice enthalpies fall at different rates because of the different sizes of the two negative ions - oxide and carbonate
...
140 nm), whereas the carbonate ion is large (no figure available)
...
205 nm (0
...
065) to 0
...
140 + 0
...

In the carbonates, the inter-ionic distance is dominated by the much larger carbonate ion
...

Some made-up figures show this clearly
...
3 nm
...
365 nm to 0
...

The rates at which the two lattice energies fall as you go down the Group depends on the percentage change as you go from
one compound to the next
...

What about the nitrates?
The nitrate ion is bigger than an oxide ion, and so its radius tends to dominate the inter-ionic distance
...
if you constructed a cycle like that further up the page, the same arguments
would apply
...


Magnesium
Magnesium has a very slight reaction with cold water, but burns in steam
...
Magnesium hydroxide is formed as a very thin layer on the magnesium and this tends to stop the reaction
...


Aluminium
Aluminium powder heated in steam produces hydrogen and aluminium oxide
...


Silicon
There is a fair amount of disagreement in the books and on the web about what silicon does with water or steam
...

The common shiny grey lumps of silicon with a rather metal-like appearance are fairly unreactive
...


But it is also possible to make much more reactive forms of silicon which will react with cold water to give the same
products
...
Cotton and Wilkinson's Advanced Inorganic Chemistry (third
edition - page 316) suggests that the reactivity of one of these could be due to a very high surface area, or perhaps because the
silicon exists in a graphite-like structure
...

Chlorine

64

Chlorine dissolves in water to some extent to give a green solution
...


In the presence of sunlight, the chloric(I) acid slowly decomposes to produce more hydrochloric acid, releasing oxygen gas,
and you may come across an equation showing the overall change:

Argon
There is no reaction between argon and water
...

For the simple oxide:

For the peroxide:

Magnesium
Magnesium burns in oxygen with an intense white flame to give white solid magnesium oxide
...
If you sprinkle aluminium powder into a Bunsen flame, you get white sparkles
...


Silicon
Silicon will burn in oxygen if heated strongly enough
...


Phosphorus
White phosphorus catches fire spontaneously in air, burning with a white flame and producing clouds of white smoke - a
mixture of phosphorus(III) oxide and phosphorus(V) oxide
...
In an excess of oxygen, the product will be almost
entirely phosphorus(V) oxide
...
It produces colourless sulphur dioxide gas
...
Argon doesn't react either
...
White solid sodium chloride is produced
...


Aluminium
Aluminium is often reacted with chlorine by passing dry chlorine over aluminium foil heated in a long tube
...
This sublimes (turns straight from solid to
vapour and back again) and collects further down the tube where it is cooler
...
This is a colourless liquid
which vaporises and can be condensed further along the apparatus
...

Phosphorus(III) chloride is a colourless fuming liquid
...


Sulphur
If a stream of chlorine is passed over some heated sulphur, it reacts to form an orange, evil-smelling liquid, disulphur
dichloride, S2Cl2
...


SOME BERYLLIUM CHEMISTRY UNTYPICAL OF GROUP 2
Beryllium chloride is covalent
Physical properties
Beryllium chloride, BeCl2, melts at 405°C and boils at 520°C
...

Notice how much dramatically lower the boiling point of beryllium chloride is compared with magnesium chloride
...

Because its boiling point is much lower, it follows that beryllium chloride can't contain ions - it must be covalent
...
There must be something more complicated
going on!
Reaction with water
Beryllium chloride reacts vigorously and exothermically with water with the evolution of acidic, steamy hydrogen chloride
gas
...

In the first instance, it reacts to give hydrated beryllium ions, [Be(H2O)4]2+, and chloride ions
...
The small beryllium ion at the
centre attracts the electrons in the bonds towards itself, and that makes the hydrogen atoms in the water even more positive

67

than they usually are
...

All the other ionic chlorides in Group 2 dissolve in water without any obvious reaction
...

Beryllium chloride, BeCl2, is a linear molecule with all three atoms in a straight line
...

As a solid
...
It is a
very small molecule, and so the intermolecular attractions would be expected to be fairly weak
...
They do this by
forming coordinate bonds (dative covalent bonds) between lone pairs on chlorine atoms
and adjacent beryllium atoms
...


