In this article, we shall study hydroxides of third row elements.

Different types of hydroxides of third-row elements are classified according to their mode of dissociation. According to Arrhenius’s theory, the acid is the substance which gives H⁺ ions in an aqueous medium, while the base is the substance which gives OH^– ions in an aqueous medium.

Let us use M-O-H be the general formula to represent a hydroxy compound of third row elements. The mode of ionization decides the nature of hydroxide whether it is acidic or basic.

Types of Hydroxy Compounds:

Basic hydroxy compound:

The hydroxy compounds which give OH^– ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq.)    →    M⁺   +   OH^–

This is possible when the element has very low ionization potential. Valence electrons are loosely held by the atom. Due to very low ionization enthalpy and electronegativity metal atom cannot hold valence electrons. Electron pair between M and O is pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong. Thus the bond between M and OH breaks

e.g. hydroxide NaOH of sodium and hydroxide Mg(OH)₂ of magnesium are basic hydroxy compounds.

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg(OH) _{2  (aq.)} →    Mg ^{2 +}  +  2 OH ^–

Both hydroxides give OH^– ions in aqueous medium.

Acidic Hydroxy Compounds (Oxyacids): 

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq.)    →   MO ^–   +   H ⁺

This is possible when the element has greater ionization potential. Valence electrons are strongly held by atom. Due to higher ionization enthalpy and electronegativity, metal atom hold valency electrons strongly. Electron pair between M and O is pulled more towards more electronegative element M. As a result of which bond M-O-H behaves as a base.

M-O bond becomes strong while the O-H bond becomes weak. Thus the bond between MO and H breaks, which results in the production of H + ions in aqueous medium.

The hydroxy compounds which give H ⁺ ions in an aqueous medium are called acidic hydroxy compounds or oxyacids. If M-O-H is assumed oxyacid then O-H bond break in an aqueous medium and it will give H ⁺ ions.

e.g. Si(OH)₄ of Silicon, P (OH)₃ and PO (OH)₃ of Phosphorous, SO (OH)₂ and SO₂(OH)₂ of sulphur , ClOH, ClO(OH) , ClO₂(OH) and ClO₃(OH) of chlorine are acidic hydroxy compounds.

Amphoteric Hydroxy Compound:

Hydroxy compound which acts as acid as well as base and can neutralize the acid, as well as base producing salt and water, is called amphoteric hydroxide.

Al(OH)₃ of aluminium is an amphoteric oxide. It neutralizes acid as well as base producing a salt and water.

Al (OH)₃ (as a base) +   3HCl    →  AlCl ₃  + 3 H₂O

Al(OH)₃ (as an acid) + NaOH  → NaAlO₂   (Sodium meta-aluminate) +2 H₂O

Thus the nature of hydroxides or oxyacids is mainly governed by the ionization potential of elements.

Trends in acid -base behaviour of hydroxy compounds:

It is seen that as we move from Na to Cl along the third row, the basic character of hydroxy compounds gradually decreases while acidic character gradually increases.

The trend in acid-base behaviour of hydroxy compounds of the third row can be summarized as follows.

Explanation:

The nature of the hydroxy compound mainly depends upon the ionization potential of elements. The trend is so because along third-row ionization potential increase, electronegative character increases, atomic size decreases. Along the third row difference in electronegativity of the element M and that of Oxygen decreases.

If ionization potential of elements is low, then such hydroxy compound give OH ^– ions in an aqueous medium and hence is basic in nature.

M — O — H _(aq)        →   M ⁺  +  OH ^–

NaOH and Mg(OH)₂ are basic. Na and Mg have low ionization potential. Na – O and Mg – O bonds are weaker than O-H bond. And thus the bond between M and O breaks to give OH^– ions.

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg(OH) _{2  (aq.)} →    Mg ^{2 +}  +  2 OH ^–

If ionization potential of an element is greater, then such hydroxy compound gives H ⁺ ions in an aqueous medium and hence is acidic in nature.

M — O — H _(aq)        →   M O^–  +  H ⁺

Si(OH)₄ of Silicon, P(OH)₃ and PO(OH₃ of Phosphorous, SO(OH)₂ and SO₂(OH)₂ of sulphur, ClOH, ClO(OH), ClO₂(OH) and ClO₃(OH) of chlorine are acidic hydroxy compounds.

Al(OH)₃ is amphoteric. It neutralizes acid as well as base producing salt and water. Hence it exhibits both the properties hence it is an amphoteric oxide.

Scientific Reasons:

Hydroxy compound of Sodium NaOH is a strong base.

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq)        →   M ⁺  +  OH ^–

This is possible when the element has very low ionization potential. Valency electrons are loosely held by atom. Due to very low ionization potential and electronegativity, metal atom cannot hold valency electrons. Electron pair between M and O is pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong.

Sodium is the strongest electropositive element. Na has lower ionization potential and lower electronegativity. Na -O bond breaks more readily in an aqueous medium and Sodium hydroxide ionizes as

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg (OH)₂ is weakly basic than NaOH.

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq)        →   M ⁺  +  OH ^–

This is possible when the element has very low ionisation potential. Valency electrons are loosely held by atom. Due to very low ionisation potential and electronegatively metal atom cannot hold valency electrons. Electron pair between M and O is Pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong.

Na has lower ionisation potential and lower electronegativity than that of Mg. So Na – O bond is relatively weak than Mg – O bond.  Na -O bond breaks more readily than the Mg-O bond in an aqueous medium. Thus Mg(OH)₂ is weakly basic than NaOH.

Aluminium hydroxide is amphoteric compound.

Hydroxy compound which acts as acid as well as base and can neutralise acid, as well as base producing salt and water, is called amphoteric hydroxide.

Al(OH) ₃ of aluminium is an amphoteric oxide. It neutralises acid as well as base producing a salt and water.

Al (OH)₃ (as a base) +   3HCl    →  AlCl ₃  + 3 H₂O

Al(OH)₃ (as an acid) + NaOH → NaAlO₂   (Sodium meta-aluminate) + 2 H₂O

It acts as a base when treated with strong acid. It acts as an acid when treated with a strong base. Due to its dual character, Al(OH)₃ is amphoteric in nature.

Atomic size of Aluminium is smaller than Sodium and Magnesium and larger than Silicon, Phosphorous, Sulphur and Chlorine. The ionisation potential of Aluminium is larger than Sodium and Magnesium and lesser than Silicon, Phosphorous, Sulphur and Chlorine. Electronegativity of Aluminium is larger than Sodium and Magnesium and lesser than Silicon, Phosphorous, Sulphur and Chlorine. Thus in aluminium hydroxide, both Al – O and O – H bonds have equal strength. Hence the fission of bond depends on the attacking reagent. Hence Aluminium hydroxide is amphoteric compound.

Orthosilicic acid Si(OH)₄ is very weak acid. 

