Chemistry Form 6 Sem 2 03

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

P

ERIOD

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 (o_C) 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

P

ERIOD

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 (o_C) 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

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, Z_eff

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.

A

to

m

ic

ra

d

iu

s

d

e

c

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a

s

e

Atomic radius increase

Atomic 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

Ion Anion Cation

Diagram

Going

When going across P3- _{, S}2- _{, Cl}- _{, K}+ _{, Ca}2+ _{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

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 1s2_2s2_2p6_3s2_{. When 2 electrons were} donated and form Mg2+ _(1s2_2s2_2p6_3s2_), 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

1s2_2s2_2p6_3s2_3p3_{. As P accept 3 electrons} and form P3- _(1s2_2s2_2p6_3s2_3p6_{), the}

effective nuclear charge decrease as the number of shell increase, which will decrease the effective nuclear charge.

3. Melting point

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 2s}2

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 _{, 3s}2

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 3s2_3p1 _{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 3s}2

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 2s2_2p1 _{& 2s}2_2p2 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 3s2_3p2_{. 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 3s2_3p2_{. 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 2s2_2p2 _{& 3s}2_3p2 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

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

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

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- ∆H}

1st 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) → A}2+ _{(g) + e- ∆H}

2nd 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) → A}3+ _{(g) + e- ∆H}

3rd 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

_2.87

₂nd _{1 451}

_3.16

3.89

4.02

3rd _{7 733 4}th _{10 541} 5th _{13 629 6}th _{17 995} 7th _{21 704 8}th _{25 657} 9th _{31 644 10}th _{35 463} 11th _{169 996 12}th _{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}

Group 2 :

Valence electron : ns2 Group : 15

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

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 _3s2_3p1 _3s2_3p2 _3s2_3p3 _3s2_3p4 _3s2_3p5 _3s2_3p6

Ionic form Na+ _Mg2+ _Al3+ _{-- P}3- _S2- _Cl-

--Bonding Metallic Bonding Giant

covalent Simple covalent

Mono-atomic

Oxidising /

reducing agent Reducing 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, 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) E}o _{= - 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.

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 covalent

Bonding Ionic

Covalent Simple covalent

Acid-Base Basic 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) + O_{2 (g)  Na2}O₂ (s) When Na₂O₂ is further heated, it decomposed to form Na₂O

2 Na₂O_{2 (s)  2 Na2}O (s) + O₂ (g)

(a) When sodium oxide dissolves in water, a strong alkali, sodium hydroxide is formed. Na₂O₂ (s) + H_{2O (l)  2 NaOH (aq) + H2}O₂ (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) + O_{2 (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) + H_{2O (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 O_{2 (g)  2 Al2}O₃ (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. Al₂O_{3 (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 H_{2O (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) + O_{2 (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)  SiO}

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 O_{2 (g)  P4}O₁₀ (s)

Phosphorous burned with limited oxygen :

P₄ (s) + 3 O_{2 (g)  P4}O₆ (s)

(a) Both oxides are acidic and will dissolve in water to form the corresponding acids.

Phosphorous pentoxide :

P₄O₁₀ (s) + 6 H_{2O (l)  4 H3}PO₄ (aq) [Phosphoric acid] Phosphorous trioxide :

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) + O_{2 (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) + H_{2O (l)  H2}SO₃ (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) + O_{2 (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 :

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

400o_C.

Equation : 2 HgO (s) + 2 Cl_{2 (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

Element Oxide

formula Reaction equation with oxygen

Melting point (o_C) Oxida -tion state Ionic/ covalent bond Acidic / basic oxide Na 1275 Mg 2852 Al 2072 Si 1610 Na₂O₂ 2 Na + O_{2  Na2}O₂ +1 ionic basic MgO 2 Mg + O_{2  2 MgO} +2 ionic basic Al₂O₃ 4 Al + 3 O_{2  2 Al2}O₃ +3 ionic ampho

teric SiO₂ Si + O_{2  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 O_{2  P4}O₆ +3 covalent acidic P₄O₁₀ P₄ + 5 O_{2  P4}O₁₀ +5 covalent acidic SO₂ S + O_{2  SO2} +4 covalent acidic SO₃ 2 SO₂ + O_{2  2 SO3} +6 covalent acidic Cl₂O 2 HgO (s) + 2 Cl2 (g) 

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

Cl₂O₇ 2 HClO4 (aq) 

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

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 water Acidic / basic

properties of solution Na Mg

2 Na + 2 H

₂

_{O  2 NaOH + H}

₂

basic

Mg + 2 H

₂

_{O  Mg(OH)}

₂

+ H

₂

basic

Al Si P S Cl

Does not react with water

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 H_{2O (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 H_{2O (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

: