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Important Compound of Transition Elements

Oxides and Oxoanions of Metals

Generally, transition metal oxides are formed by the reaction of transition metals with molecular oxygen at high temperatures. Except the first member of 3d series, Scandium, all other transition elements form ionic metal oxides. The oxidation number of metal in metal oxides ranges from +2 to +7. As the oxidation number of a metal increases, ionic character decreases, for example, Mn₂O₇ is covalent.

Mostly higher oxides are acidic in nature, Mn₂O₇ dissolves in water to give permanganic acid (HMnO₄), similarly CrO₃ gives chromic acid (H₂CrO₄) and dichromic acid (H₂Cr₂O₇). Generally lower oxides may be amphoteric or basic, for example, Chromium (III) oxide – Cr₂O₃, is amphoteric and Chromium(II) oxide, CrO, is basic in nature.

Potassium Dichromate K₂Cr₂O₇

Preparation:

Potassium dichromate is prepared from chromate ore. The ore is concentrated by gravity separation. It is then mixed with excess sodium carbonate and lime and roasted in a reverbratory furnace.

The roasted mass is treated with water to separate soluble sodium chromate from insoluble iron oxide. The yellow solution of sodium chromate is treated with concentrated sulphuric acid which converts sodium chromate into sodium dichromate.

The above solution is concentrated to remove less soluble sodium sulphate. The resulting solution is filtered and further concentrated. It is cooled to get the crystals of Na₂SO₄.2H₂O.

The saturated solution of sodium dichromate in water is mixed with KCl and then concentrated to get crystals of NaCl. It is filtered while hot and the filtrate is cooled to obtain K₂Cr₂O₇ crystals.

Physical Properties:

Potassium dichromate is an orange red crystalline solid which melts at 671K and it is moderately soluble in cold water, but very much soluble in hot water. On heating it decomposes and forms Cr₂O₃ and molecular oxygen. As it emits toxic chromium fumes upon heating, it is mainly replaced by sodium dichromate.

Structure of Dichromate Ion:

Both chromate and dichromate ion are oxo anions of chromium and they are moderately strong oxidizing agents. In these ions chromium is in +6 oxidation state. In an aqueous solution, chromate and dichromate ions can be interconvertible, and in an alkaline solution chromate ion is predominant, whereas dichromate ion becomes predominant in acidic solutions. Structures of these ions are shown in the figure.

Chemical Properties:

1. Oxidation

Potassium dichromate is a powerful oxidising agent in acidic medium. Its oxidising action in the presence of H⁺ ions is shown below. You can note that the change in the oxidation state of chromium from Cr⁶⁺ to Cr³⁺. Its oxidising action is shown below.

Cr₂O₇^2- + 14H⁺ + 6e^– → 2Cr³⁺ + 7H₂O

The oxidising nature of potassium dichromate (dichromate ion) is illustrated in the following examples.

(i) It oxidises ferrous salts to ferric salts.

Cr₂O₇^2- + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

(ii) It oxidises iodide ions to iodine

Cr₂O₇^2- + 6I^– + 14H⁺ → 2Cr³⁺ + 3I₂ + 7H₂O

(iii) It oxidises sulphide ion to sulphur

Cr₂O₇^2- + 3S^2- + 14H⁺ → 2Cr³⁺ + 3S + 7H₂O

(iv) It oxidises sulphur dioxide to sulphate ion

Cr₂O₇^2- + 3SO₂ + 2H⁺ → 2Cr³⁺ + 3SO^2-₄ + H₂O

(v) It oxidises stannous salts to stannic salt

Cr₂O₇^2- + 3Sn²⁺ + 14H⁺ → 2Cr³⁺ + 3Sn⁴⁺ + 7H₂O

(vi) It oxidises alcohols to acids.

2K₂Cr₂O₇ + 8H₂SO₄ + 3CH₃CH₂OH →
2K₂SO₄ + 2Cr₂(SO₄)₃ + 3CH₃COOH + 11H₂O

Chromyl Chloride Test:

When potassium dichromate is heated with any chloride salt in the presence of Conc H₂SO₄, orange red vapours of chromyl chloride (CrO₂Cl₂) is evolved. This reaction is used to confirm the presence of chloride ion in inorganic qualitative analysis.

The chromyl chloride vapours are dissolved in sodium hydroxide solution and then acidified with acetic acid and treated with lead acetate. A yellow precipitate of lead chromate is obtained.

CrO₂Cl₂ + 4NaOH → Na₂CrO₄ + 2NaCl + 2H₂O

Uses of Potassium Dichromate:

Some important uses of potassium dichromate are listed below.

1. It is used as a strong oxidizing agent.
2. It is used in dyeing and printing.
3. It used in leather tanneries for chrome tanning.
4. It is used in quantitative analysis for the estimation of iron compounds and iodides.

Potassium Permanganate – KMnO₄

Preparation:
Potassium permanganate is prepared from pyrolusite (MnO₂) ore. The preparation involves the following steps.

(i) Conversion of MnO₂ to potassium manganate:

Powdered ore is fused with KOH in the presence of air or oxidising agents like KNO₃ or KClO₃. A green coloured potassium manganate is formed.

(ii) Oxidation of potassium manganate to potassium permanganate:

Potassium manganate thus obtained can be oxidised in two ways, either by chemical oxidation or electrolytic oxidation.

Chemical Oxidation:

In this method potassium manganate is treated with ozone (O₃) or chlorine to get potassium permanganate.

2MnO₄^2- + O₃ + H₂O → 2MnO₄^– + 2OH^– + O₂
2MnO₄^2- + Cl₂ → 2MnO₄^– + 2Cl^–

Electrolytic Oxidation

In this method aqueous solution of potassium manganate is electrolyzed in the presence of little alkali.

K₂MnO₄ ⇄ 2K⁺ + MnO₄^2-
H₂O ⇄ H⁺ + OH^–

Manganate ions are converted into permanganate ions at anode.

H₂ is liberated at the cathode.

2H⁺ + 2e^– → H₂↑

The purple coloured solution is concentrated by evaporation and forms crystals of potassium permanganate on cooling.

Physical Properties:

Potassium permanganate exists in the form of dark purple crystals which melts at 513 K. It is sparingly soluble in cold water but, fairly soluble in hot water.

Structure of Permanganate ion

Permanganate ion has tetrahedral geometry in which the central Mn⁷⁺ is sp³ hybridised.

Chemical Properties:

1. Action of Heat:

When heated, potassium permanganate decomposes to form potassium manganate and manganese dioxide.

2KMnO₄ → 2K₂MnO₄ + MnO₂ + O₂

2. Action of conc H₂SO₄

On treating with cold conc H₂SO₄, it decomposes to form manganese heptoxide, which subsequently decomposes explosively.

But with hot conc H₂SO₄, Potassium permanganate give MnSO_4. 

3. Oxidising Property:

Potassium permanganate is a strong oxidising agent, its oxidising action differs in different reaction medium.

(a) In neutral medium:

In neutral medium, it is reduced to MnO₂

MnO₄^– + 2H₂O + 3e^– → MnO₂ + 4OH^–

(i) It oxidises H₂S to sulphur

2MnO₄^– + 2H₂O + 3e^– → MnO₂ + 4OH^–

(ii) It oxidises thiosulphate into sulphate

8MnO₄^– + 3S₂O₃^2- + H₂O → 6SO₄^2- + 8MnO₂ + 2OH^–

(b) In alkaline medium:

In the presence of alkali metal hydroxides, the permanganate ion is converted into manganate.

MnO₄^– + e^– → MnO₄^2-

This manganate is further reduced to MnO2 by some reducing agents.

MnO₄^2- + H₂O → MnO₂ + 2OH^– + [O]

So the overall reaction can be written as follows.

MnO₄^– + 2H₂O + 3e^– → MnO₂ + 4OH^–

This reaction is similar as that for neutral medium.

