Further explanation on Period table elements

1. PRE-U CHEMISTRY
SEMESTER 2
CHAPTER 3 :
PERIODIC TABLE :
PERIODICITY

2. 3.0 Introduction to Inorganic Chemistry
Inorganic chemistry deals with the properties of all of the
elements in the periodic table. These elements range from
highly reactive metals, such as sodium, to noble metals, such
as gold. The nonmetals include solids, liquids, and gases, and
range from the aggressive oxidizing agent fluorine to
unreactive gases such as helium. Although this variety and
diversity are features of any study of inorganic chemistry,
there are underlying patterns and trends which enrich and
enhance our understanding of the discipline. These trends inenhance our understanding of the discipline. These trends in
reactivity, structure, and properties of the elements and their
compounds provide an insight into the landscape of the
periodic table and provide a foundation on which to build
understanding.
The periodic table provides an organizing principle that
coordinates and rationalizes the diverse physical and chemical
properties of the elements. Periodicity is the regular manner in
which the physical and chemical properties of the elements
vary with atomic number

3. PERIOD 2
Element Li Be B C N O F Ne
Proton number 3 4 5 6 7 8 9 10
Atomic radius
(nm)
0.152 0.111 0.086 0.077 0.073 0.062 0.032 0.029
Melting pointMelting point
(oC)
180 1287 2076 3500 -220 -218 -210 -249
1st ionisation
energy
(kJ/mol)
519 900 799 1090 1400 1310 1680 2080
Electronegativ
ity
0.98 1.57 2.04 2.55 3.04 3.44 3.90 --
Classification Metal
Metal
-loid
Non metal

4. PERIOD 3
Element Na Mg Al Si P S Cl Ar
Proton number 11 12 13 14 15 16 17 18
Atomic radius
(nm)
0.186 0.160 0.143 0.118 0.108 0.106 0.099 0.088
Melting point
(oC)
98 650 660 1423 44 120 -101 -189
1st ionisation
energy
(kJ/mol)
494 736 577 786 1060 1000 1260 1520
Electronegativ
ity
0.9 1.2 1.5 1.8 2.1 2.5 3.0 -
Classification Metal
Metal
-loid
Non metal

5. 3.1 Variation of physical properties of group and period
1. Atomic radius – half of the distance between the nuclei of
the two closest@ identical atom (or half of the closest internuclear
distance)Type Diagram Explanation
Covalent
radius (for
metalloid and
non-metal)
• In the case of covalent molecule,
atomic radius is also known as
covalent radius. Covalent radius
is half the distance between the
nuclei of 2 identical atoms
covalently bonded.
• Or simply, covalent radius is half
Atomic
nucleus
non-metal)
• Or simply, covalent radius is half
of the bond length between 2
covalently bonded identical atoms.
Metallic
Radius
(for metal)
• Metallic radius is define as half
the distance between the nuclei
of neighbouring metal ion in a
crystal lattice of a metal.
• Usually, the metallic radius is
greater than covalent radius.

6. The atomic size of an element is determined by 2 factors.
The screening effect of the inner shell electrons which makes
the atomic size larger. The screening effect is the result of the
mutual repulsion between the electrons in the inner shell with
those in the outer shell. Filled inner shells “shield” the outer
electrons more effectively than electrons in the same sub-shell
shield each other.
The nuclear charge which pulls all the electrons closer to
nucleus. As a result of the increasing nuclear charge, atomic
size becomes smaller.size becomes smaller.
Hence when 2 factors combine effective nuclear charge, Zeff
Zeff = Z (nuclear charge) – σ (screening constant)

7. The trend of atomic radius when gases down to group atomic
radius ……………………
Explanation : When going down to group, nuclear charge
increase as number of proton increase together with number of
electrons. However, as more electrons filling the shells, the
screening effect also increase. Consequently caused the effect
nuclear charge decrease and outer most shell electrons are not
hold tightly by the nucleus. For these reason, atomic radius
increase
The trend of atomic radius across the period. When across the
increase
The trend of atomic radius across the period. When across the
period 3, atomic radius ……………….
Explanation : When going across period, nuclear charge
increases as number of protons increase together with the number
of electrons. However, the screening effect remain almost constant
because electrons are filling in the same shell. This will caused the
effective nuclear charge increases gradually resulting the atomic
radius to decrease.
decrease

8. Atomicradiusdecrease
Atomic radius increaseAtomic radius increase

9. 2. Ionic radius
Ionic radius measures from the ion’s nucleus to the outermost
shell.
Diagram below shows the ionic radius for 2 cations from Period 4
and 3 anions from Period 3.

