Trends in period 3
Atomic radius
First ionization energy
Melting and boiling
points
Atomic Radius
• The atomic radius of an element is difficult to precisely define because
of the uncertainty over the size of the electron cloud. Several definitions
are used.
One definition is half the shortest internuclear distance found
in the structure of the element.
For non-metallic elements, the covalent radius is often used as the
atomic radius. This is half the internuclear distance between two
identical atoms in a single covalent bond.
• In Period 3 of the Periodic Table,
the 3s and 3p orbitals are filling
with electrons
• The shortened versions of the
electronic structures for the eight
elements are:
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+.
• X(g) → X+(g) + e• It is the energy needed to carry out this change per mole of X.
First ionisation energy
• The general trend is
upwards, but this is
broken by falls between
magnesium and
aluminium, and between
phosphorus and sulphur
First ionisation energy
• Explaining the pattern
• First ionisation energy is governed by:
• the charge on the nucleus;
• the distance of the outer electron from the nucleus;
• the amount of screening by inner electrons;
• whether the electron is alone in an orbital or one of a pair.
First ionisation energy
• The upward trend
• In the whole of period 3, the outer electrons are in 3-level orbitals
• These are all the same sort of distances from the nucleus, and are
screened by the same electrons in the first and second levels.
First ionisation energy
• The major difference is the increasing number of protons in the nucleus
as you go from sodium across to argon
• That causes greater attraction between the nucleus and the electrons
and so increases the ionisation energies.
• The increasing nuclear charge also drags the outer electrons in closer
to the nucleus
• That increases ionisation energies still more as you go across the
period.
First ionisation energy
The fall at aluminium
• Aluminium's outer electron is in a 3p orbital rather than a 3s.
• The 3p electron is slightly more distant from the nucleus than the 3s,
and partially screened by the 3s electrons as well as the inner electrons
• Both of these factors offset the effect of the extra proton.
First ionisation energy
The fall at sulphur
• The screening is identical in phosphorus and sulphur (from the inner
electrons and, to some extent, from the 3s electrons), and the electron is
being removed from an identical orbital.
• The difference is that in the sulphur case the electron being removed is one
of the 3px2 pair
• The repulsion between the two electrons in the same orbital means that the
electron is easier to remove than it would otherwise be.
Atomic radius
• The diagram shows how the
atomic radius changes as you go
across Period 3.
Atomic radius
• 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.
Atomic radius
• 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 increasing number of protons in the nucleus as you go across the
period pulls the bonding electrons more tightly to it
• The amount of screening is constant for all of these elements.
Electronegativity
• Electronegativity is a measure of the tendency of an atom to attract a
bonding pair of electrons
• The Pauling scale is the most commonly used
• Fluorine (the most electronegative element) is assigned a value of 4.0,
and values range down to caesium and francium which are the least
electronegative at 0.7.
Electronegativity
The trend across Period 3 looks like this:
Electronegativity
• Electronegativity is about the tendency of an atom to attract a bonding
pair of electrons
• Since argon doesn't form covalent bonds, you can't assign it an
electronegativity.
Electronegativity
• The trend is explained in exactly the same way as the trend in atomic
radii
• As you go across the period, the bonding electrons are always in the
same level - the 3-level
• They are always being screened by the same inner electrons.
• All that differs is the number of protons in the nucleus. As you go from
sodium to chlorine, the number of protons steadily increases and so
attracts the bonding pair more closely.
Physical Properties
• Physical Properties include
• The electrical conductivity
• The melting
• The boiling points of the elements
• To understand these, you first have to understand the structure of each
of the elements.
Structures of the elements
The structures of the elements change as you go across the period
The first three are metallic, silicon is giant covalent, and the rest are
simple molecules.
Structures of the elements
• Three metallic structures
• Sodium, magnesium and aluminium all have metallic structures
• In sodium, only one electron per atom is involved in the metallic bond the single 3s electron
• In magnesium, both of its outer electrons are involved, and in
aluminium all three.
Structures of the elements
• The other difference is the way the atoms are packed in the metal crystal.
• Sodium is 8-co-ordinated - each sodium atom is touched by only 8 other atoms.
• Both magnesium and aluminium are 12-co-ordinated (although in slightly different
ways). 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.
Structures of the elements
• A giant covalent structure
• Silicon has a giant covalent
structure just like diamond
• The structure is held together by
strong covalent bonds in all three
dimensions.
Structures of the elements
• Four simple molecular structures
• The structures of phosphorus and sulphur vary depending on the type of
phosphorus or sulphur you are talking about
• For phosphorus, assuming the common white phosphorus
• For sulphur, assuming one of the crystalline forms - rhombic or
monoclinic sulphur.