Electronegativity of Silicon is 1.8 units. The ionisation potential of Silicon is higher than that of Aluminium.

Si – O is covalent bond with the ionic character. Hence Si – O bond is a strong bond. Oxygen atom pulls the shared electron in O – H bond towards itself and on fission produces H⁺ ions. However, O-H bond is not broken in water.

The orthosilicic acid reacts with strong alkali on heating.

H₄SiO₄  +  2 NaOH  →  Na₂Si ₃   +  3 H₂O

Hydroxy compounds of Phosphorous, Sulphur And Chlorine are acidic. 

Phosphorous, Sulphur And Chlorine are non-metals with smaller atomic size, high nuclear charge, High ionisation potential and high electronegativity. These elements have very little or practically no tendency to give an electron to Oxygen.

In the structure of M – O – H Oxygen, therefore, tries to pull electron pair between O – H towards itself, This releases H⁺ ions in aqueous solution.

Structures of Hydroxy Compounds of Phosphorous or Oxyacids of Phosphorous:

Phosphorous acid (P(OH)₃ OR H₃PO₃) and  Phosphoric  acid (PO(OH)₃ OR (H₃PO₄)

H₃PO₃  +  2 NaOH  →  Na₂HPO₃  +  2 H₂O

H₃PO₄  +  3 NaOH  →  Na₃PO₄  +  3 H₂O

Phosphoric acid has an un-hydrogenated oxygen atom so O-H bond breaks readily. Hence phosphorous acid is stronger than phosphorous acid.

Structures of Hydroxy compounds of Sulphur or Oxyacids of Sulphur:

Sulphurous acid SO(OH)₂ OR (H₂SO₃) and Sulphuric  acid SO₂(OH)₂ OR (H₂SO₄)

H₂SO₃  +  2 NaOH  → Na₂SO₃  +  2 H₂O

H₂SO₄  +  2 NaOH  → Na₂SO₄  +  2 H₂O

These oxyacids of sulphur are strongly acidic due to greater IP of sulphur.  Sulphuric acid has two un-hydrogenated oxygen atoms while in sulphurous acid there has one un-hydrogenated oxygen atom hence, sulphuric acid is stronger than sulphurous acid.

” Greater the Oxidation no. More the Acidic Nature”. In H₂SO₄ the oxidation no. of the central atom Sulphur is +6 and that in H₂SO₃ is +4. The oxidation number of sulphur is greater in sulphuric acid than the sulphurous acid. Hence sulphuric acid is stronger than sulphurous acid.

Hydroxy compounds of Chlorine or Oxyacids of Chlorine:

Hypochlorous acid Cl(OH) OR (HOCl), Chloric acid ClO₂(OH) OR (HClO₃), Perchloric  acid ClO₃(OH) OR (HClO₄)

HOCl   +  NaOH  →   NaOCl    +   H₂O

HClO₄  +  NaOH  →  NaClO₄  +  —– H₂O

These are oxyacids of chlorine which are very strongly acidic. Chlorine has greater I.P and electronegativity. Due to the presence of 3 un-hydrogenated oxygen atoms attached to Chlorine atom,  HClO₄ is strongest amongst these oxyacids of chlorine.

The strengths are in the order are ClO₃(OH) > ClO₂ (OH)   > ClO(OH)  >  ClOH.

Science > Chemistry > Third Row Elements > Hydroxides of Third Row Elements

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A binary compound of an element with oxygen, in which the oxygen atom is electronegative is called an oxide. e.g. MgO, Al₂O_3, etc. The oxide in which Oxygen exhibits the normal oxidation state of -2 is called normal oxide. In this article, we shall study the oxides of third-row elements.

Classification of Oxides:

Depending upon the chemical behaviour, oxides of third-row elements are Classified into three types, namely i) acidic oxides  Na₂O, MgO ii) basic oxides SiO₂, SO₃, Cl₂O_7, P₂O₅  and iii) amphoteric oxides Al₂O₃.

Basic Oxides:

The oxides which react with water and produce alkali and can neutralise acids forming salt and water are called basic oxides.

Na₂O of sodium and MgO of magnesium of the third row are basic oxides. Because both oxides produce alkali when treated with water and can neutralise acid producing salt and water.

Na₂O    +     H₂O        →    2 NaOH (strong base)

Na₂O    +   2 HCl        →   2 NaCl    +   H2O

MgO    +   2 H₂O        →   Mg (OH)₂ (base)

MgO    +   2 HCl        →   MgCl₂    +   H₂O

Na and Mg have bigger atomic size. 1 and 2 valence electrons respectively and very low ionisation potential value.

Amphoteric Oxide:

The oxide which acts as an acid as well as base and neutralizes acid as well as base to give salt and water is called an amphoteric oxide.

Al₂O₃ of aluminium is an amphoteric oxide. It neutralizes acid like HCl as well as a base like NaOH.

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    2 NaOH  →   2 NaAlO₂  (sodium aluminate)   +      H₂O

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    6 HCl  → 2 AlCl₃    +   3 H₂O

Aluminium has higher I.P. than sodium and its electronegativity is greater than Mg and Na metals. So Al-O bond in  Al₂O₃  shows amphoteric in nature.

Acidic Oxides:

The oxides of electronegative elements which give acid when treated with water and can neutralise bases to produce salt and water called acidic oxides.

Oxides of Phosphorous, Sulphur, Chlorine are acidic.

SiO₂   +  2 NaOH →   Na₂SiO₃   + H₂O

P₄O₁₀ +   6 H₂O  → 4 H₃PO₄

SO₃   +  H₂O    →     H₂SO₄

SO₃    +  2 Na OH  → Na₂SO₄  +  H₂O

Cl₂O₇   + H₂O →  2 HClO₄ (perchloric acid)

Cl₂O₇  + 2 Na OH  →  2 NaClO₄ +   H₂O

Si, P ,S, and Cl have greater ionisation potential , electronegativity , tendency to attract electrons.  They have smaller atomic size and more valency electrons.

Trend in Acidic and Basic Character:

Trend:

Generally, the oxides of metals are basic in nature while that of non-metals are acidic. As we move from Na to Cl along the third period, the acidic character of elements goes on gradually increasing while basic character goes on decreasing.

Reasons:

• In third row elements, those elements who have low ionisation potential, bigger atomic size, less electronegativity, less number of valence electrons form basic oxides as their oxides give alkalies when treated with water.
• On the other hand, those elements who have greater ionisation potential, smaller atomic size, greater electronegativity, more valency electrons form acidic oxides, Such oxides give acids when treated with water.
• The trend is so because from Na to Cl while going along the third period, ionisation potential gradually increases, atomic size decreases, the number of valence electrons increases.

Examples:

Na₂O and MgO of Na and Mg respectively are highly basic in nature. They produce alkalies when treated with water.  They neutralise acids. Na and Mg have very low ionisation potential amongst third-row elements.