10. Ion Anion Cation
Diagram
Going
When going across P3- , S2- , Cl- , K+ , Ca2+ the ionic radius decrease
When going across these ions, the nuclear charge increase, since the number
Going
across
Period
of protons increase. However, all these ions are isoelectronic (have the same
number of electrons), hence the screening effect of these ions remain
constant. This will caused the effective nuclear charge to increase, which
result the ionic radius decrease when going across these ions.
Going
down
Group
When going down to any group, ionic radius decrease (e.g. :Group 1 Li+,
Na+, K+, Rb+, Cs+ )
This is due to, as nuclear charge increase, more electrons filling in the shell,
which caused the screening effect to increase gradually. As a result, the
effective nuclear charge decrease, hence caused the ionic radius to
decrease

11. When the atomic and ionic radius of an element were to compare,
student must know why does the atomic radius of an element is
greater/smaller than its ionic radius, by using the screening
effect and effective nuclear charge
Cation Anion
Trend
Atomic radius is larger than cation
radius
Atomic radius is smaller than anionic
radius
Explanation
Using Mg and Mg2+ as example ;
Electronic configuration of Mg is
1s22s22p63s2. When 2 electrons were
donated and form Mg2+ (1s22s22p63s2),
the effective nuclear charge increase as
the number of shell decrease, which will
decrease the screening effect.
Using P and P3- as example ;
Electronic configuration of P is
1s22s22p63s23p3. As P accept 3 electrons
and form P3- (1s22s22p63s23p6), the
effective nuclear charge decrease as the
number of shell increase, which will
decrease the effective nuclear charge.

12. 3. Melting point
Table below shows the melting point of elements across Period 3

13. Bonding and Forces Period Explanation
Metallic bonding - formed
when electrostatic forces is
formed between the
delocalised electrons and
the positive ion. When
electrons were delocalised
from a metal, it formed an
electron sea thus interacting
2
Elements : Lithium (Li) and Beryllium (Be)
Valence electrons of Li and Be are 2s1 and 2s2
respectively. Since Be delocalise more electrons
than Li, so melting point of Be is higher than Li
3
Elements : Sodium (Na) , Magnesium (Mg) and
Aluminum (Al)
Valence electrons of Na, Mg and Al are 3s1 , 3s2
electron sea thus interacting
with the positive ion formed
as a result of donating
electrons. Thus, the more
the electrons delocalised
by the metal, stronger the
electrostatic forces,
stronger the metallic bond
3
Valence electrons of Na, Mg and Al are 3s , 3s
and 3s23p1 respectively. Since Na, Mg and Al
delocalised 1, 2 and 3 electrons respectively, so
melting point increase Na < Mg < Al
Between
Period
Example : Between Be and Mg
Valence electrons of Be and Mg are 2s2 and 3s2
respectively, indicate they are from the same
Group. Since Be has smaller metallic radius than
Mg, hence greater electrostatic forces, so higher
the melting point.

14. Bonding and Forces Period Explanation
Gigantic structure - each
atom are strongly held by
using covalent bond
(depending on the number of
valence electrons that are able
2
Elements : Boron and Carbon
Valence electrons of B and C are 2s22p1 & 2s22p2
respectively, hence B form strong covalent with 3
other boron atoms(via sp2 hybridisation), while C
form strong covalent bond with 4 other carbon atoms
(via sp3 hybridisation). More energies required to
break more covalent bonds form between C, hence C
has a higher melting point than B
Element : Silicon
Valence electron of Si is 3s23p2. So, each Si can form
to form covalent bond) hence
forming a gigantic network
which are very stable and
required high temperature to
break the covalent bond
within the gigantic network.
3
Valence electron of Si is 3s23p2. So, each Si can form
strong covalent bond with 4 other Si atom (via sp3) to
form a gigantic covalent network, hence required high
temperature to break it.
Betwee
n
Period
Example : C and Si
Valence electrons of C and Si are 2s22p2 & 3s23p2
respectively. Both of them form sp3 hybridisation
between each atom. Since bond length between C-C is
shorter than Si-Si, hence stronger covalent bond is
form between C-C. That is why carbon has a higher
melting point than silicon

15. Bonding and Forces Period Explanation
Simple molecules - Non-
metal (except for C) tend
to form simple molecule
between them by using
covalent bond. These
molecules are hold weakly
2
Elements : Nitrogen, Oxygen, Fluorine, Neon.
Nitrogen, Oxygen, Fluorine exist as diatomic
molecule, as N2 O2 and F2, while Neon exist as
monoatomic Ne. Boiling point increase from
Ne<N2<O2<F2 as the molecular mass increased
in the order arranged which increased weak Van
Der Waals forces.
Phosphorous (P), Sulphur (S), Chlorine (Cl),
Argon (Ar)
Phosphorous exist as P , Sulphur exist as S ,molecules are hold weakly
by using weak Van Der
Waals forces between
them, hence it required
relatively low
temperature to break the
weak intermolecular forces
between them
3
Phosphorous exist as P4, Sulphur exist as S8,
Chlorine exist as Cl2, while Argon exist as
monoatomic Ar. Boiling point increase from Ar
< Cl2 < P4 < S8, as the molecular mass increase
in order arranged which increased weak Van Der
Waals forces
Between
Period

16. 3.1.4 First ionisation energy
The first ionisation energy is the minimum energy required to
remove 1 mole of electron from 1 mole atom at gaseous state to
form a unipositive ion. M (g) M+(g) + e
Three factors are involved in determining in ionisation energy
of an element :
The distance of valence electrons from the nucleus.
The magnitude of the nuclear charge.
The effectiveness of the shielding among the orbitals.The effectiveness of the shielding among the orbitals.
GENERALLY – The nuclear charge increases from sodium
to chlorine while the atomic size decreases. Hence, the
distance between the valence electrons and the nucleus is
getting shorter. In addition, the shielding or screening effect
remains almost constant across the period since electrons are
filled in the same shell . All these factors contribute to an
increase in ionisation energy across the period as valence
electrons become more difficult to be removed

18. Extra note
When going down to Group, Ionisation energy decrease.
This is due to, when going down to group, nuclear charge
increased with the number of electrons. As a result, more
shells are used to fill in the electrons. This will cause the
screening effect increase, which gradually increase the
atomic radius. Hence, the effective nuclear charge decrease,
causing the ionisation energy decreased.

19. When across period, the first ionisation energy generally
……………, since the nuclear charge across the period …………..
while the screening effect …………………..…………. as electrons
are filling in the same shell. As a result, the atomic radius
………….. , which cause the effective nuclear charge ……………..
thus ………….. the ionisation energy.
There are some anomalies of the trend of ionisation energy when
across period. For example, in Period 3, The anomaly occur
between ionisation energy of magnesium – aluminium and also
phosphorous – sulphur.
increase increase
remain almost constant
decrease increase
increase
phosphorous – sulphur.
Supposedly, the ionisation energy of magnesium is lower than
aluminium, since the atomic radius of magnesium is …………….
than aluminium. The orbital diagram of electron valence for Mg
and Al are as below
3 s 3 p 3 s 3 p
magnesium aluminium
higher

20. Since the …………. 3s orbital are more …………than a …………
filled orbital of 3p in aluminium, thus the energy required
………….. to draw out an electron from a single electron in the 3p
orbital.
For the anomaly occur among the 1st ionisation energy of
phosphorous and sulphur, it can be explained by using the
orbital diagram of phosphorous and sulphur
full-filled stable partially
is lesser
3s 3p 3s 3p
phosphorous sulphur
The ………………. 3p orbital in phosphorous are more stable
than …………………………. 3p orbital in silicon. Thus the energy
required to withdraw the electron from sulphur ……………….
than expected.
half-filled
partially – filled
is lesser

21. Element Na Mg Al Si P S Cl Ar
Electronegativity 0.9 1.2 1.5 1.8 2.1 2.5 3.0 -
1.4 Electronegativities & Electron Affinity
Na Mg Al Si P S Cl