Bayer’s Reagent:

Cold dilute alkaline KMnO₄ is known as Bayer’s reagent. It is used to oxidise alkenes into diols. For example, ethylene can be converted into ethylene glycol and this reaction is used as a test for unsaturation.

(c) In acid medium:

In the presence of dilute sulphuric acid, potassium permanganate acts as a very strong oxidising agent. Permanganate ion is converted into Mn²⁺ ion.

MnO₄^– + 8H⁺ + 5e^– → Mn²⁺ + 4H₂O

The oxidising nature of potassium permanganate (permanganate ion) in acid medium is illustrated in the following examples.

(i) It oxidises ferrous salts to ferric salts.

2MnO₄^– + 10Fe₂₊ + 16H⁺ → 2Mn²⁺ + 10Fe³⁺ + 8H₂O

(ii) It oxidises iodide ions to iodine

2MnO₄^– + 10 I^– + 16H⁺ → 2Mn²⁺ + 5I₂ + 8H₂O

(iii) It oxidises oxalic acid to CO₂

2MnO₄^– + 5(COO)^2-₂ + 16H⁺ → 2Mn²⁺ + 10CO₂ + 8H₂O

(iv) It oxidises sulphide ion to sulphur

2MnO^–₄ + 5 S^2- + 16H⁺ → 2Mn²⁺ + 5 S + 8H₂O

(v) It oxidises nitrites to nitrates

2MnO₄^– + 5 NO₂^– + 6H⁺ → 2Mn²⁺ + 5NO^–₃ + 3H₂O

(vi) It oxidises alcohols to aldehydes.

2KMnO₄ + 3H₂SO₄ + 5CH₃CH₂OH → K₂SO₄ + 2MnSO₄ + 5CH₃CHO + 8H₂O

(vii) It oxidises sulphite to sulphate

2MnO₄^– + 5SO₃^2- + 6H⁺ → 2Mn²⁺ + 5SO₄^2- + 3H₂O

Uses of Potassium Permanganate:

Some important uses of potassium permanganate are listed below.

1. It is used as a strong oxidizing agent.
2. It is used for the treatment of various skin infections and fungal infections of the foot.
3. It used in water treatment industries to remove iron and hydrogen sulphide from well water.
4. It is used as Bayer’s reagent for detecting unsaturation in an organic compound.
5. It is used in quantitative analysis for the estimation of ferrous salts, oxalates, hydrogen peroxide and iodides.

F-Block Elements – Inner Transition Elements

In the inner transition elements there are two series of elements.

1. Lanthanoids (previously called lanthanides)
2. Actinoids (previously called actinides)

Lanthanoid series consists of fourteen elements from Cerium (₅₈Ce) to Lutetium (₇₁Lu) following Lanthanum (₅₇La). These elements are characterised by the preferential filling of 4f orbitals, Similarly actinoids consists of 14 elements from Thrium (₉₀Th) to Lawrencium (₁₀₃Lr) following Actinium (₈₉Ac). These elements are characterised by the preferential filling of 5f orbital.

The position of Lanthanoids in the periodic table

The actual position of Lanthanoids in the periodic table is at group number 3 and period number 6. However, in the sixth period after lanthanum, the electrons are preferentially filled in inner 4f sub shell and these fourteen elements following lanthanum show similar chemical properties. Therefore these elements are grouped together and placed at the bottom of the periodic table. This position can be justified as follows.

1. Lanthanoids have general electronic configuration [Xe] 4f^1-14 5d^10-1 6s²
2. The common oxidation state of lanthanoides is +3
3. All these elements have similar physical and chemical properties.

Similarly the fourteen elements following actinium resemble in their physical and chemical properties. If we place these elements after Lanthanum in the periodic table below 4d series, the properties of the elements belongs to a group would be different and it would affect the proper structure of the periodic table. Hence a separate position is provided to the inner transition elements as shown in the figure.