You can see from the last diagram that the beryllium atoms are still electron deficient
...
The next diagram shows the coordinate bonds in the conventional
way using arrows
...


2

Why isn't beryllium chloride ionic?
Beryllium has quite a high electronegativity compared with the rest of the Group
...

In order for an ionic bond to form, the beryllium has to let go of its electrons
...

Beryllium forms 4-coordinated complex ions
Some simple background
Although beryllium doesn't normally form simple ions, Be2+, it does form ions in solution
...


68

The ion is said to be 4-coordinated, or to have a coordination number of 4, because there are four water molecules arranged
around the central beryllium
...
For example, magnesium ions in solution exist as [Mg(H2O)6]2+
...
One of
the lone pairs on each water molecule is used to form a bond with an empty orbital in the metal ion
...
It would seem logical for the
metal ion to form as many bonds like this as it possibly can
...
Beryllium has the electronic structure 1s22s2
...
That leaves
the 2-level completely empty
...
In the next diagram the 1s
electrons have been left out
...


Each water molecule, of course, has two lone pairs of electrons
...

Notice that once four water molecules have bonded in this way, there isn't any more space available at the bonding level
...

The water molecules arrange themselves to get as far apart as possible - which is pointing towards the corners of a
tetrahedron
...

The hydration of magnesium
You might think that magnesium would behave just the same, but at the 3-level there are 3d orbitals available as well as 3s
and 3p
...
When that ion is hydrated, it uses the 3s orbital,
all three of the 3p orbitals and two of the 3d orbitals
...


69

Why does magnesium stop at attaching six waters? Why doesn't it use the remaining 3d orbitals as well? You can't physically
fit more than six water molecules around the magnesium - they take up too much room
...
The ions
become so big that they aren't sufficiently attractive to the lone pairs on the water molecules to form formal bonds - instead
the water molecules tend to cluster more loosely around the positive ions
...

Beryllium hydroxide is amphoteric
Amphoteric means that it can react with both acids and bases to form salts
...
They react with acids to form salts
...

Beryllium hydroxide
Beryllium hydroxide reacts with acids, forming solutions of beryllium salts
...
Beryllium hydroxide reacts with the sodium hydroxide to
give a colourless solution of sodium tetrahydroxoberyllate
...
The name describes this ion
...
The "ate" ending always shows that the ion is negative
...
It would probably have been made by
adding sodium hydroxide solution to a solution of a beryllium salt like beryllium sulphate
...
The beryllium has such a strongly
polarising effect on the water molecules that hydrogen ions are very easily removed from them
...
If you add just the right amount of sodium
hydroxide solution, you get a precipitate of what is normally called "beryllium hydroxide" - but which is a shade more
complicated than that!

The product (other than water) is a neutral complex, and it is covalently bonded
...

You get a precipitate of the neutral complex because of the lack of charge on it
...

What happens if you add an acid to this?
The hydrogen ions that were originally removed are simply replaced
...


What happens if you add a base?
Adding more hydroxide ions to the neutral complex pulls more hydrogen ions off the water molecules to give the
tetrahydroxoberyllate ion:

The beryllium hydroxide dissolves because the neutral complex is converted into an ion which will be sufficiently attracted to
water molecules
...
These react with hydrogen ions from an acid to
form water - and so the hydroxide reacts with acids
...
Adding more hydroxide ions from a base has no effect because
they haven't got anything to react with
...
Just as a reminder, the shortened versions of
the electronic structures for the eight elements are:
Na
[Ne] 3s1
Mg
[Ne] 3s2
Al
[Ne] 3s2 3px1
Si
[Ne] 3s2 3px1 3py1
P
[Ne] 3s2 3px1 3py1 3pz1
S
[Ne] 3s2 3px2 3py1 3pz1
Cl
[Ne] 3s2 3px2 3py2 3pz1
Ar
[Ne] 3s2 3px2 3py2 3pz2
In each case, [Ne] represents the complete electronic structure of a neon atom
...