22. 1.4 Electronegativities
Electronegativities is the relative strength of an atom to attract
electrons in a covalent bond which it is bonded.
Going across the third period, the increase in the nuclear
charge results in a greater attraction for the electrons in the
outermost shell. This increase tendency to attract electrons result
in an increase in electronegativity
1.5 Electron affinity
Electron affinity is the amount of energy being liberated when anElectron affinity is the amount of energy being liberated when an
atom receive one mole of electron in gaseous state.
F (g) + e- F- H = - ve kJ/mol
Unlike electronegativity (which has no unit), electron affinity
explained on how ‘easy’ an atom receive the electron and form
anion (mostly applied when forming lattice crystal)
Across period 3, the electron affinity increase, meaning the
tendency of the atom to receive an electron (Chlorine is the
easiest to form chloride ion)

24. 3.1.7 Predicting position of element using successive
ionisation energy
When 1st electron is ionised under the following expression :
A (g) → A+ (g) + e- ∆H1st IE = + a kJ mol-1 ;
the energy required is known as the 1st ionisation energy
It is possible for A+ to further ionised to form ion with greater
charge. When A+ is further ionised, the equation can be expressed
as : A+ (g) → A2+ (g) + e- ∆H2nd IE = + b kJ mol-1 and the energies
used is known as 2nd ionisation energy. It is expected that 2nd
ionisation energy is greater than 1st ionisation energy since theionisation energy is greater than 1st ionisation energy since the
effective nuclear charge of A+ (g) is greater than in A (g).
The A2+ can further be ionised when 3rd ionisation energies is
applied, where
A2+ (g) → A3+ (g) + e- ∆H3rd IE = + c kJ mol-1
For the energies used to remove each electron, it is known as
successive ionisation energies. So, it is possible to remove all
electrons in an atom when a massive amount of energies is
applied.

25. No. of
electron
removed
Ionisation
energy
(kJ/mol)
lg IE
No. of
electron
removed
Ionisation
energy
(kJ/mol)
lg IE
1st 738 2nd 1 4512.87 3.16
3.89 4.023rd 7 733 4th 10 541
5th 13 629 6th 17 995
7th 21 704 8th 25 657
9th 31 644 10th 35 463
11th 169 996 12th 189 371
3.89 4.02
4.13 4.26
4.34 4.41
4.50 4.55
5.23 5.28

26. 4.5
5.0
5.5
1 2 3 4 5 6 7 8 9 10 11 12
2.5
3.0
3.5
4.0

27. Note the following points of the graph of lg IE against no of
electron removed.
Each successive ionisation energies increased gradually,
indicates for each electron removed, the effective nuclear
charge also increased gradually.
The 2nd and 3rd ionisation energies difference significantly.
This is due to the 3rd electron is removed from an inner
shell. Therefore, the screening effect decreased significantly,
hence increase the effective nuclear charge greatly. So,
greater amount of energies were required to remove the 3rdgreater amount of energies were required to remove the 3rd
electron.
The same explanation occur between the 10th and 11th
ionisation energies, where there were a huge difference
between them, indicate that 11th electron were removed from
another inner shell

28. Group 2 :
Valence electron : ns2
Group : 15
Valence electron : ns2 np3

29. Group : 1
Valence electron : ns1
Group : 18
Valence electron : ns2 np6
Group : 13
Valence electron : ns2 np1
Group : 17 or 18
Valence electron : ns2 np5/6
IE 1 2 3 4 5 6 7 IE 1 2 3 4 5 6 7
∆H 459 1400 2717 7205 8720 10020 11400 ∆H 653 1925 3420 4860 6130 7670 9090

30. IE 1 2 3 4 5 6 7
∆H 362 1693 3102 4604 10350 11890 13700
Element R
Group : 14
Valence electron : ns2 np2
Element S
IE 1 2 3 4 5 6 7
∆H 259 1320 2890 4200 5492 9970 11020
Group : 15
Valence electron : ns2 np3