Electronic Configuration of Lanthanoids:

We know that the electrons are filled in different orbitals in the order of their increasing energy in accordance with Aufbau principle. As per this rule after filling 5s, 5p and 6s and 4f level begin to fill from lanthanum, and hence the expected electronic configuration of Lanthanum(La) is [Xe] 4f¹ 5d° 6s² but the actual electronic configuration of Lanthanum is [Xe] 4f° 5d¹ 6s² and it belongs to d block.

Filling of 4f orbital starts from Cerium (Ce) and its electronic configuration is [Xe] 4f¹ 5d¹ 6s². As we move from Cerium to other elements the additional electrons are progressively filled in 4f orbitals as shown in the table.

Table: Electronic Configuration of Lanthanum and Lanthanoids

In Gadolinium (Gd) and Lutetium (Lu) the 4f orbitals, are half-filled and completely filled, and one electron enters 5d orbitals. Hence the general electronic configuration of 4f series of elements can be written as [Xe] 4f^1-14 5d^0-1 6s²

Oxidation State of Lanthanoids:

The common oxidation state of lanthanoids is +3. In addition to that some of the lanthanoids also show either +2 or +4 oxidation states. Gd³⁺ and Lu³⁺ ions have extra stability, it is due to the fact that they have exactly half filled and completely filled f-orbitals respectively their electronic configurations are

Gd³⁺: [Xe]4f⁷
Lu³⁺: [Xe]4f¹⁴

Similarly Cerium and terbium attain 4f° and 4f⁷ configurations respectively in the +4 oxidation states. Eu²⁺
and Yb²⁺ ions have exactly half filled and completely filled f orbitals respectively.

The stability of different oxidation states has an impact on the properties of these elements the following table shows the different oxidation states of lanthanoids.

Atomic and Ionic Radii:

As we move across 4f series, the atomic and ionic radii of lanthanoids show gradual decrease with increse in atomic number. This decrease in ionic size is called lanthanoid contraction.

Cause of Lanthanoid Contraction:

As we move from one element to another in 4f series (Ce to Lu) the nuclear charge increases by one unit and an additional electron is added into the same inner 4f sub shell. We know that 4f sub shell have a diffused shapes and therefore the shielding effect of 4f elelctrons relatively poor hence, with increase of nuclear charge, the valence shell is pulled slightly towards nucleus. As a result, the effective nuclear charge experienced by the 4f elelctorns increases and the size of Ln³⁺ ions decreases. Lanthanoid contraction of various lanthanoids is shown in the graph.

Consequences of Lanthanoid Contraction:

1. Basicity Differences

As we from Ce³⁺ to Lu³⁺, the basic character of Ln³⁺ ions decrease. Due to the decrease in the size of Ln³⁺ ions, the ionic character of Ln – OH bond decreases (covalent character increases) which results in the decrease in the basicity.

2. Similarities Among Lanthanoids:

In the complete f – series only 10 pm decrease in atomic radii and 20 pm decrease in ionic radii is observed because of this very small change in radii of lanthanoids, their chemical properties are quite similar.

The elements of the second and third transition series resemble each other more closely than the elements of the first and second transition series. For example

Series	Element	Atomic Radius
3d Series	Ti	132 pm
4d Series	Zr	145 pm
5d Series	Hf	144 pm

Actinoids:

The fourteen elements following actinium, i.e., from thorium (Th to lawrentium (Lr) are called actinoids. Unlike the lanthanoids, all the actinoids are radioactive and most of them have short half lives. Only thorium and uranium (U) occur in significant amount in nature and a trace amounts of Plutonium (Pu) is also found in Uranium ores. Neptunium (Np) and successive heavier elements are produced synthetically by the artificial transformation of naturally occuring elements by nuclear reactions. Similar to lanthanoids, they are placed at the bottom of the periodic table.

Electronic Configuration:

The electronic configuration of actinoids is not definite. The general valence shell electronic configuration of 5f elements is represented as [Rn]5f^1-146d^0-27s². The following table show the electronic configuration of actinoids.