The figures used to construct this diagram are based on:
•
•

metallic radii for Na, Mg and Al;
covalent radii for Si, P, S and Cl;
the van der Waals radius for Ar because it doesn't form any strong bonds
...
It isn't
fair to compare these with a van der Waals radius, though
...
You aren't comparing like with like
...

Explaining the trend
A metallic or covalent radius is going to be a measure of the distance from the nucleus to the bonding pair of electrons
...

From sodium to chlorine, the bonding electrons are all in the 3-level, being screened by the electrons in the first and second
levels
...
The amount of screening is constant for all of these elements
...


72

The Pauling scale is the most commonly used
...
0, and
values range down to caesium and francium which are the least
electronegative at 0
...

The trend
The trend across Period 3 looks like this:

Notice that argon isn't included
...
Since argon doesn't form covalent bonds, you obviously can't assign it an
electronegativity
...

As you go across the period, the bonding electrons are always in the same level - the 3-level
...

All that differs is the number of protons in the nucleus
...

Physical Properties
This section is going to look at the electrical conductivity and the melting and boiling points of the elements
...

Structures of the elements
The structures of the elements change as you go across the period
...

Three metallic structures
Sodium, magnesium and aluminium all have metallic structures
...
In magnesium, both of its outer
electrons are involved, and in aluminium all three
...

Sodium is 8-co-ordinated - each sodium atom is touched by only 8 other atoms
...
This is a more efficient way to
pack atoms, leading to less wasted space in the metal structures and to stronger bonding in the metal
...
A tiny part of the structure looks like this:

The structure is held together by strong covalent bonds in all three dimensions
...
For
phosphorus, I am assuming the common white phosphorus
...


The atoms in each of these molecules are held together by covalent bonds (apart, of course, from argon)
...

Electrical conductivity
•

Sodium, magnesium and aluminium are all good conductors of electricity
...

None of the rest conduct electricity
...

In the silicon case, explaining how semiconductors conduct electricity is beyond the scope of A level chemistry courses
...
There are no electrons free to move around
...
The figures are
plotted in kelvin rather than °C to avoid having negative values
...

The metallic structures
Melting and boiling points rise across the three metals because of the increasing strength of the metallic bonds
...
The atoms also get
smaller and have more protons as you go from sodium to magnesium to aluminium
...

The "sea" is getting more negatively charged
...


Silicon
Silicon has high melting and boiling points because it is a giant covalent structure
...

Because you are talking about a different type of bond, it isn't profitable to try to directly compare silicon's melting and
boiling points with aluminium's
...
Their melting or boiling points will be lower than those of the first four members of the period which have giant
structures
...
Remember the structures of
the molecules:

75

Phosphorus
Phosphorus contains P4 molecules
...

Sulphur
Sulphur consists of S8 rings of atoms
...

Chlorine
Chlorine, Cl2, is a much smaller molecule with comparatively weak van der Waals attractions, and so chlorine will have a
lower melting and boiling point than sulphur or phosphorus
...
The scope for van der Waals attractions between these is very limited and
so the melting and boiling points of argon are lower again
...
These are the oxides where the Period 3
elements are in their highest oxidation states
...

The structures
The trend in structure is from the metallic oxides containing giant structures of ions on the left of the period via a giant
covalent oxide (silicon dioxide) in the middle to molecular oxides on the right
...

The oxides of phosphorus, sulphur and chlorine consist of individual molecules - some small and simple; others polymeric
...
These vary in
size depending on the size, shape and polarity of the various molecules - but will always be much weaker than the ionic or
covalent bonds you need to break in a giant structure
...

The metallic oxides
The structures
Sodium, magnesium and aluminium oxides consist of giant structures containing metal ions and oxide ions
...
Melting and boiling points
There are strong attractions between the ions in each of these oxides and these attractions need a lot of heat energy to break
...

Electrical conductivity
None of these conducts electricity in the solid state, but electrolysis is possible if they are molten
...
The only important example of this is in the electrolysis of
aluminium oxide in the manufacture of aluminium
...
If it sublimes, you won't get any liquid to electrolyse!
Magnesium and aluminium oxides have melting points far too high to be able to electrolyse them in a simple lab
...
Silicon dioxide is a giant
covalent structure
...
The easiest one to remember and
draw is based on the diamond structure
...
To turn it into silicon dioxide, all you
need to do is to modify the silicon structure by including some oxygen atoms
...
Don't forget that this is just a tiny part of a giant
structure extending in all 3 dimensions
...
Very strong silicon-oxygen covalent bonds
have to be broken throughout the structure before melting occurs
...