31. 1.3 Chemical Properties of Period 3
Element Na Mg Al Si P S Cl Ar
Proton number 11 12 13 14 15 16 17 18
Valance Electron 3s1 3s2 3s23p1 3s23p2 3s23p3 3s23p4 3s23p5 3s23p6
Ionic form Na+ Mg2+ Al3+ -- P3- S2- Cl- --
Bonding Metallic Bonding
Giant
covalent
Simple covalent
Mono-
atomic
Oxidising /
reducing agent
Reducing agent
Oxidising
agent

32. 3.3.1 Oxidising and reducing ability of Period 3 element.
Since the ionisation of sodium, magnesium and aluminium are
relatively low, they tend to release electron. In the other word,
they tend to be oxidised.
By the angle of standard reduction potential, Eo
red, sodium has
the highest tendency to be oxidise as the Eo value is the most
negative. Thus metal are strong reducing agent
Na+ (aq) + e- ↔ Na (s) Eo = - 2.71 V
Mg2+ (aq) + 2 e- ↔ Mg (s) Eo = - 2.38 V
Al3+ (aq) + 3 e- ↔ Al (s) Eo = - 1.67 VAl3+ (aq) + 3 e- ↔ Al (s) Eo = - 1.67 V
It is because of the high oxidising ability, it is used in the
extraction for some metal. Example
Extracting titanium metal : TiCl4 + 2 Mg 2 MgCl2 + Ti
Extracting chromium metal : Cr2O3 + 2 Al Al2O3 + 2 Cr
As for chlorine, since it has a high electron affinity, it has a
tendency to receive an electron. Thus, chlorine is preferably to be
reduced.
Cl2 (g) + 2e- ↔ 2 Cl- (aq) Eo = + 1.36 V

33. 1.3.1 Trend of oxide of Period 3
Element Na Mg Al Si P S Cl
Oxide of element
When burned with
oxygen
Na2O2 MgO Al2O3 SiO2 P4O10 SO3 Cl2O7
Bonding Ionic
Giant
Covalent
Simple covalentBonding Ionic
Covalent
Simple covalent
Acid-Base Basic Oxide
Ampho-
teric
Acidic Oxide

34. 1. In the laboratory, sodium and potassium are normally kept under
paraffin oil to avoid contact with air. This is because alkali metals are
extremely reactive. Sodium burns brilliantly in air (limited supply of
oxygen) to form sodium oxide, a white powder.
Reaction of sodium with oxygen : 2 Na (s) + O2 (g) Na2O2 (s)
When Na2O2 is further heated, it decomposed to form Na2O
2 Na2O2 (s) 2 Na2O (s) + O2 (g)
(a) When sodium oxide dissolves in water, a strong alkali, sodium
hydroxide is formed. Na2O2 (s) + H2O (l) 2 NaOH (aq) + H2O2 (aq)
Sodium hydroxide and potassium hydroxide have similar properties.
They are both prepared industrially through the electrolysis of sodium
chloride and potassium chloride solutions.
Sodium hydroxide is used in the manufacture of soap and many organic
and inorganic compounds whereas potassium hydroxide is used as an
electrolyte in some storage batteries
2. Even though magnesium is not as reactive as sodium, it still burns
brilliantly in air with a bright light to from magnesium oxide (white
powder).
Reaction of magnesium with oxygen : 2 Mg (s) + O2 (g) 2 MgO (s)
(a) Magnesium oxide is a strong base and will dissolve slowly in water
to form magnesium hydroxide, a white solid suspension used to treat
acid indigestion MgO (s) + H2O (l) Mg(OH)2 (aq)

35. 3.Aluminium is another reactive metal that, when exposed to air,
will react easily with oxygen to form a white oxide coating.
4 Al (s) + 3 O2 (g) 2 Al2O3 (s)
This layer of aluminium oxide coating causing the metal to be
insoluble in water. Due to its amphoteric porperties, it can
react with both acids and alkalis.
(a) With acids, it behaves as a base to produce salt and water only.
Al2O3 (s) + 6 HCl (aq) 2 AlCl3 (aq) + 3H2O (l)
(b) With alkalis, it behaves as an acid and a complexs salt is
produced.produced.
Al2O3 (s) + 2 NaOH (aq) + 3 H2O (1) 2 NaAl(OH)4 (aq)
4.Silicon, a metalloid, only reacts with oxygen slowly at very high
temperature. Silicon dioxide is formed in the reaction.
Si (s) + O2 (g) SiO2 (s)
(a)Due to its gigantic molecular structure, silicon dioxide does not
react with water, but still, it reacted with concentrated alkalis to
form silicate ion.
SiO2 (s) + 2OH- (aq) SiO3
2- (aq) + H2O (l)