Na₂O    +     H₂O        →    2 NaOH (strong base)

Na₂O    +   2 HCl        →   2 NaCl    +   H2O

MgO    +   2 H₂O        →   Mg (OH)₂ (base)

MgO    +   2 HCl        →   MgCl₂    +   H₂O

The oxides of electronegative elements such as Si, P, S. and Cl are acidic in nature.  Their oxides give acids when treated with water can neutralise base producing salt and water. Si, P, S and Cl elements have greater ionisation potential, more valency electrons, greater electronegativity.

SiO₂   +  2 NaOH →   Na₂SiO₃   + H₂O

P₄O₁₀ +   6 H₂O  → 4 H₃PO₄

SO₃   +  H₂O    →     H₂SO₄

SO₃    +  2 Na OH  → Na₂SO₄  +  H₂O

Cl₂O₇   + H₂O →  2 HClO₄ (perchloric acid)

Cl₂O₇  + 2 Na OH  →  2 NaClO₄ +   H₂O

Al₂O₂ of aluminium is amphoteric because it neutralises acid as well as base producing salt and water. Due to its dual nature, it is called amphoteric.

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    2 NaOH  →   2 NaAlO₂  (sodium aluminate)   +      H₂O

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    6 HCl  → 2 AlCl₃    +   3 H₂O

The oxides of third-row elements are summarised in the following table.

Scientific Reasons:

Sodium oxide is more basic than Magnesium oxide. 

A binary compound of an element with oxygen, in which Oxygen atom is electronegative is called an oxide. The oxides which react with water and produce alkali and can neutralise acids forming salt and water are called basic oxides.

Na₂O, MgO are basic oxides. These oxides produce alkali when treated with water and can neutralize acid-producing salt and water. Sodium is a more electropositive element than Magnesium. Sodium has lower ionization potential and lower electronegativity than Magnesium.

Due to the above reasons Sodium loses its electron very easily to oxygen than Magnesium and acts as more basic than Magnesium.

Science > Chemistry > Third Row Elements > Oxides of the Third Row Elements

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in this article, we shall study the crystal structure of molecular solids of third perid of periodic table.

Molecular solid:

The substance in which lattice points are molecules which are held together by means of weak physical forces (van der Waal’s forces) are called molecular solid. Phosphorous, sulphur, chlorine, and argon are molecular solids because lattice points are molecules.

They have greater ionization potential and vacant valency orbitals are not available.

Characteristics of Molecular Solids:

• In the crystal structure of these elements, the units occupying lattice points are molecules
• They have a greater ionization potential
• The molecules are attached to each other by weak van der Wall’s forces of attraction.
• In these solids, the atoms are joined together within the molecule by strong covalent bonds.
• In these solids, vacant valency orbitals are not available. All the valence orbitals are used for intra-molecular strong covalent bonding.

Scientific Reasons:

Phosphorous, Sulphur, Chlorine, and Argon as solid are soft easily compressible and distorted. They are volatile their boiling points and melting points are very low.

In the crystal structure of these elements, the units occupying lattice points are molecules. They have a greater ionization potential. The molecules are attached to each other by weak van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds. Thus they are soft easily compressible and distorted.

Similarly less energy is required to separate the molecules from each other in a molecular crystal. Hence they are volatile and possess low boiling points and melting points.

Phosphorous and Sulphur are solids at room temperature while Chlorine and Argon are gases at room temperature.

In the crystal structure of these elements, the units occupying lattice points are molecules. They have a greater ionization potential. The molecules are attached to each other by weak van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

Sulphur is Octaatomic, Phosphorous is tetra atomic, Chlorine is diatomic, Argon is monoatomic. The size of the molecule decreases in the order S₈; P₄ > Cl₂ > Ar. van der Wall’s forces of attraction decrease in the same order. They are stronger in phosphorous and sulphur while negligible in chlorine and argon. Thus Chlorine and Argon have very low boiling and melting points. Hence Phosphorous and Sulphur are solids at room temperature while Chlorine and Argon are gases at room temperature.

Structure of Molecular Solids:

Structure of Phosphorous:

Phosphorous is molecular solid because lattice points are P₄ molecules which are held together by means of van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Phosphorous is 15. Electronic configuration of phosphorous is, 1s2, 2s22p6, 3s2 , 3px1 3py1 3pz1. Phosphorous has greater ionization potential. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Phosphorous atoms are sp³ hybridized.

One phosphorous atom forms three covalent bonds with three other phosphorous atoms while one sp³ hybrid orbital contained paired electrons (lone pair) remain non-bonded. Thus tetrahedral P₄ molecule is formed.  The P-P-P bond angle is 60°. The phosphorous molecule consists of four phosphorous atoms in it.

In white phosphorous, these tetrahedral molecules are joined together by means of weak van der Waal’s forces of attraction due to smaller molecular size.  So white phosphorous has a low melting point.

Red phosphorous is a polymeric form. In red phosphorous one P-P bond of P₄ unit gets ruptured and freed bonds link up to form a chain. Thus in red phosphorous, the tetrahedral molecules are joined to one another by means of strong covalent bonds, forming a chain like structure. Hence red phosphorous exists in polymeric form.

In white phosphorous, the molecules are joined together by weak van der Wall’s forces while in red phosphorous molecules are bound by strong covalent bonds. Therefore more energy is required for breaking bonds between molecules of red phosphorous than that in white phosphorous. Hence red phosphorous shows a high melting point and less reactivity than white phosphorous.

Structure of Sulphur:

Sulphur is molecular solid because lattice points are S₈ molecules which are held together by means of van der Waal’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Sulphur is 16. Electronic configuration of phosphorous is 1s2, 2s22p6, 3s2 , 3px2 3py1 3pz1. Sulphur has greater ionization potential. Its two half-filled orbitals are used for intermolecular bonding. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each Sulphur molecule consists of eight sulphur atoms in it.

The sulphur molecule has a puckered ring structure or crown structure. Sulphur has two half-filled 3p orbitals in the valency shell. Within each S₈ molecule, each s atom is linked to two adjacent S atoms by the covalent single bond. Thus if seen from the top it forms a ring-like structure with four sulphur atoms in one plane and alternate other four in a parallel plane. Each Sulphur atom has a lone pair of electrons. The S-S-S bond angle is 107.8 ⁰ and S-S bond length is 2.04 ^o

Structure of Chlorine:

Chlorine is molecular solid because lattice points are Cl₂ molecules which are held together by means of van der Waal’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Chlorine is 17. Electronic configuration of phosphorous is 1s², 2s²2p⁶, 3s² , 3p_x² 3p_y² 3p_z¹. Chlorine has greater ionization potential. Hence it forms covalent bonding. Its one half-filled orbital is used for intermolecular bonding. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each chlorine molecule consists of two Chlorine atoms in it bonded by a covalent bond.

Chlorine in solid state consists of layers of chlorine molecules which are held by weak Wander Wall’s forces. Hence Chlorine is gas at room temperature.

Structure of Argon:

The substance in which lattice points are molecules which are held together by means of weak physical forces (vander waal’s forces) are called molecular solids.

Argon is a molecular solid or atomic solid because lattice points are Ar atoms (molecules) which are held together by means of van der Waal’s forces of attraction. The atomic number of Argon is 18. Electronic configuration of Argon is 1s2, 2s22p6, 3s2 , 3px2 3py2 3pz1. Thus its octet is complete. It has no unpaired electrons. Argon has greater ionization potential. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each Argon molecule consists of one Argon atom.

Argon crystal in its solid-state consists of a continuous pattern of atoms giving rise to a face centred closed pack cubic lattice like aluminium. But the difference is that in the case of aluminium lattice points are aluminium ions while in case of argon lattice points are Argon atoms. Melting and boiling points of argon are very low.

Science > Chemistry > Third Row Elements > Molecular Solids of the Third Row

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In this article, we shall study properties metallic solids of third row.

The metallic bond is defined as the force of attraction that binds metal cations to a number of mobile or delocalized electrons within its sphere of influence which holds the metallic cations together in a definite pattern. On the basis of metallic bonding and type of crystal structure, the general properties of metals like high melting point, thermal conductivity, electrical conductivity, malleability, and ductility can be explained.

Characteristics of Metallic Solids:

• They possess high melting and boiling points
• Thy show high thermal conductivity
• They show high electrical conductivity
• They are malleable and ductile

Reasons for High Melting and Boiling Points of Metallic Solids:

• Metals possess a very strong metallic bond.
• In order to break a strong metallic bond more heat is required.
• Due to strong metallic bonding and closely packed crystal lattices, metals have higher melting and boiling points.
• Metallic bond strength depends on the quantity of charge carried by metal ions in the metallic bonding and hence the melting and boiling points of Na, Mg and Al are in the order. Al  >  Mg  > Na.
• Among Na, Mg, and Al, aluminium has greater melting and boiling point due to a very strong metallic bond and closely packed crystal lattice.

Reasons for High Thermal Conductivity of Metallic Solids:

In metals, cations are surrounded by many mobile electrons. The loosely bound delocalized electrons can move freely in a crystal lattice. When one end of the metal is heated, delocalized electrons acquires the heat energy, their Kinetic energy increases. The mobility of these delocalized electrons increases. They convey heat moving rapidly throughout the crystal lattice from the hot end to the cooler end. During this process, they collide with adjacent electrons and thus transfer part of their energy to adjacent electrons. Thus heat is transferred from one end to other due to mobile electrons and collisions.

Thus due to the presence of free, mobile electrons, metals have high thermal conductivity. Thermal conductivity depends on the number of mobile electrons. Thus the thermal conductivity increases from Na to Al as the number of valence electrons increases. Thermal conductivity increases in the order Na < Mg < Al

Reasons for High Electrical Conductivity of Metallic Solids:

In metals, cations are surrounded by many mobile electrons. The loosely bound delocalised electrons can move freely in a crystal lattice. The delocalised electrons in metals account for the electrical conductivity.

When an electric field i.e. potential difference is applied across the two ends of a metallic wire, the free electrons are displaced from the negative end to the positive end and conducts electricity. Thus high electrical conductivity of metals is due to the presence of mobile valence electrons.

Electrical conductivity depends on the number of mobile electrons. Thus the electrical conductivity increases from Na to Al as the number of valence electrons increases. Electrical conductivity increases in the order Na < Mg < Al

Malleability and Ductility:

The property of metals by which they can be flattened into thin shits by hammering is known as malleability. The property of metals by which they can be drawn out into wires by stretching is known as ductility. Malleability takes places under compression and ductility takes place under tension.

In metals, cations are surrounded by many mobile electrons. The loosely bound delocalised electrons can move freely in a crystal lattice. The metallic bond in the metal crystal is non-directional and not rigid because each metal ion is surrounded by many mobile electrons. The mutual repulsion between metal ions is neutralized by the sea of electrons that moves around them.

When stress is applied to the metal surface, one layer of metal ions slides over the other. This relative movement of layers with respect to each other is called slip. After slipping the ions settle into new positions keeping the distance between the metal ions same. At the same time, the mobile electrons also move along the metal ions. The crystal structure of the metal, therefore, does not change but the shape of the metal changes.

Thus a metal can be flattened into sheets without breaking of any strong metallic bonds and any change in its internal structure. The same thing happens when metal is drawn into wire.

Malleability and ductility of metal depend upon the crystal structure. Amongst Na, Mg, and Al, aluminium is more malleable and ductile because aluminium has more closely packed face centred cubic structure in which empty space is just about 26% of the available space in the unit cell and has a strong metallic bond. The order of malleability or ductility is Na < Mg < Al

Metallic lustre:

The clean surface of nearly all metals have shinning surfaces and reflects light giving the metals a silvery colour. i.e. lustre. In metals, cations are surrounded by many mobile electrons. The loosely bound delocalized electrons can move freely in a crystal lattice. Metallic lustre is due to the presence of free mobile electrons in the crystal lattice of metals.

When a beam of light falls on the surface of the metal, the mobile electrons present in the surface absorb the photons of incident light and get excited to the higher energy state. But this absorbed light energy is immediately re-emitted in the form of electromagnetic radiations of same frequencies by returning to original energy level. Consequently, due to this momentary exchange of light energy, the metal surface exhibit shinning appearance.

Na, Mg, and Al of the third row show bright lustre when freshly cut, because being chemically active, the surface of these metals gets coated with oxide, hydroxide, carbonate etc. formed due to their interaction with the atmospheric gases.

Electrical Conductivity of Silicon:

Non-Conductivity at absolute zero:

The atomic number of Silicon is 14. Electronic configuration of silicon in its ground state is 1s², 2s²2p⁶, 3s² 3p². It has four valence electrons in their valence orbitals. Silicon undergoes sp³ hybridisation forming four sp³ hybridised orbitals of equal energy.

Each silicon atom forms four covalent bonds with four other neighbouring silicon atoms due to SP³– SP³ Thus tetrahedral Si₄  unit is formed which is extended to form a three-dimensional giant molecule. Due to overlapping of hybrid orbitals, Si-Si covalent bonds are very strong and are directional. In each bond, there are 2 electrons one contributed by each atom.

Thus all valence electrons are used in the formation of covalent bonds. Thus the electrons are localised. Due to which no free electrons are available for conduction. Thus Silicon is an insulator at absolute zero.

Intrinsic Conductivity of Silicon:

Due to heat or at room temperature, some of the covalent bonds in the crystal lattice break up and release free electrons. Thus free electrons are available for conduction of electricity. When a potential difference is applied across the crystal these free electrons move towards the positive terminal of battery leaving behind a hole in the bond.