Because you are talking about a different form of bonding, it doesn't make sense to try to compare these values directly with
the metallic oxides
...
Electrical conductivity
Silicon dioxide doesn't have any mobile electrons or ions - so it doesn't conduct electricity either as a solid or a liquid
...
Some of these molecules are fairly simple others are polymeric
...

Melting and boiling points of these oxides will be much lower than those of the metal oxides or silicon dioxide
...
The strength of these will vary depending on the size of the molecules
...
None of them contains ions or free electrons
...

Phosphorus(III) oxide
Phosphorus(III) oxide is a white solid, melting at 24°C and boiling at 173°C
...

Phosphorus(V) oxide
Phosphorus(V) oxide is also a white solid, subliming (turning straight from solid to vapour) at 300°C
...


78

Solid phosphorus(V) oxide exists in several different forms - some of them polymeric
...

This is most easily drawn starting from P4O6
...


The sulphur oxides
Sulphur has two common oxides, sulphur dioxide (sulphur(IV) oxide), SO2, and sulphur trioxide (sulphur(VI) oxide), SO3
...
It consists of simple SO2
molecules
...
The bent shape of SO2 is due to this lone pair
...
It reacts very rapidly with water vapour in the air
to form sulphuric acid
...

Gaseous sulphur trioxide consists of simple SO3 molecules in which all six of the sulphur's outer electrons are involved in the
bonding
...
The simplest one is a trimer, S3O9, where three SO3 molecules are joined up
and arranged in a ring
...
For example:

The fact that the simple molecules join up in this way to make bigger structures is what makes the sulphur trioxide a solid
rather than a gas
...
It consists of simple small molecules
...

Chlorine(VII) oxide
In chlorine(VII) oxide, the chlorine uses all of its seven outer electrons in bonds with oxygen
...

Chlorine(VII) oxide is a colourless oily liquid at room temperature
...
In fact, the shape is tetrahedral around both
chlorines, and V-shaped around the central oxygen
...

Aluminium hydroxide
Aluminium hydroxide, like aluminium oxide, is amphoteric - it has both basic and acidic properties
...
These compounds are all acidic - ranging from
the very weakly acidic silicic acids (one of which is shown below) to the very strong sulphuric or chloric(VII) acids
...
In each case the -OH group is covalently bound to the Period 3 element, and in each
case it is possible for the hydrogens on these -OH groups to be removed by a base
...

But they vary considerably in strength:
•
•

Orthosilicic acid is very weak indeed
...

Sulphuric acid and chloric(VII) acids are both very strong acids
...
This in turn depends on how much the negative charge can be spread around the rest of the ion
...
The lost hydrogen ion will be easily recaptured and the acid will be weak
...
The acid will then be strong
...

There are other acids (also containing -OH groups) formed by these elements, but these are the ones where the Period 3
element is in its highest oxidation state
...

Like sodium or magnesium hydroxides, it will react with acids
...

With dilute hydrochloric acid, a colourless solution of aluminium chloride is formed
...
It will react with sodium hydroxide solution to give a
colourless solution of sodium tetrahydroxoaluminate
...


The structure of the ion doesn't stay like this:
Instead, the negative charge is delocalised over the whole ion, and all four chlorine-oxygen bonds are
identical
...
That's still an effective
delocalisation, and sulphuric acid is almost as strong as chloric(VII) acid
...

In orthosilicic acid, there aren't any silicon-oxygen double bonds to delocalise the charge
...


CHEMICAL REACTIONS OF THE PERIOD 3 ELEMENTS
Reactions with water Sodium
Sodium has a very exothermic reaction with cold water producing hydrogen and a colourless solution of sodium hydroxide
...

A very clean coil of magnesium dropped into cold water eventually gets covered in small bubbles of hydrogen which float it
to the surface
...


