PRE-U CHEMISTRY

SEMESTER 2

CHAPTER 3 :

PERIODIC TABLE :

PERIODICITY

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 in enhance 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

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 point Melting point

(oC)180 1287 2076 3500 -220 -218 -210 -249

1st ionisation

energy

(kJ/mol)519 900 799 1090 1400 1310 1680 2080

Electronegativ

ity0.98 1.57 2.04 2.55 3.04 3.44 3.90 --

Classification MetalMetal

-loidNon metal

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

ity0.9 1.2 1.5 1.8 2.1 2.5 3.0 -

Classification MetalMetal

-loidNon metal

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.

� 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)

� 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

Atomic ra

dius decrease

Atomic radius increaseAtomic radius increase

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.

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

� 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

TrendAtomic 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.

3. Melting point

� Table below shows the melting point of elements across Period 3

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 , 3s2electron 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

3Valence 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.

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

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

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

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.

� 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

� 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

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

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 an Electron 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)

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 the ionisation 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.

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

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

� 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 3rd greater 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

Group 2 :

Valence electron : ns2Group : 15

Valence electron : ns2 np3

Group : 1

Valence electron : ns1

Group : 18

Valence electron : ns2 np6

Group : 13

Valence electron : ns2 np1Group : 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

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

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 BondingGiant

covalentSimple covalent

Mono-

atomic

Oxidising /

reducing agentReducing agent

Oxidising

agent

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, Eored, 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 V� Al3+ (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

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 OxideAmpho-

tericAcidic Oxide

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)

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)

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]

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)

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)

ElementOxide

formulaReaction 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 ionicampho

teric

SiO2 Si + O2 � SiO2 +4 covalent acidic

P O P + 3 O � P O +3 covalent acidicP

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

Cl2O2 HgO (s) + 2 Cl2 (g) �

HgO • HgCl2 (s) + Cl2O (g)+1 covalent acidic

Cl2O72 HClO4 (aq) �

Cl2O7 (l) + H2O (l)+7 covalent acidic

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 charge 3.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.

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 ions and 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

MELTING POINT against Period 3 oxide

Na2O

MgO

Al2O3

SiO2

P O

Na Mg Al Si P S Cl

P4O10

SO3

Cl2O7

1.3.0 Reaction of Period 3 element with water

Element Equation of reaction with waterAcidic / basic

properties of solution

Na

Mg

2 Na + 2 H2O � 2 NaOH + H2basic

Mg + 2 H2O � Mg(OH)2 + H2 basic

Al

Si

P

S

Cl

Does not react with water --

Cl2 + H2O � HClO + HCl acidic

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

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 the HCl, 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)