Third Period elements
R.B.Mahajan
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 The trend
 The diagram
shows how the atomic radius
changes as you go across Period 3.
 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 because the all of these
valence electrons are in 3rd level.
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Ionic radius of cation is less than its atom- one shell is
lost & less screening.
Ionic radius of anion is more than its atom for the
contrary reasons to cation.
Let's look at the radii of the simple ions formed by
elements as you go across Period 3 of the Periodic
Table - the elements from Na to Cl.
ion
Na
Mg
Al
Si
P
S
Cl
Radius.
mm
0.095
0.065
0.050
0.041
0.212
0.184
0.181
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 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.
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The trend of variation can be explained on the basis of difference
in their structures.
Sodium, magnesium and aluminum 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 aluminum all three.
Silicon has a giant covalent structure just like diamond. A tiny part
of the structure looks like this:
Phosphorus, Sulfur, Chlorine and Argon has following molecular
structures
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Sodium, magnesium and aluminum are all good conductors
of electricity. Conductivity increases as you go from sodium
to magnesium to aluminum.
Silicon is a semiconductor.
None of the rest conduct electricity.
The three metals, of course, conduct electricity because the
delocalized 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.
With a diamond structure, you mightn't expect it to conduct
electricity, but it does!
The rest don't conduct electricity because they are simple
molecular substances. There are no electrons free to move
around.
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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.
The trend across Period 3 looks like this:
The trend is explained in exactly the same way as the trend
in atomic radii.
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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+.
It is the energy needed to carry out this change per
mole of X.
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Notice that the general trend is upwards, but this is
broken by falls between magnesium and aluminum,
and between phosphorus and sulphur.
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
R.B.Mahajan
pair.
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Name of Element
Symbol
Sodium
Na
Magnesium
Mg
Aluminium
Al
Silicon
Si
Phosphorus
P
Sulfur
S
Chlorine
Cl
Argon
Ar
Atomic Number, z
11
12
13
14
15
16
17
18
Electronic Configuration
2,8,1
2,8,2
2,8,3
2,8,4
2,8,5
2,8,6
2,8,7
2,8,8
Atomic Radius
(picometers)
186
160
143
118
110
102
99
192
1stIonization Energy
(kJ/mol)
502
744
584
793
1017
1006
1257
>1526
Electronegativity (Pauling)
0.93
1.31
1.61
1.9
2.19
2.58
3.16
-
Melting Point (oC)
98
639
660
1410
44
113
-101
-189
Boiling Point (oC)
883
1090
2467
2680
280
445
-35
-186
Metallic Character
metal
metal
metal
semi-metal
(metalloid)
non-metal
non-metal
non-metal
non-metal
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Sodium
Sodium burns in oxygen with an orange flame to
produce a white solid mixture of sodium oxide and
sodium peroxide.
For the simple oxide:
4 Na + O2
2Na2O
For the peroxide:
2Na + O2
Na2O2
Magnesium
Magnesium burns in oxygen with an intense white
flame to give white solid magnesium oxide.
2Mg + O2
MgO
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Aluminum
Aluminum will burn in oxygen if it is powdered,
otherwise the strong oxide layer on the aluminum tends
to inhibit the reaction. If you sprinkle aluminum powder
into a Bunsen flame, you get white sparkles. White
aluminum oxide is formed.
4 Al + 3O2
2Al2O3
Silicon
Silicon will burn in oxygen if heated strongly enough.
Silicon dioxide is produced.
Si + O2
SiO2
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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. 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
Sulphur
Sulphur burns in air or oxygen on gentle heating with a pale blue
flame. It produces colourless sulphur dioxide gas.
S + O2
SO2
Chlorine and argon
Despite having several oxides, chlorine won't react directly with
oxygen.
Argon doesn't react either.
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Sodium- Sodium burns in chlorine with a bright orange flame. White
solid sodium chloride is produced.
2Na + Cl2 → 2NaCl
Magnesium- Magnesium burns with its usual intense white flame to
give white magnesium chloride
Mg + Cl2
→ MgCl2
Aluminium- 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
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Silicon
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
Phosphorus
White phosphorus burns in chlorine to produce a mixture of two chlorides,
phosphorus(III) chloride and phosphorus(V) chloride (phosphorus trichloride and
phosphorus pentachloride).
Phosphorus(III) chloride is a colourless fuming liquid.
P4 + 6Cl2 → 4PCl3
Phosphorus(V) chloride is an off-white (going towards yellow) solid.