Magnesium burns in steam with its typical white flame to produce white magnesium oxide and hydrogen
...
The reaction is relatively slow because of the
existing strong aluminium oxide layer on the metal, and the build-up of even more oxide during the reaction
...
The truth
seems to depend on the precise form of silicon you are using
...
Most sources suggest
that this form of silicon will react with steam at red heat to produce silicon dioxide and hydrogen
...

Phosphorus and sulphur
These have no reaction with water
...
A reversible reaction takes place to produce a mixture of
hydrochloric acid and chloric(I) acid (hypochlorous acid)
...

Reactions with oxygen Sodium
Sodium burns in oxygen with an orange flame to produce a white solid mixture of sodium oxide and sodium peroxide
...


Aluminium
Aluminium will burn in oxygen if it is powdered, otherwise the strong oxide layer on the aluminium tends to inhibit the
reaction
...
White aluminium oxide is formed
...
Silicon dioxide is produced
...

The proportions of these depend on the amount of oxygen available
...


84

For the phosphorus(III) oxide:

For the phosphorus(V) oxide:

Sulphur
Sulphur burns in air or oxygen on gentle heating with a pale blue flame
...


Chlorine and argon
Despite having several oxides, chlorine won't react directly with oxygen
...

Reactions with chlorine
Sodium
Sodium burns in chlorine with a bright orange flame
...


Magnesium
Magnesium burns with its usual intense white flame to give white magnesium chloride
...
The aluminium
burns in the stream of chlorine to produce very pale yellow aluminium chloride
...


Silicon
If chlorine is passed over silicon powder heated in a tube, it reacts to produce silicon tetrachloride
...


Phosphorus

85

White phosphorus burns in chlorine to produce a mixture of two chlorides, phosphorus(III) chloride and phosphorus(V)
chloride (phosphorus trichloride and phosphorus pentachloride)
...


86

Phosphorus(V) chloride is an off-white (going towards yellow) solid
...


Chlorine and argon
It obviously doesn't make sense to talk about chlorine reacting with itself, and argon doesn't react with chlorine
...

However, as you go down the Group, there are more and more examples where the oxidation state is +2, such as SnCl 2, PbO,
and Pb2+
...


An example from carbon chemistry
The only common example of the +2 oxidation state in carbon chemistry occurs in carbon monoxide, CO
...

For example, carbon monoxide reduces many hot metal oxides to the metal - a reaction which is used, for example, in the
extraction of iron in a blast furnace
...
However, tin(IV) is still the more stable oxidation state of tin
...
This is best shown in the fact that
Sn2+ ions in solution are good reducing agents
...
In the process, the tin(II) ions are oxidised to tin(IV) ions
...
For example, tin(II) chloride solution will reduce iron(III) chloride
solution to iron(II) chloride solution
...


Tin(II) ions will also, of course, be easily oxidised by powerful oxidising agents like acidified potassium manganate(VII)
solution (potassium permanganate solution)
...


And as a final example
...
This reaction involves the tin first being oxidised to tin(II) ions and then further to the preferred tin(IV) ions
...
This time, the lead(II) oxidation state is the more stable, and there is a strong tendency for
lead(IV) compounds to react to give lead(II) compounds
...
and lead(IV) oxide decomposes on heating to give lead(II) oxide and oxygen
...
Once again, the lead is reduced from the +4 to the more stable +2 state
...

All of the elements in the group have the outer electronic structure ns2npx1npy1, where n varies from 2 (for carbon) to 6 (for
lead)
...

As you get closer to the bottom of the Group, there is an increasing tendency for the s2 pair not to be used in the bonding
...

However, just giving it a name like "inert pair effect" explains nothing
...


The inert pair effect in the formation of ionic bonds
If the elements in Group 4 form 2+ ions, they will lose the p electrons, leaving the s2 pair unused
...

You would normally expect ionisation energies to fall as you go down a Group as the electrons get further from the nucleus
...

This first chart shows how the total ionisation energy needed to form the 2+ ions varies as you go down the Group
...


89

Notice the slight increase between tin and lead
...

However, if you look at the pattern for the loss of all four electrons, the discrepancy between tin and lead is much more
marked
...