Structures of the elements
Structures of the elements
• The atoms in each of these molecules are held together by covalent
bonds (apart, of course, from argon).
• In the liquid or solid state, the molecules are held close to each other by
van der Waals dispersion forces.
Electrical conductivity
• Sodium, magnesium and aluminium are all good conductors of electricity
• Conductivity increases as you go from sodium to magnesium to aluminium
• Silicon is a semiconductor.
• None of the rest conduct electricity.
Electrical conductivity
• The three metals, of course, conduct electricity because the delocalised
electrons (the "sea of electrons") are free to move throughout the solid or
the liquid metal.
• In the silicon case, explaining how semiconductors conduct electricity is
beyond the scope of A level chemistry courses.
• The rest don't conduct electricity because they are simple molecular
substances. There are no electrons free to move around
Melting and boiling points
The chart shows how the melting
and boiling points of the elements
change as you go across the
period
The figures are plotted in kelvin
rather than °C to avoid having
negative values.
Melting and boiling points
• Melting and boiling points rise across the three metals because of the
increasing strength of the metallic bonds.
• The number of electrons which each atom can contribute to the
delocalised "sea of electrons" increases
• The atoms also get smaller and have more protons as you go from
sodium to magnesium to aluminium.
Melting and boiling points
• The attractions and therefore the melting and boiling points increase because:
• The nuclei of the atoms are getting more positively charged.
• The "sea" is getting more negatively charged.
• The "sea" is getting progressively nearer to the nuclei and so more strongly
attracted.
Melting and boiling points
• Silicon has high melting and boiling points because it is a giant covalent
structure
• You have to break strong covalent bonds before it will melt or boil.
• 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.
• Phosphorus, sulphur, chlorine and argon are
simple molecular substances with only van der
Waals attractions between the molecules
• Their melting or boiling points will be lower
than those of the first four members of the
period which have giant structures.
• The sizes of the melting and boiling points are
governed entirely by the sizes of the
molecules. Remember the structures of the
molecules:
• Phosphorus contains P4 molecules
• To melt phosphorus you don't have to break any covalent bonds - just
the much weaker van der Waals forces between the molecules.
• Sulphur consists of S8 rings of atoms. The molecules are bigger than
phosphorus molecules, and so the van der Waals attractions will be
stronger, leading to a higher melting and boiling point.
• 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.
• Argon molecules are just single argon atoms, Ar. The scope for van der
Waals attractions between these is very limited and so the melting and
boiling points of argon are lower again.
Reaction with water
• Sodium has a very exothermic reaction with cold water producing
hydrogen and a colourless solution of sodium hydroxide.
• 2Na + 2H2O → 2NaOH + 2H2
Reaction with water
Magnesium has a very slight reaction with cold water, but burns in
steam.
 A very clean coil of magnesium dropped into cold water eventually gets
covered in small bubbles of hydrogen which float it to the surface.
Reaction with water
• Magnesium hydroxide is formed as a very thin layer on the magnesium and
this tends to stop the reaction.
• Mg + 2H2O → Mg(OH)2 + H2
• Magnesium burns in steam with its typical white flame to produce white
magnesium oxide and hydrogen.
• Mg + H2O → MgO + H2
Reaction with water
• Aluminium powder heated in steam produces hydrogen and aluminium
oxide
• 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.
• 2Al + 3H2O →
Al2O3 + 3H2
Reaction with water
• Silicone, phosphorus and sulphur do not react. They are insoluble in
water
Reaction with water
• Chlorine dissolves in water to some extent to give a green solution
• A reversible reaction takes place to produce a mixture of hydrochloric
acid and chloric(I) acid (hypochlorous acid).
• Cl2 + H2O → HCl + HOCl
Reaction with oxygen
 Sodium burns in oxygen with an orange flame to produce a white solid mixture of sodium oxide and
sodium peroxide.
• For the simple oxide:
• 4Na + O2 → 2Na2O
• For the peroxide:
• 2Na + O2 → Na2O2
Reaction with oxygen
• Magnesium burns in oxygen with an intense white flame to give white
solid magnesium oxide.
• 2Mg + O2 →
2MgO
Reaction with oxygen
• Aluminium will burn in oxygen if it is powdered, otherwise the strong
oxide layer on the aluminium tends to inhibit the reaction
• If you sprinkle aluminium powder into a Bunsen flame, you get white
sparkles
• White aluminium oxide is formed.
• 4Al + 3O2 → 2Al2O3
Reaction with oxygen
• Silicon will burn in oxygen if heated strongly enough. Silicon dioxide is
produced.
• Si + O2 →
SiO2
Reaction with oxygen
• 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.
• The proportions of these depend on the amount of oxygen available
Reaction with oxygen
• . In an excess of oxygen, the product will be almost entirely phosphorus(V) oxide.
• For the phosphorus(III) oxide:
• P4 + 3O2
P4O6
• For the phosphorus(V) oxide:
• P4 + 5O2
P4O10
Reaction with oxygen
• Sulphur burns in air or oxygen on gentle heating with a pale blue flame
• It produces colourless sulphur dioxide gas.
• S + O2 → SO2
Reaction with oxygen
• Despite having several oxides, chlorine won't react directly with oxygen
• Argon doesn't react either.
Reaction with chlorine
• Sodium burns in chlorine with a bright orange flame
• White solid sodium chloride is produced.
• 2Na + Cl2 → 2NaCl
Reaction with chlorine
• Magnesium burns with its usual intense white flame to give white
magnesium chloride.
• Mg + Cl2
MgCl2
Reaction with chlorine
• Aluminium is often reacted with chlorine by passing dry chlorine over
aluminium foil heated in a long tube
• The aluminium burns in the stream of chlorine to produce very pale
yellow aluminium chloride. This sublimes (turns straight from solid to
vapour and back again) and collects further down the tube where it is
cooler.
• 2Al + 3Cl2 →
2AlCl3
Reaction with chlorine
• If chlorine is passed over silicon powder heated in a tube, it reacts to
produce silicon tetrachloride
• This is a colourless liquid which vaporises and can be condensed
further along the apparatus.
• Si + 2Cl2 → SiCl4
Reaction with chlorine
• If a stream of chlorine is passed over some heated sulphur, it reacts to
form an orange, evil-smelling liquid, disulphur dichloride, S2Cl2.
• 2S + Cl2 → S2Cl2
Reactions with chlorine
• Chlorine and argon
• Obviously chlorine will not react with itself, and argon doesn't react
with chlorine
PHYSICAL PROPERTIES OF THE PERIOD 3
OXIDES
• The trend in structure is from the metallic oxides containing giant
structures of ions on the left of the period a giant covalent oxide (silicon
dioxide) in the middle to molecular oxides on the right
• These oxides tend to be gases, liquids or low melting point solids.
• The attractive forces between these molecules will be van der Waals
dispersion and dipole-dipole interactions
• 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 oxides of phosphorus, sulphur and chlorine consist of individual
molecules - some small and simple; others polymeric
• Melting and boiling points
• The giant structures (the metal oxides and silicon dioxide) will have high
melting and boiling points because a lot of energy is needed to break
the strong bonds (ionic or covalent) operating in three dimensions.
Metallic oxides structure
Sodium, magnesium and aluminium oxides consist of giant structures
containing metal ions and oxide ions
Magnesium oxide has a structure just like sodium chloride
The other two have more complicated arrangements of the ions beyond
the scope of syllabuses
Melting and boiling points of metallic oxides
• There are strong attractions between the ions in each of these oxides
and these attractions need a lot of heat energy to break
• These oxides therefore have high melting and boiling points.
Electrical conductivity
• None of these conducts electricity in the solid state, but electrolysis is
possible if they are molten
• They conduct electricity because of the movement and discharge of the ions
present. The only important example of this is in the electrolysis of aluminium
oxide in the manufacture of aluminium.
• Whether you can electrolyse molten sodium oxide depends, of course, on
whether it actually melts instead of subliming or decomposing under ordinary
circumstances. 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 (silicon(IV) oxide)
• The structure
• The electronegativity of the elements increases as you go across the
period, and by the time you get to silicon, there isn't enough
electronegativity difference between the silicon and the oxygen to form
an ionic bond. Silicon dioxide is a giant covalent structure.
• There are three different crystal forms of silicon dioxide. The easiest one
to remember and draw is based on the diamond structure.
• Crystalline silicon has the same structure as diamond. To turn it into
silicon dioxide, all you need to do is to modify the silicon structure by
including some oxygen atoms.
• Notice that each silicon atom is bridged to its neighbours by an oxygen
atom. Don't forget that this is just a tiny part of a giant structure
extending in all 3 dimensions.
• Melting and boiling points
• Silicon dioxide has a high melting point - varying depending on what the
particular structure is (remember that the structure given is only one of three
possible structures), but they are all around 1700°C. Very strong siliconoxygen covalent bonds have to be broken throughout the structure before
melting occurs. Silicon dioxide boils at 2230°C.
• 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. What
you can safely say is that because the metallic oxides and silicon dioxide have
giant structures, the melting and boiling points are all high.
• 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.
• The molecular oxides
• Phosphorus, sulphur and chlorine all form oxides which consist of
molecules. Some of these molecules are fairly simple - others are
polymeric. We are just going to look at some of the simple ones.