P4 + 10Cl2 → 4PCl5
Sulphur
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 or S + Cl2
→ SCl2
Chlorine and argon
It obviously doesn't make sense to talk about chlorine reacting with itself, and argon
doesn't react with chlorine.
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Sodium- Sodium reacts vigorously with water, melting,
giving hydrogen gas leaving strongly alkaline solution of
sodium hydroxide.
2Na + 2H2O → 2NaOH
Magnesium- Magnesium reacts very slowly with cold water
producing weakly alkaline solution with slightly soluble
Magnesium hydroxide.
Mg + 2H2O
→ Mg(OH)2
Hot Magnesium reacts with steam vigorously to form
Magnesium oxide and hydrogen gas.
Mg + 2H2O → MgO +H2
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Basic oxides and hydroxides: Na2O, NaOH,
MgO and Mg(OH)2
Sodium oxide:
Na2O (s) + H2O (l) → 2NaOH (aq)
Na+ (aq) + OH - (aq) + H+ (aq) → Na+ (aq) + H2O (l)
Magnesium oxide:
MgO (s) + 2H+ (aq) → Mg2+ (aq) + H2O (l)
Mg(OH)2 (s) + 2H+ (aq) → Mg2+ (aq) + 2H2O (l)
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Amphoteric oxide and hydroxide: Al2O3, Al(OH)3.
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Basic properties:
Aluminium oxide is very resistant to attack by acids and so the first
reaction is slow. Aluminium hydroxide reacts easily with dilute
acids.
Al2O3 (s) + 6H+ (aq) " 2Al3+ (aq)+3H2O(l)
Al(OH)3 (s) + 3H+ (aq) " Al3+ (aq) + 3H2O (l)
Acidic properties:
Alkaline attack on aluminium oxide is slow unless the alkali is hot
and concentrated or molten. The formula of the aluminate ion
shown is accepted in exams; you may also see [Al(OH)4]- or AlO2-,
though the latter only really comes from reaction with molten
alkali.
Al2O3(s) + 6OH – (aq) + 3H2O(l) " 2 [Al(OH)6]3 –(aq)
Aluminium hydroxide reacts readily with sodium hydroxide
solution:
Al(OH)3 (s) + 3OH – (aq) " [Al(OH)6]3 – (aq)
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Acidic oxides (acid anhydrides) SiO2, P4O10,
SO2, SO3, Cl2O.
Silicon dioxide, silica. Silica’s giant covalent lattice means it is very
resistant to attack; the alkali solution must be concentrated and needs
heating.
SiO2 (s) + 2 OH – (aq) " SiO3 2 – (aq) + H2O(l)
Phosphorus(V) oxide. The reaction is violent with cold water.
P4O10 (s) + 2H2O (l) " 4HPO3(aq)
On heating the solution further reaction occurs forming phosphoric(V) acid:
HPO3(aq) + H2O(l) " H3PO4(aq)
Sulphur dioxide:
SO2(g) + H2O (l) " H2SO3 (aq)
Sulphur trioxide: This reaction is violent, and is not used directly to make
sulphuric acid.
SO3(g) + H2O (l) " H2SO4 (aq)
Dichlorine oxide:
Cl2O (aq) + H2O(l) " 2 HClO (aq)
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The chlorides we'll be looking at are:
NaCl, MgCl2, AlCl3, SiCl4, PCl5, PCl3, S2Cl2
The structures
Sodium chloride and magnesium chloride consist of giant ionic
lattices at room temperature
Aluminium chloride and phosphorus(V) chloride are tricky! They
change their structure from ionic to covalent when the solid turns
to a liquid or vapour. Melting and boiling points
Sodium and magnesium chlorides are solids with high melting
and boiling points because of the large amount of heat which is
needed to break the strong ionic attractions.
The rest are liquids or low melting point solids. Leaving aside the
aluminium chloride and phosphorus(V) chloride cases where the
situation is quite complicated, the attractions in the others will be
much weaker intermolecular forces such as van der Waals
dispersion forces. These vary depending on the size and shape of
the molecule, but will always be far weaker than ionic bonds.
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
Oxidation number of elements in chlorides, reaction
with water
Chloride
Na
Mg
Al
Si
P
S
Oxidation
number
+1
+2
+3
+4
+5 or +3
+2 or +1
With water
Dissolves
Reacts & gives of HCL fumes
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