Table: Electronic configuration of actinoids

Oxidation State of Actinoids:

Like lanthanoids, the most common state of actinoids is +3. In addition to that actinoids show variable oxidation states such as +2 , +3 , +4 ,+5,+6 and +7. The elements Americium(Am) and Thrium (Th show +2 oxidation state in some compounds, for example thorium iodide (ThI₂). The elements T, Pa, U, Np, Pu and Am show +5 oxidation states. Np and Pu exhibit +7 oxidation state.

Differences Between Lanthanoids and Actinoids:

Lanthanoids	Actinoids
1. Differentiating electron enters in 4f orbital	1. Differentiating electron enters in 5f orbital
2. Binding energy of 4f orbitals are higher	2. Binding energy of 5f orbitals are lower
3. They show less tendency to form complexes	3. They show greater tendency to form complexes
4. Most of the lanthanoids are colourless	4. Most of the actinoids are coloured. For example. U3+(red), U4+(green), UO22+(yellow)
5. They do not form oxo cations	5. They do form oxo cations such as UO22+, NPO22+ etc
6. Besides +3 oxidation states lanthanoids show +2 and +4 oxidation states in few cases	6. Besides +3 oxidation states actinoids show higher oxidation states such as +4, +5, +6 and +7

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General Trend in Properties

Metallic Behaviour:

All the transition elements are metals. Similar to all metals the transition metals are good conductors of heat and electricity. Unlike the metals of Group-1 and group-2, all the transition metals except group 11 elements are hard. Of all the known elements, silver has the highest electrical conductivity at room temperature. Most of the transition elements are hexagonal close packed, cubic close packed or body centrered cubic which are the characteristics of true metals.

Figure 4.2 Lattice Structures of 3d, 4d, and 5d transistion metals

As we move from left to right along the transition metal series, melting point first increases as the number of unpaired d electrons available for metallic bonding increases, reach a maximum value and then decreases, as the d electrons pair up and become less available for bonding.

For example, in the first series the melting point increases from Scandium (m.pt 1814K) to a maximum of 2183 K for vanadium, which is close to 2180K for chromium. However, manganese in 3d series and Tc in 4d series have low melting point. The maximum melting point at about the middle of transition metal series indicates that d⁵ configuration is favourable for strong interatomic attraction. The following figure shows the trends in melting points of transition elements.

Variation of Atomic and Ionic Size:

It is generally expected a steady decrease in atomic radius along a period as the nuclear charge increases and the extra electrons are added to the same sub shell. But for the 3d transition elements, the expected decrease in atomic radius is observed from Sc to V, thereafter up to Cu the atomic radius nearly remains the same.

As we move from Sc to Zn in 3d series the extra electrons are added to the 3d orbitals, the added 3d electrons only partially shield the increased nuclear charge and hence the effective nuclear charge increases slightly. However, the extra electrons added to the 3d sub shell strongly repel the 4s electrons and these two forces are operated in opposite direction and as they tend to balance each other, it leads to constancy in atomic radii.

At the end of the series, d – orbitals of Zinc contain 10 electrons in which the repulsive interaction between the electrons is more than the effective nuclear charge and hence, the orbitals slightly expand and atomic radius slightly increases. Generally as we move down a group atomic radius increases, the same trend is expected in d block elements also. As the electrons are added to the 4d sub shell, the atomic radii of the 4d elements are higher than the corresponding elements of the 3d series.

However there is an unexpected observation in the atomic radius of 5d elements which have nearly same atomic radius as that of corresponding 4d elements. T is is due to lanthanoide contraction which is to be discussed later in this unit under inner transition elements.

Ionization Enthalpy:

Ionization energy of transition element is intermediate between those of s and p block elements. As we move from left to right in a transition metal series, the ionization enthalpy increases as expected. This is due to increase in nuclear charge corresponding to the filling of d electrons. The following figure show the trends in ionisation enthalpy of transition elements.

The increase in first ionisation enthalpy with increase in atomic number along a particular series is not regular. The added electron enters (n-1)d orbital and the inner electrons act as a shield and decrease the effect of nuclear charge on valence ns electrons. Therefore, it leads to variation in the ionization energy values.