• Melting and boiling points of these oxides will be much lower than those
of the metal oxides or silicon dioxide. The intermolecular forces holding
one molecule to its neighbours will be van der Waals dispersion forces
or dipole-dipole interactions. The strength of these will vary depending
on the size of the molecules.
• None of these oxides conducts electricity either as solids or as liquids.
None of them contains ions or free electrons.
• The phosphorus oxides
• Phosphorus has two common oxides, phosphorus(III) oxide, P4O6, and phosphorus(V) oxide,
P4O10.
• Phosphorus(III) oxide
• Phosphorus(III) oxide is a white solid, melting at 24°C and boiling at 173°C.
• The structure of its molecule is best worked out starting from a P4 molecule which is a little
tetrahedron.
• Pull this apart so that you can see the bonds
• and then replace the bonds by new bonds linking the phosphorus
atoms via oxygen atoms. These will be in a V-shape (rather like in
water), but you probably wouldn't be penalised if you drew them on a
straight line between the phosphorus atoms in an exam
• and then replace the bonds by new bonds linking the phosphorus
atoms via oxygen atoms. These will be in a V-shape (rather like in
water), but you probably wouldn't be penalised if you drew them on a
straight line between the phosphorus atoms in an exam
Oxidation states
• The oxides of the elements in period 3 all exist in a positive oxidation
states because oxygen is more electronegative than any of these
elements
• The maximum oxidation state of each of the elements in period 3 in their
oxides rises across the periodd
• This corresponds to the ability of the elements to use all the electrons in
their outermost electron shells in bonding
• This leads to an expanded octet for oxides of P, S and Cl
Oxide
Na2O
MgO
Al2O3
SiO2
P4O10
SO3
Cl2O7
Max.
oxidation
state
+1
+2
+3+
+4
+5
+6+
+7
• The non-metallic oxides can also form oxides in lower oxidation states
e.g. SiO, P2O3 , SO2
• Several lower oxides of chlorine (Cl2O, ClO2 and Cl2O6)
• The chlorides of period 3-elements also exist in positive oxidation states
because chlorine is more electronegative than the other elements in the
period
• the oxidation states of the chlorides show a similar pattern rising to +5 in
PCL 5
• the maximum oxidation state for sulphur in its chlorides However is 2+
Chloride
NaCl
MgCl2
AlCl3
SiCl4
PCL5
SCl2
Maximum
Oxidation
state
+1
+2
+3
+4
+5
+2
• Some of the non-metallic elements can form chlorides in lower oxidation
states e.g. PCl3 S2Cl2
Bonding in chlorides and oxides of period 3
• The structure and bonding of these compounds can be related to the
relative electronegativity of the atoms involved in bonding.
• The greater the difference in electronegativity the more likely it is that
the oxide or chloride will be ionic
• the electronegativity difference between oxygen and sodium
magnesium or aluminum it's so great that 1 two or three electrons
respectively are transferred from the metal atom to the nonmetal atom
• these oxides are therefore ionic
• the other period 3 oxides alright covalently bonded
• silicon dioxide has a giant covalent structure
• although in theory an Si4+ ion can I exist, the fourth ionization energy is
so large that the lattice energy can not compensate for it to make ionic
SiO2 Stable
• Chlorides of P,S and Cl have a simple molecular structure
• Electronegativity differences can also be used to Explain the structure of
period 3 chlorides
• Sodium chloride and magnesium chloride are ionic
• anhydrous aluminum chloride Is it covalently bonded molecule Al2Cl6
• The ionic radius of aluminum is very small and the Al3+ ion is highly
charged
• this high positive charge density tends to pull the electrons in the larger
chlorine ion towards it to such an extent that the electron pairs are
shared between Al and Cl atoms.
• This is called ion polarization
• Chlorides of Si, P and S have a simple molecular structure
Reactions of period 3 oxides and hydroxides
• The oxides of sodium and magnesium rapid water to form hydroxides
• Na2O(s) + H20(l) →2NaOH(aq)
• Magnesium hydroxide is less alkaline because it is less soluble than
sodium hydroxide
• Oxides and hydroxides of sodium and magnesium are basic
• They react with acids to form a salt and water
• MgO (s) + 2HCl (aq) → MgCl2 (aq) + H2O(l)
• Aluminium oxide does not dissolve in water but reacts with acids and alkali
• Al2O3(s) + 3 H2SO4(aq) →Al2(SO4)3 +3 H2O
• Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2NaAl(OH)4(aq) (sodium aluminate)
• A substance Which acts both as an acid and a base is said to be amphoteric
Silicon dioxide is insoluble in water but like Al2O3 reacts with hot Alkali
SiO2(s) + 2NaOH(aq) → Na2SiO3(aq) + H2O(l)
SiO2 is an acidic oxide it does not react with acids