Again, the values are all in kJ mol-1, and the two charts are to approximately the same scale
...
With the heavier elements like lead, there is what is known as a
relativistic contraction of the electrons which tends to draw the electrons closer to the nucleus than you would expect
...
The heavier the element, the greater this effect
...

In the case of lead, the relativistic contraction makes it energetically more difficult to remove the 6s electrons than you might
expect
...
That means that it doesn't make energetic sense for lead to form 4+ ions
...

Using the electrons-in-boxes notation, the outer electronic structure of carbon looks like this:

There are only two unpaired electrons
...


That leaves 4 unpaired electrons which (after hybridisation) can go on to form 4 covalent bonds
...

Each covalent bond that forms releases energy, and this is more than enough to supply the energy needed for the promotion
...
Bond energies tend to fall as atoms get bigger and the bonding pair is further from the two nuclei and better screened
from them
...

This would would be made worse, of course, if the energy gap between the 6s and 6p orbitals was increased by the relativistic
contraction of the 6s orbital
...
Carbon dioxide is a gas whereas silicon dioxide is a hard high-melting solid
...

This obviously reflects a difference in structure between carbon dioxide and the dioxides of the rest of the Group
...
Carbon can form simple molecules with
oxygen because it can form double bonds with the oxygen
...


When carbon forms bonds with oxygen, it first promotes one of the electrons in the 2s level into the empty 2p level
...


It now reshuffles those electrons slightly by hybridising the 2s electron and one of the 2p electrons to make two sp1 hybrid
orbitals of equal energy
...


What these look like in the atom (using the same colour coding) is:

Notice that the two green lobes are two different hybrid orbitals - arranged as far apart from each other as possible
...


92

So that's how the carbon is organised just before it bonds
...

Oxygen's electronic structure is 1s22s22px22py12pz1
...
This time, sp2 hybrids are formed with the s orbital and two of the p orbitals being
rearranged to give 3 orbitals of equal energy - leaving a temporarily unaffected p orbital
...


Now line up the two oxygens and the carbon prior to bonding them
...
These are
properly called sigma bonds, and are shown as orange in the next diagram
...


93

Sideways overlap between the two sets of p orbitals produces two pi bonds - similar to the pi bond found in, say, ethene
...


So
...


The structure of silicon dioxide
Silicon doesn't double bond with oxygen
...
That means that silicon-oxygen bonds will be longer than carbon-oxygen bonds
...
With the
longer silicon-oxygen bonds, the p orbitals on the silicon and the oxygen aren't quite close enough together to allow enough
sideways overlap to give a stable pi bond
...

There are various different structures for silicon dioxide
...

This means that silicon dioxide is a giant covalent structure
...

The acid-base behaviour of the Group 4 oxides
The oxides of the elements at the top of Group 4 are acidic, but acidity of the oxides falls as you go down the Group
...

An oxide which can show both acidic and basic properties is said to be amphoteric
...


Carbon and silicon oxides
Carbon monoxide
Carbon monoxide is usually treated as if it was a neutral oxide, but in fact it is very, very slightly acidic
...


The fact that the carbon monoxide reacts with the basic hydroxide ion shows that it must be acidic
...
With water
Silicon dioxide doesn't react with water, because of the difficulty of breaking up the giant covalent structure
...

Overall, this reaction is:

95

The solution of carbon dioxide in water is sometimes known as carbonic acid, but in fact only about 0
...
The position of equilibrium is well to the left-hand side
...


Silicon dioxide also reacts with sodium hydroxide solution, but only if it is hot and concentrated
...


You may also be familiar with one of the reactions happening in the Blast Furnace extraction of iron - in which calcium oxide
(from the limestone which is one of the raw materials) reacts with silicon dioxide to produce a liquid slag, calcium silicate
...


Gernamium, tin and lead oxides
The monoxides
All of these oxides are amphoteric - they show both basic and acidic properties
...

For example, they all react with concentrated hydrochloric acid
...
where X can be Ge and Sn, but unfortunately needs modifying a bit for lead
...


However, in this example we are talking about using concentrated hydrochloric acid
...
These ionic complexes are soluble in water and so the problem disappears
...
This time we can generalise without exception:

Lead(II) oxide, for example, would react to give PbO22- - plumbate(II) ions
...