As the temperature increases the electron-hole pairs formed increase. Electrons move to positive pole and hole moves towards the negative pole of the battery. This conduction of electric current by pure semiconductor at high temperature is called intrinsic conductivity.

On lowering the temperature freed electrons fall back to their position in the hole. Thus due to less availability or unavailability of free electrons, electrical conductivity decreases.

Note:

Due to overlapping of hybrid orbitals in silicon, Si-Si covalent bonds are very strong and are directional. Presence of network of strong covalent bonds accounts for the hardness and its high melting point. Similarly it shows no malleability and no ductility.

Science > Chemistry > Third Row Elements > Metallic Solids of the Third Row

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In this article, we shall study the concept of metallic bond and metallic bonds in sodium, magnesium and aluminium crystals.

Concept of Metallic Bond:

The metallic bond is defined as the force of attraction that binds metal cations to a number of mobile or delocalized electrons within its sphere of influence which holds the metallic cations together in a definite pattern. To explain the nature of metallic bond many theories were proposed.  Free electron theory or electron sea theory is one of the simplest theory proposed by Drude and Lorentz. Some of the important postulates of this theory are as follows:

Free Electron Theory of Metallic Bond:

Metal atoms have less number of valence electrons so they have many vacant valence orbitals. Na, Mg and Al have 3p orbitals vacant.

Ionization potential values of metals are low, hence valence electrons are loosely held and can be easily removed. The closely packed structure of the metallic crystal consists of the atoms of the metal which are identical in all respects.

The unoccupied orbitals of closely packed atoms of the metal overlap with the similar orbital of adjacent atoms through the crystal lattice. Valence electrons removed from their orbitals and can move freely from vacant valency orbits of one atom to other. Since these valence electrons do not belong to any single atom but to crystal as a whole, they are called delocalized or mobile electrons. The metal ions (cations) produced due to delocalization are called kernels. The metal ions have fixed positions in the crystal lattice while delocalized electrons are free to move in the crystal lattice. Thus metal can be considered as an aggregation of metal cations immersed in a sea of mobile electrons.

Since the electrons in the metals are delocalized, and they are assumed to be uniformly distributed throughout the crystal lattice. The forces of attraction between the metal ions and delocalized electrons are uniform in all the direction. Hence the metallic bond is non-directional.

The units occupying lattice points in Sodium, Magnesium, Aluminium are positive ions of them respectively and are surrounded by mobile electrons. Thus Sodium, Magnesium, and Aluminium are metallic solids.

As we move from left to right i.e. from Sodium to Aluminium the no. of valence electrons increases and hence the strength of the bond increases from Sodium to Aluminium.

The properties like electrical and thermal conductivity, metallic lustre, malleability and ductility can be explained on the basis of free electron theory.

Characteristics of Metallic Bonds:

• The metallic bond is defined as the force of attraction that binds metal cations to a number of mobile or delocalized electrons within its sphere of influence which holds the metallic cations together in a definite pattern.
• The metallic bond is non-directional.
• They are weaker than the covalent bond but stronger than van der Waal’s forces.
• The bonds are not rigid.
• The strength of the metallic bond is directly related to the positive charge on the metal ion. So the strength of metallic bond increases as Na < Mg < Al.

Metallic Solids:

Metallic solids are crystalline solids in which the units occupying lattice points are positive ions surrounded by a pool of electrons. (Concept of metallic bond)

Crystal Structures of Metals:

X-ray analysis of different metallic crystals have shown that metals adopt either of the following crystal structures.

• Body centred cubic structure. (BCC)
• Face centred cubic structure. (FCC)
• Hexagonal close-packed structure. (HCP)

Sodium (Na):

Sodium metal has body centred cubic (BCC) open packed crystal structure. Bonding is non-directional metallic bonding.

The arrangement of ions in one plane- Cubic array – open or square packed structure. In this arrangement, each metal ion is touching four adjacent ions in one plane. The sequence of the layers is AB, AB, AB, ……….

Sodium is a metallic solid. In the unit cell of sodium, each sodium ion is surrounded by eight other equidistant sodium ions. Hence co-ordination number is 8. These sodium ions are arranged at the corners of an imaginary cube and at the centre of the cube, one sodium ion is present. There are 2 ions present in one unit cell of sodium.

The sodium ions occupy only about 68% of the available space in a unit cell. So 32% of the unit cell remains empty (void). Since this structure has more empty space metals which adopt this structure are soft. Sodium is thus soft metal because of more empty space (about 32% ) in its crystal structure and rather weak metallic bond due to just one valence electron per Na atom in its crystal.

Magnesium (Mg):

Magnesium has hexagonal close-packed (HCP) crystal structure. Bonding is non-directional metallic bonding.

Magnesium is metallic solid. The units occupying lattice sites are Mg ions and these ions are surrounded by mobile or delocalized electrons.

The arrangement of ions in one plane the arrangement of ions is a hexagonal array or closed packed layer. Thus each metal ion is touching six adjacent ions in one plane. Each magnesium ion touches six magnesium ions in its own layer, three in the layer above and three in the layer below. In hexagonal packed structure, the closed packed layers of ions are stacked in an alternating sequence usually called AB ABA ….. Every third layer of ions is exactly the same as and lie directly above the first layer.

Each Mg atom is surrounded by 12 other equidistant Mg ions. Hence co-ordination number is 12. In the unit cell, about 26% of the available space is empty (void). This structure is more closely packed. Due to less empty space in the crystal structure, more electron cloud due to Mg⁺² and strong metallic bond, magnesium is harder than sodium metal. It is more malleable and ductile than sodium.

Aluminium (Al):

Aluminium has face centred cubic (FCC) crystal structure. Bonding is non-directional metallic bonding.

Aluminium is metallic solid. The units occupying lattice sites are Al ions and these ions are surrounded by mobile or delocalised electrons.

The arrangement of ions in one plane the arrangement of ions is a hexagonal array or closed packed layer. Thus each metal ion is touching six adjacent ions in one plane.

In a cubic close-packed or face centred cubic crystal structure, the sequence of close-packed layers of ion repeats every fourth layer. I.e. every fourth layer of ions is exactly the same as and lies directly above the first layer. It is thus called ABC, ABC, ABC, …..  cubic close packing. Cubic close packing is also called as face centred close-packed structure because if viewed from a particular angle, the ions can be considered as being at the eight corners of the cube and at the centre of each of the six faces of the cube i.e. unit cell.

Each aluminium ion (Al⁺³) is surrounded by 12 other equidistant aluminium atoms. Hence co-ordination number is 12. In the unit cell of aluminium, about 26% of the available space is empty (void). Due to less empty space in the crystal structure, strong metallic bonding, aluminium is harder, more malleable and ductile than sodium and magnesium.