36. 5.Phosphorus burns readily in air (oxygen) to form acidic oxides.
White phosphorus is a highly toxic substance and will burst into
flames spontaneously when exposed to oxygen to form
phosphorus pentoxide, P4O10. If a limited supply of oxygen
is used during burning, a lower form of oxide, phosphorus
trioxide, P4O6, is produced.
Phosphorous burned with excess oxygen :
P4 (s) + 5 O2 (g) P4O10 (s)
Phosphorous burned with limited oxygen :
P4 (s) + 3 O2 (g) P4O6 (s)
(a) Both oxides are acidic and will dissolve in water to form the
corresponding acids.
Phosphorous pentoxide :
P4O10 (s) + 6 H2O (l) 4 H3PO4 (aq) [Phosphoric acid]
Phosphorous trioxide :
P4O6 (s) + 6 H2O (l) 4 H3PO3 (aq) [Phosphorous acid]

37. 6.Sulphur can from two important oxides, sulphur dioxide, SO2,
and sulphur trioxide, SO3. Sulphur burns in air to form sulphur
dioxide.
When sulphur burn in air : S (s) + O2 (g) SO2 (g)
(a) Sulphur dioxide is a pungent, colourless and toxic gas. Being a
non-metallic gas, sulphur dioxide dissolves in water to form
sulphurous acid. When sulphur dioxide dissolve in water :
SO2 (g) + H2O (l) H2SO3 (aq)
(b) In excess oxygen, sulphur dioxide will slowly be oxidised to(b) In excess oxygen, sulphur dioxide will slowly be oxidised to
sulphur trioxide. The reaction can be enhanced with the presence
of a catalyst like platinum or vanadium (V) oxide. When sulphur
burn in excess air : 2 SO2 (g) + O2 (g) 2 SO3 (g)
This process in important in Contact Process in industries as
sulphuric acid is made in such way. When sulphur trioxide
dissolves in water to form sulphuric acid.
When sulphur trioxide dissolve in water :
SO3 (g) + H2O (l) H2SO4 (aq)

38. 7.Chlorine does not react with oxygen gas under any condition.
(a) The oxide of chlorine, Cl2O, is a yellow gas made up by passing
dry chlorine gas over fresh precipitated mercury (II) oxide at
400oC.
Equation : 2 HgO (s) + 2 Cl2 (g) HgO • HgCl2 (s) + Cl2O (g)
(b) another oxide of chlorine, Cl2O7, is prepared by adding chloric
(VII) acid to phosphorous (V) oxide (act as dehydrating agent)
cooled in ice salt. The chlorine (VII) oxide can be distilled off
from the mixture
Equation : 2 HClO4 (aq) Cl2O7 (l) + H2O (l)

39. Element
Oxide
formula
Reaction equation with oxygen
Melting
point
(oC)
Oxida
-tion
state
Ionic/
covalent
bond
Acidic /
basic oxide
Na 1275
Mg 2852
Al 2072
Si 1610
Na2O2 2 Na + O2 Na2O2 +1 ionic basic
MgO 2 Mg + O2 2 MgO +2 ionic basic
Al2O3 4 Al + 3 O2 2 Al2O3 +3 ionic
ampho
teric
SiO2 Si + O2 SiO2 +4 covalent acidic
P O P + 3 O P O +3 covalent acidic
P
24
580
S
-73
17
Cl
-20
45
P4O6 P4 + 3 O2 P4O6 +3 covalent acidic
P4O10 P4 + 5 O2 P4O10 +5 covalent acidic
SO2 S + O2 SO2 +4 covalent acidic
SO3 2 SO2 + O2 2 SO3 +6 covalent acidic
Cl2O
2 HgO (s) + 2 Cl2 (g)
HgO • HgCl2 (s) + Cl2O (g)
+1 covalent acidic
Cl2O7
2 HClO4 (aq)
Cl2O7 (l) + H2O (l)
+7 covalent acidic