The basic nature of the dioxides
The dioxides react with concentrated hydrochloric acid first to give compounds of the type XCl 4:

These will react with excess chloride ions in the hydrochloric acid to give complexes such as XCl62-
...
If the reaction is done any warmer,
the lead(IV) chloride decomposes to give lead(II) chloride and chlorine gas
...

The acidic nature of the dioxides
The dioxides will react with hot concentrated sodium hydroxide solution to give soluble complexes of the form [X(OH)6]2-
...
In that case, the equation is different
...


97

They are all simple covalent molecules with a typical tetrahedral shape
...

(Although at room temperature, lead(IV) chloride will tend to decompose to give lead(II) chloride and chlorine gas - see
below
...
It is very slightly soluble in cold water, but more soluble in hot water
...

Stability
At the top of Group 4, the most stable oxidation state shown by the elements is +4
...
These therefore have no tendency to split up to give dichlorides
...

Lead(IV) chloride decomposes at room temperature to give the more stable lead(II) chloride and chlorine gas
...
If you add it to water, it simply forms a separate layer underneath the layer of
water
...
The reaction would have to start by the water
molecule's oxygen attaching itself to the carbon atom via the oxygen's lone pair
...

There are two problems with this
...



...
There is going to be a lot of repulsion between the various lone pairs on all the atoms
surrounding the carbon
...
A very unstable
transition state means a very high activation energy for the reaction
...

If it could attach before the chlorine starts to break away, that would be an advantage
...
But in the case of a carbon atom, that
isn't possible
...

The silicon atom is bigger, and so there is more room around it for the water molecule to attack, and the transition state will
be less cluttered
...
Carbon doesn't have 2d orbitals because there are no such things
...

This means that the oxygen can bond to the silicon before the need to break a silicon-chlorine bond
...

So
...


Liquid SiCl4 fumes in moist air for this reason - it is reacting with water vapour in the air
...
You will get lead(IV) oxide produced as
a brown solid and fumes of hydrogen chloride given off
...
)

Lead(II) chloride
Unlike the tetrachlorides, lead(II) chloride can be thought of as ionic
...
Looked at simply, solubility in water involves the break-up of the ionic lattice and the hydration of the lead(II) and
chloride ions to give Pb2+(aq) and Cl-(aq)
...


If more sodium hydroxide solution is added, the precipitate redissolves to give a colourless solution which might be called
sodium plumbate(II) solution - but could be called by a lot of alternative names depending on exactly how the formula is
written!

Making lead(II) chloride
Lead(II) chloride can be made as a white precipitate by adding a solution containing chloride ions to lead(II) nitrate solution
...


Making lead(II) iodide
If you add colourless potassium iodide solution (or any other source of iodide ions in solution) to a solution of lead(II) nitrate,
a bright yellow precipitate of lead(II) iodide is produced
...
The
easiest thing to add is usually dilute sulphuric acid - but any other soluble sulphate would do
...
It also explains the concept of a conjugate pair - an acid and its
conjugate base, or a base and its conjugate acid
...

Bases are substances which produce hydroxide ions in solution
...


Limitations of the theory
Hydrochloric acid is neutralised by both sodium hydroxide solution and ammonia solution
...

These are clearly very similar reactions
...

However, in the ammonia case, there don't appear to be any hydroxide ions!
You can get around this by saying that the ammonia reacts with the water it is dissolved in to produce
ammonium ions and hydroxide ions:

This is a reversible reaction, and in a typical dilute ammonia solution, about 99% of the ammonia
remains as ammonia molecules
...

However, this same reaction also happens between ammonia gas and hydrogen chloride gas
...
The Arrhenius theory wouldn't count this as an acid-base reaction, despite the fact that it is
producing the same product as when the two substances were in solution
...

A base is a proton (hydrogen ion) acceptor
...

Hydroxide ions are still bases because they accept hydrogen ions from acids and form water
...

When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride
molecule gives a proton (a hydrogen ion) to a water molecule
...
Hydroxonium ions,
H3O+, are produced
...
For example, a proton is transferred from a hydroxonium ion to a hydroxide ion to make water
...