Silicon (Si):

Covalent solids are crystalline solids in which unit lattice points are atoms. The major binding force is covalent bonds between atoms. Silicon is a covalent solid in which lattice points are occupied by the atoms of the element. Silicon is a network solid. There is a network of Si-Si covalent bonds.

The atomic number of Silicon is 14. Electronic configuration of silicon in its ground state is 1s², 2s²2p⁶, 3s² 3p². It has four valence electrons in their valence orbitals. Silicon undergoes sp³ hybridisation forming four sp³ hybridised orbitals of equal energy. Each silicon atom forms four covalent bonds with four other neighbouring silicon atoms due to SP³– SP³ Thus tetrahedral Si₄  unit is formed which is extended to form a three-dimensional giant molecule.

The Si -Si bonds run continuously throughout the crystal. Thus a crystal of silicon is regarded as a giant three-dimensional molecule having a tetrahedral network of silicon atoms bonded together by strong covalent bonds. Solid containing such a structure is called network solid. Si-Si bond angle is 109 ^o 28 ’ while the bond length is 2.35 ^o

Due to overlapping of hybrid orbitals, Si-Si covalent bonds are very strong and are directional. Presence of network of strong covalent bonds accounts for the hardness and its high melting point.

The electrons in the covalent bond are localized hence silicon is not a good conductor of electricity and heat. But it is a semiconductor. Due to the non-availability of free electrons, silicon is an insulator at absolute zero temperature. However, silicon is a semiconductor. If the temperature is increased, covalent bonds break electrons set free which can conduct electricity. Thus conductivity increases with increase in temperature.

Science > Chemistry > Third Row Elements > Concept of Metallic Bond and Metallic Solids

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In this article, we shall discuss the trend in oxidizing and reducing property of third row elements.

Oxidation:

The process in which an atom, a molecule or an ion loses one or more electrons is called oxidation. It is also known as de-electronation.

e.g.            Na      →    Na⁺   +   e^–

In this case, the oxidation of sodium is taking place.

Reduction:

The process in which an atom, a molecule or an ion gains one or more electrons is called reduction. It is also known as electronation.

e.g.             CI      +    e^–    →     Cl^–

In this case, the reduction of chlorine is taking place.

Reducing Agent:

A substance (an atom, a molecule or an ion) which forces another substance to accept electrons and it itself undergoes oxidation by losing electrons is called the reducing agent. Reducing agent is electron donor. e.g.  Na, Al, Mg etc.

Oxidizing Agent:

A substance (an atom, a molecule or an ion) which forces another substance to lose electrons and it itself undergoes reduction by accepting electrons is called oxidizing agent. The oxidizing agent is an electron acceptor e.g. Cl, F, Br, O etc.

Reducing Property:

The tendency of an element to lose electrons is called its reducing property. By virtue of this property, the substance itself undergoes oxidation.

Oxidizing Property:

The tendency of an element to gain electrons is called its oxidizing property. By virtue of this property, the substance itself undergoes reduction.

Factors Affecting the Oxidizing and Reducing Property:

• Ionisation potential
• Electropositivity or electronegativity
• Atomic size
• Metallic and non-metallic character
• Number of valence electrons

Trend in oxidizing and Reducing Property:

• As we move from left to right i.e. from sodium to chlorine along the third row, the oxidizing property goes on increasing while reducing property goes on decreasing.
• Sodium, magnesium, aluminium are good reducing agents. Silicon, phosphorous and sulphur are weak reducing agents.
• Sodium is the strongest reducing agent. Chlorine is the strongest oxidizing agent. Argon is neither oxidizing agent nor reducing agent.

Scientific Reasons:

As we move from left to right i.e. from Sodium to Chlorine along the third row, the oxidizing property goes on increasing while reducing property goes on decreasing.

Oxidizing and reducing strengths of elements depend upon the atomic size, ionization enthalpy, electropositive and electronegative character, and the number of valence electrons.

Those elements who have bigger atomic size, lower ionization enthalpy, and few valence electrons tend to donate electrons and hence are reducing agents. Hence sodium, magnesium, aluminium are reducing agents.

Those elements who have greater ionization enthalpy, smaller atomic size and more valency electrons tend to accept electrons and hence are the oxidizing agent. Hence Chlorine is an oxidizing agent.

It is observed that, as we move from left to right along the third period atomic size gradually decreases, ionization enthalpy increases and the number of valence electrons increases. Hence electron donating tendency of elements goes on decreasing and that of electron gaining tendency of elements goes on increasing. Hence as we move from left to right i.e. from sodium to chlorine along the third row, the oxidizing property goes on increasing while reducing property goes on decreasing.

Sodium, magnesium and aluminium are good reducing agents.

Oxidizing and reducing strength of elements depends upon the atomic size, ionization enthalpy, electropositive and electronegative character and a number of valence electrons.

Atomic numbers of sodium, magnesium and aluminium are 11, 12 and 13 respectively. They have 1, 2, and 3, valence electrons respectively. Thus removal of 1, 2, 3 electrons from sodium, magnesium, aluminium respectively would give the stable inert gas configuration of neon.

Compared to other third row elements these elements have bigger atomic sizes and lower ionization potentials.  Hence they readily lose their valence electrons and are thus strong reducing agents. Reducing strength decrease in the order.  Na > Mg > Al.

Sodium is the strongest reducing agent. 

Oxidizing and reducing strength of elements depends upon the atomic size, ionization enthalpy, electropositive and electronegative character and the number of valence electrons.

The atomic number of sodium is 11. It consists of a single unpaired electron (3s¹). Thus removal of 1 valence electron would give sodium a stable inert gas configuration of neon.

Sodium has the largest atomic size among third low elements and the lowest ionization potential among third-row elements. Hence Sodium readily loses its valence electron and is thus strongest reducing agents.

Na      →    Na⁺   +   e^–

Silicon, phosphorous and sulphur are weak reducing agents.

Oxidizing and reducing strength of elements depends upon the atomic size, ionization enthalpy, electropositive and electronegative character and the number of valence electrons.

Atomic numbers of Silicon, phosphorous and sulphur are 14, 15 and 16 respectively. They have 4, 5, and 6, valence electrons respectively.

They have comparatively smaller atomic size and higher ionization potential. Hence they show less tendency to lose their valence electrons. They are less electropositive. Hence they are weak reducing agents.

 They act as weak reducing agents when treated with strong oxidizing agents like fluorine. They also act as a weak oxidizing agent with strong reducing agents like sodium.

Chlorine is the strongest oxidizing agent in third-row elements.

Oxidizing and reducing strength of elements depends upon the atomic size, ionization enthalpy, electropositive and electronegative character and the number of valence electrons.

Chlorine has high electronegativity. The atomic number of chlorine is 17. It consists of seven electrons in the valence shell. It requires only one electron to complete its octet and attain a stable electronic configuration of argon.