40. 3.3.1 The melting point trend of Period 3 oxides.
1.Sodium oxide, magnesium oxide and aluminium oxide are ionic
oxide. So, when it is concerning ionic substance, the strength of
ionic bond is influenced by charge of both cation and anion, and
ionic radius between the oppositely charged ions.
2.Usually, cation with high charge and small radius and
anion with high charge small radius has a greater
electrostatic attraction forces between them, hence a higher
melting point
3.Since sodium ion (Na+) in sodium oxide has a smaller charge3.Since sodium ion (Na+) in sodium oxide has a smaller charge
and greater cationic radius compare to magnesium ion (Mg2+)
in magnesium oxide, so the melting point of sodium oxide is
expected to be lower than magnesium oxide.

41. 4. When it comes to aluminium oxide and magnesium oxide,
supposedly aluminium oxide has a higher melting point than
magnesium oxide (as aluminium has a smaller radius and higher
charge compare to magnesium) but magnesium oxide is observed
to have much higher melting point compare to aluminium oxide.
This is due to the charge density of aluminium is very high, that
it caused the aluminium oxide formed has high covalency
properties which greatly reduce the ionic strength of the
aluminium oxide. The oxide ion is highly polarised by aluminium
and reduce the electrostatic forces between the 2 ionsand reduce the electrostatic forces between the 2 ions
5.For silicon oxide, SiO2, it has a gigantic molecular structure. The
covalent bond between silicon and oxygen are strong thus
requiring a high energy to break the strong covalent bond. That’s
why the melting point of silicon oxide is high.
6.As for phosphorous oxide, sulphur oxide, chlorine oxide, they are
held by weak Van Der Waal forces. The weak Van Der Waals
forces increased as the molecular mass increase, so the trend of
the non-metal oxide is as follow
SO2 < Cl2O < SO3 < P4O6 < Cl2O7 < P4O10

42. MELTING POINT against Period 3 oxide
Na2O
MgO
Al2O3
SiO2
P O
Na Mg Al Si P S Cl
P4O10
SO3
Cl2O7

43. 1.3.0 Reaction of Period 3 element with water
Element Equation of reaction with water
Acidic / basic
properties of solution
Na
Mg
2 Na + 2 H2O 2 NaOH + H2 basic
Mg + 2 H2O Mg(OH)2 + H2 basic
Al
Si
P
S
Cl
Does not react with water --
Cl2 + H2O HClO + HCl acidic

44. 1. Alkali metal such as sodium and potassium are very
electropositive metal. It reacts vigorously with water to form
basic hydroxide solution and releases hydrogen gas. The
reactivity increases when goes down to Group 1.
Reaction of sodium with water :
2 Na (s) + 2 H2O (l) 2 NaOH (aq) + H2 (g)
2. Since Group 2 metal (earth alkali metal) is less reactive than
alkali metal (Group 1), so a certain condition must be obeyed in
order for Group 2 to react. For magnesium, it reacted slowly
with steam to form magnesium hydroxide and hydrogen gas
Reaction of magnesium with steam :Reaction of magnesium with steam :
Mg (s) + 2 H2O (g) Mg(OH)2 (aq) + H2 (g)
3 Aluminium is a Group 13 element. In nature, its principal ore is
bauxite, Al2O3.2H2O. As we move across the Periodic Table from
left to right in a given period, there is a gradual decrease in
metallic properties. So, although aluminium is regarded as
reactive metal, it is not as reactive as sodium or
magnesium. It does not react with water because it has a
protective layer (oxide) on its surface

45. 4. Silicon, phosphorous and sulphur does not react with water
under any condition. So nothing will be produced.
5. Chlorine, a halogen, is a reactive non-metal. The
magnitudes of reactivity and toxicity decrease down the
group from fluorine to iodine. Halogen dissolves partially in
water to form acids. Chlorine is used to purify water
and disinfect swimming pools. When chlorine dissolves
in water,it disproportionate to form hydrochloric acid,
HCl, and hypochlorous acid, HC1O, are formed. It is theHCl, and hypochlorous acid, HC1O, are formed. It is the
ClO- ions that kill the bacteria in the water.
Reaction of chlorine in water :
Cl (g) + H2O (1) HCl (aq) + HOCl (aq)