The hydrogen chloride / ammonia problem
This is no longer a problem using the Bronsted-Lowry theory
...
The
hydrogen becomes attached to the lone pair on the nitrogen of the ammonia via a co-ordinate bond
...


Conjugate pairs
When hydrogen chloride dissolves in water, almost 100% of it reacts with the water to produce
hydroxonium ions and chloride ions
...


103

In fact, the reaction between HCl and water is reversible, but only to a very minor extent
...


Thinking about the forward reaction:
•

The HA is an acid because it is donating a proton (hydrogen ion) to the water
...


But there is also a back reaction between the hydroxonium ion and the A- ion:
•

The H3O+ is an acid because it is donating a proton (hydrogen ion) to the A- ion
...


The reversible reaction contains two acids and two bases
...


When the acid, HA, loses a proton it forms a base, A-
...
These two are a conjugate pair
...

If you are thinking about HA as the acid, then A- is its conjugate base
...

The water and the hydroxonium ion are also a conjugate pair
...

Thinking about the hydroxonium ion as an acid, then water is its conjugate base
...

A second example of conjugate pairs
This is the reaction between ammonia and water that we looked at earlier:
104

Think first about the forward reaction
...
The ammonium ion is its conjugate acid - it can release that hydrogen ion again to reform the
ammonia
...
The hydroxide ion can accept
a hydrogen ion to reform the water
...
The
hydroxide ion is a base and water is its conjugate acid
...

A substance which can act as either an acid or a base is described as being amphoteric
...
The two words are related and easily confused
...
Water is a good
example of such a compound
...

The "protic" part of the word refers to the hydrogen ions (protons) either being donated or accepted
...

But as well as being amphiprotic, these compounds are also amphoteric
...
So what is the difference between the two terms?
All amphiprotic substances are also amphoteric - but the reverse isn't true
...
There is a whole new definition of acid-base
behaviour that you are just about to meet (the Lewis theory) which doesn't necessarily involve hydrogen ions at all
...

Some metal oxides (like aluminium oxide) are amphoteric - they react both as acids and bases
...
That's not a problem as far as the definition of amphiprotic
is concerned - but the reaction as an acid is
...


105

You can think of lone pairs on hydroxide ions as forming dative covalent (coordinate) bonds with empty orbitals in the
aluminium ions
...
So aluminium oxide can act as both an
acid and a base - and so is amphoteric
...


The Lewis Theory of acids and bases
This theory extends well beyond the things you normally think of as acids and bases
...

A base is an electron pair donor
...
Three Bronsted-Lowry bases we've looked at are hydroxide ions, ammonia and water,
and they are typical of all the rest
...
The reason they are combining with hydrogen ions is that they have lone pairs of
electrons - which is what the Lewis theory says
...

So how does this extend the concept of a base? At the moment it doesn't - it just looks at it from a
different angle
...

Here is a reaction which you will find talked about on the page dealing with co-ordinate bonding
...


107

As far as the ammonia is concerned, it is behaving exactly the same as when it reacts with a hydrogen
ion - it is using its lone pair to form a co-ordinate bond
...


Note: If you haven't already read the page about co-ordinate bonding you should do so now
...

Use the BACK button on your browser to return quickly to this page
...
In the above example, the BF3 is acting as the Lewis acid by
accepting the nitrogen's lone pair
...

This is an extension of the term acid well beyond any common use
...
Textbooks often write this as if the
ammonia is donating its lone pair to a hydrogen ion - a simple proton with no electrons around it
...
They are so reactive
that they are always attached to something else
...

There isn't an empty orbital anywhere on the HCl which can accept a pair of electrons
...
The electrons in the hydrogen-chlorine bond will be attracted towards the chlorine end,
leaving the hydrogen slightly positive and the chlorine slightly negative
...
As it approaches it, the electrons in the hydrogen-chlorine bond are repelled still further
towards the chlorine
...

This is best shown using the "curly arrow" notation commonly used in organic reaction mechanisms
...
It is accepting a pair of electrons from the ammonia,
and in the process it breaks up
...


A final comment on Lewis acids and bases
A Lewis acid is an electron pair acceptor
...


Note: Remember this by thinking of ammonia acting as a base
...
Ammonia is basic because of its lone pair
...
Acids are the opposite

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