Chlorine has the smallest atomic size among third low elements.

Chlorine has very high ionization potential and very high electron affinity.

Hence it has a strong tendency to gain an electron and it is highly electronegative. Hence Chlorine is the strong oxidizing agent.

CI      +    e^–    →     Cl^–

Argon is neither oxidizing agent nor reducing agent.

Oxidizing and reducing strength of elements depends upon the atomic size, ionization enthalpy, electropositive and electronegative character and the number of valence electrons.

Argon is neither electropositive nor electronegative element Atomic number of Argon is 18. It consists of eight electrons in the valence shell. Thus it has completed octet. s orbital and p orbitals are completely filled. Hence it has a stable electronic configuration.

Argon has very high ionization enthalpy. Hence it has a strong tendency not to gain or lose electrons. Hence Argon is neither oxidizing agent nor reducing agent.

Science > Chemistry > Third Row Elements > Oxidizing and Reducing Property of the Third Row Elements

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In this article, we shall study the trend im metallic character of third row elements.

Metallic Character:

The tendency of an atom to lose electrons to form positively charged ion is called its metallic character or electropositive character. Typically metals show luster and are good conductors of heat and electricity.

Non-Metallic Character:

The tendency of an atom to gain electrons to form negatively charged ion is called its non-metallic character or electronegative character.

Factors affecting metallic and non-metallic characters:

• Size of an atom
• Ionization enthalpy
• Nuclear charge

Trends in metallic and non-metallic characters:

As we move across the periodic table from left to right i.e. from sodium to argon in the third period metallic character decreases and non-metallic character increases. Sodium, magnesium, aluminium are typical metals. Silicon is metalloid. (weakly non- metallic). Phosphorous, Sulphur, Chlorine are non-metals. Argon is an inert gas.

Sodium is the most metallic element while chlorine is the most non-metallic element. Both the extreme elements sodium and chlorine are extremely reactive. Argon is neither electropositive nor electronegative.

Scientific Reasons:

As we move across the periodic table from left to the right i.e. from sodium to Argon in the third period metallic character decreases and non-metallic character increases.

As we move from left to right in the third period the atomic number goes on increasing. Thus as we move from sodium to argon the nuclear charge increases and from left to right additional electron is added to the same i.e. third orbit. Due to increase in the nuclear charge the attractive force on electrons in outermost orbit increases and thus the size of atom decreases.

As we move from left to right in the third period the ionization enthalpy increases. From left to right in the third-period electronegativity increases i.e. the tendency of losing electrons decreases and tendency of gaining electrons increases. As we move from left to right in the third-period number of valence increases. Hence as we move across the periodic table from left to the right i.e. from sodium to Argon in the third period metallic character decreases and non-metallic character increases.

Sodium, Magnesium, Aluminium are typical metals.

Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionization potential. Sodium, magnesium, aluminium have atomic numbers 11, 12, 13 respectively. They contain 1, 2, 3 valence electrons in their outermost shell.

Compared to other elements in the third row their atomic sizes are larger. Similarly, the nuclear charges are also less. Due to this attractive force on valence electrons is less. Compared to other elements in the third row their ionization enthalpies are lower. Thus the tendency of these elements is to lose their valence electrons. Hence sodium, magnesium, aluminium are typical metals.

Phosphorous, Sulphur, Chlorine are non-metals.

Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionization potential. Phosphorous, sulphur, chlorine have atomic numbers 15, 16, 17 respectively. They contain 5, 6, 7 valence electrons in their outermost shell.

Compared to other elements in the third row their atomic sizes are smaller. Similarly, the nuclear charges are also more. Due to this attractive force on valence electrons is more. Compared to other elements in the third row their ionization enthalpies are higher. Thus the tendency of these elements is not to lose their valence electrons but to gain electrons. Hence phosphorous, sulphur, chlorine are non-metals.

Argon is neither electropositive nor electronegative element.

Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionization potential. Argon has atomic number 18. It contains 8 valence electrons in their outermost shell. Argon has completed octet and completely filled s and p orbitals thus it has the most stable electronic configuration in third-row elements which it will not disturb by accepting or losing the electron.

Compared to other elements in the third row its atomic sizes are larger but the effect of filled s and p orbitals is dominating. Compared to other elements in the third row its ionization enthalpy is the highest. Thus the tendency of Argon to lose or gain valence electrons is almost absent. Hence Argon is neither electropositive nor electronegative element.

Sodium is the most metallic element.

Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionization potential. Sodium has atomic numbers 11. It contains 1 valence electron in its outermost shell.

Compared to other elements in the third row its atomic size is larger. Similarly, the nuclear charge is also less. Due to this attractive force on valence electron is less. Compared to other elements in the third row its ionization enthalpy is the lowest. Thus the tendency of Sodium is to lose its valence electron. Hence Sodium is the most metallic element.

Chlorine is the most non-metallic element. 

Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionization potential. Chlorine has atomic number 17. It contains 7 valence electrons in its outermost shell.

Compared to other elements in the third row its atomic size is the smallest. Similarly, the nuclear charge is also more. Due to this attractive force on valence electrons is more. Compared to other elements in the third row its ionization enthalpy is the highest. Thus the tendency of Chlorine is not to lose its valence electrons but to gain electrons. Hence Chlorine is the most non-metallic element.

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In this article, we shall study trend in the ionization enthalpy of third-row elements.

Ionization Enthalpy:

The minimum energy required to remove the most loosely attached electron from the outermost shell of a neutral gaseous, isolated atom of an element in its ground state to produce gaseous cation is known as ionization enthalpy or ionization potential of that element. Ionization enthalpy is expressed in terms of kJ per mole

M_(g)       +    I. P.       →     M_(g)⁺     +       e^–

Where M is the third-row element

Factors Affecting the Ionization Enthalpy:

• The size (atomic radius) of an atom i.e. the distance of the outermost electron from the nucleus.
• The charge on the nucleus or the nuclear charge i.e. protons present in the nucleus.
• The screening effect.
• The type or the geometry of the subshell in which the electron is present.
• For the third row elements, the screening effect for all the elements is almost the same.

Trends in the Ionization Enthalpy Across the Period:

With some minor exceptions, the trend in ionization enthalpy of third-row elements is that as we move from left to the right i.e. from Sodium to Argon along period, ionization enthalpy of these elements goes on increasing steadily.

Its values for the third-row elements are given below.

The irregularity in the trend can be observed from the following graph.

The ionization enthalpy of magnesium is more than aluminium. The ionization enthalpy of phosphorous is more than sulphur. The last element argon has the highest ionization enthalpy. Magnesium, phosphorous and argon have more ionization enthalpy than expected due to the extra stability of their exactly half filled and completely filled orbitals.

Scientific Reasons:

Ionization enthalpy of third-row elements goes on increasing steadily along the period. 

Ionization enthalpy of an element depends upon i) atomic size ii) nuclear charge iii) screening effect.

If the element has smaller atomic size, greater nuclear charge and less dense inter electronic cloud, then it has greater ionization enthalpy. If the element has bigger atomic size, low nuclear charge and denser electronic cloud then it has low ionization potential.

Along the third period i.e. from Na to Ar, as atomic number increases, the nuclear charge goes on increasing, atomic size goes on decreasing, and the number of valence electron goes on increasing. The attractive force on the valence electron of an atom increases. Hence Ionization enthalpy of the third-row elements goes on increasing steadily along the period.

The ionisation enthalpy of magnesium is more than aluminium.

The atomic number of magnesium is 12. Its electronic configuration is 2, 8, 2. Its detailed configuration is 1s²2s² 2p⁶3s². The box diagram for final orbit configuration for magnesium is as below. We can see that Mg has completely filled ‘3s’ orbitals.

The atomic number of aluminium is 13. Its electronic configuration is 2, 8, 3. Its detailed configuration is 1s²2s² p⁶3s²3p¹. The box diagram for final orbit configuration for aluminium is as below. We can see that aluminium has partially filled ‘3p’ orbitals.

It has been observed that the extra stability is associated with vacant, half filled and completely filled orbitals. Magnesium has thus extra stable electronic state due to completely filled ‘3s’ orbital. It is difficult to remove an electron from pair due to extra stability. While aluminium has no such extra stable state.  There is one un­paired electron in the 3p subshell of aluminium.

s orbitals are closer to the nucleus than p orbitals. So energy required to remove the electron from s orbital is always more than to remove it from p orbital. In the case of magnesium, the electron is to be removed from s orbital while in case of aluminium it is to be removed from p orbitals. Thus more energy will be required to remove valence electron from magnesium atom than that of aluminium. So ionization enthalpy of magnesium is greater than that of the aluminium.

The ionisation enthalpy of phosphorous is more than sulphur.

The atomic number of Phosphorous is 15. Its electronic configuration is 2, 8, 5. Its detail configuration is 1s², 2s² 2p⁶, 3s² 3p³. The box diagram for final orbit configuration for phosphorous is as below. We can see that the p orbitals are exactly half filled.

The atomic number of Sulphur is 16. Its electronic configuration is 2, 8, 6.Its detail configuration is 1s², 2s² 2p⁶, 3s² 3p⁴. The box diagram for final orbit configuration for sulphur is as below. We can see that sulphur has partially filled ‘3p’ orbitals.

It has been observed that the extra stability is associated with vacant, half-filled and completely filled orbitals. Phosphorous has thus extra stable electronic state due to exactly half filled ‘3p’ orbital. Sulphur has no such extra stable state as there are two un­paired electrons in 3p subshell. In sulphur, removing one electron from 3p orbital, it becomes half-filled and attains the stable state. Thus sulphur tries to lose electron fast. Thus more energy will be required to remove valence electron of phosphorous than that of sulphur.  So ionization potential of phosphorous is greater than that of the sulphur.

Argon has the highest ionization potential.

The atomic number of argon is 18. Its electronic configuration is 2, 8, 8. Its detail configuration is 1s², 2s² 2p⁶, 3s² 3p⁶. The box diagram for final orbit configuration for argon is as below. We can see that argon has completely filled ‘3s’ and ‘3p’ orbitals.

It has been observed that the extra stability is associated with empty, half-filled and completely filled orbitals. Argon the as most stable electronic configuration. It has a complete octet of electrons in the outer most shell. All electrons are paired. It is very difficult to remove valence electrons. Hence much more energy is required to remove an electron from such a stable configuration. So Argon has the highest ionization potential.

Concept of Higher Ionization Potentials:

Aluminium rarely forms the tri-positive cation. 

Aluminium has atomic number 13. Its electronic configuration is 1s²2s² p⁶3s²3p¹. To form a tri-positive ion of aluminium all the three electrons of the outermost shell are to be removed.

Formation of mono-positive ion i.e. cation involves the removal of an electron from a neutral atom. Hence the first ionization potential is always low. Formation of di-positive ions involves removal of an electron from mono-positive cation while the formation of tri-positive ion involves removal of an electron from di-positive cation. Due to the positive charge on the mono-positive and dipositive cations, the outgoing electron has to overcome larger attractive force.

Hence more energy is required to remove the second electron and still more energy is required to remove the third electron. Hence Aluminium rarely forms tri-positive ions.

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The elements are arranged in the order of increasing atomic numbers, in such a way that elements with similar properties fall in the same vertical column of the periodic table. There are eighteen vertical columns known as groups and seven horizontal rows are known as periods. A group of eight elements namely, sodium (Na), magnesium (Mg), aluminium (Al), silicon (Si), phosphorous (P), sulphur (S), chlorine (Cl) and argon (Ar) belongs to the third period of the periodic table are called as third-row elements.

The properties of each element are character­istics of the group to which they belong.  Each one of these elements can, therefore, be considered as the representative of the whole group to which it belongs.  Hence these elements are known as typical or representative elements.

The properties of these elements gradually change from metallic or basic to non-metallic or acidic character across the third period. Sodium, magnesium and aluminium are metals.  They have a metallic lustre, can conduct heat and electricity.  They are malleable and ductile.  Sodium is a very soft metal. Silicon is a hard solid and is metalloid. Phosphorous is a yellow waxy solid.  It is non-metal.  Sulphur is a yellow coloured solid.  It is a non-metal.  Chlorine and argon are gases at room temperature.  They are non-metals.

Position of the Third Row Elements in the Periodic table:

In the modern periodic table, third row elements have been placed in 1, 2, 13, 14, 15, 16, 17 and 18 groups. Except for Sodium and Argon, the third-row elements are the second member of their group. Sodium and Argon are the third members of their groups 1 and 18 respectively

For sodium and magnesium, the last electron to be configured enter into ‘s’ orbitals. Hence sodium and magnesium are s block elements. For all the other third row elements the last electron to be configured enter into ‘ ‘p’ orbitals hence they are p block elements.

There are three orbits and valence electrons are present in the third main energy level. All these elements have vacant 3d orbitals. These elements are capable of transferring their valence electrons to these vacant d orbitals. Thus they are capable of expanding their octet which is not possible in the second-row elements

Electronic Configuration of Third Row Elements:

Electronic configuration of third row elements of the third row is as fol­lows

It can be seen that Sodium (Z = 11) has a single electron in its 3 s valence orbital. With the increase in atomic number, in magnesium, aluminium, silicon, phosphorous, and chlorine, the electrons successively occupy 3s and 3p valence orbitals until another closed-shell configu­ration 1s²2s² 2p⁶3s²3p⁶ is reached at argon (Z = 18).

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