The ionisation enthalpy values can be used to predict the thermodynamic stability of their compounds. Let us compare the ionisation energy required to form Ni²⁺ and Pt²⁺ ions.

For Nickel, IE₁ + IE₂
= (735 + 1753)
= 2490 kJmol^-1

For Platinum, IE₁ + IE₂
= (864 + 1791)
= 2655 kJmol^-1

Since, the energy required to form Ni²⁺ is less than that of Pt²⁺, Ni(II) compounds are thermodynamically more stable than Pt(II) compounds.

Oxidation State:

The first transition metal Scandium exhibits only +3 oxidation state, but all other transition elements exhibit variable oxidation states by loosing electrons from (n-1)d orbital and ns orbital as the energy difference between them is very small. Let us consider the 3d series; the following table summarizes the oxidation states of the 3d series elements.

At the beginning of the series, +3 oxidation state is stable but towards the end +2 oxidation state becomes stable. The number of oxidation states increases with the number of electrons available, and it decreases as the number of paired electrons increases.

Hence, the first and last elements show less number of oxidation states and the middle elements with more number of oxidation states. For example, the first element Sc has only one oxidation state +3; the middle element Mn has six different oxidation states from +2 to +7. The last element Cu shows +1 and +2 oxidation states only.

The relative stability of different oxidation states of 3d metals is correlated with the extra stability of half filled and fully filled electronic confiurations. Example: Mn²⁺(3d⁵) is more stable than Mn⁴⁺(3d³)

The oxidation states of 4d and 5d metals vary from +3 for Y and La to +8 for Ru and Os. The highest oxidation state of 4d and 5d elements are found in their compounds with the higher electronegative elements like O, F and Cl. for example: RuO₄, OsO₄ and WCl₆. Generally in going down a group, a stability of the higher oxidation state increases while that of lower oxidation state decreases.

It is evident from the Frost diagram (ΔG° vs oxidation number) as shown below. For titanium, vanadium and chromium, the most thermodynamically stable oxidation state is +3. For iron, the stabilities of +3 and +2 oxidation states are similar. Copper is unique in 3d series having a stable +1 oxidation state. It is prone to disproportionate to the +2 and 0 oxidation states.

Standard Electrode Potentials of Transition Metals

Redox reactions involve transfer of electrons from one reactant to another. Such reactions are always coupled, which means that when one substance is oxidised, another must be reduced. The substance which is oxidised is a reducing agent and the one which is reduced is an oxidizing agent. The oxidizing and reducing power of an element is measured in terms of the standard electrode potentials.

Standard electrode potential is the value of the standard emf of a cell in which molecular hydrogen under standard pressure (1 atm) and temperature (273 K) is oxidised to solvated protons at the electrode. If the standard electrode potential (E°), of a metal is large and negative, the metal is a powerful reducing agent, because it loses electrons easily. Standard electrode potentials (reduction potential) of few first transition metals are given in the following table.

In 3d series as we move from Ti to Zn, the standard reduction potential (\(E^{0}{ }_{M^{2+}} /_{M}\)) value is approaching towards less negative value and copper has a positive reduction potential. i.e., elemental copper is more stable than Cu²⁺.

There are two deviations., In the general trend, Fig shows that (\(E^{0}{ }_{M^{2+}} /_{M}\)) value for manganese and zinc are more negative than the regular trend. It is due to extra stability which arises due to the half filled d⁵ configuration in Mn²⁺ and completely filled d¹⁰ configuration in Zn²⁺.

Transition metals in their high oxidation states tend to be oxidizing. For example, Fe³⁺ is moderately a strong oxidant, and it oxidises copper to Cu²⁺ ions. The feasibility of the reaction is predicted from the following standard electrode potential values.

Fe³⁺(aq) + e^– ⇄ Fe²⁺ E° = 0.77V
Cu²⁺(aq) + 2e^– ⇄ Cu(s) E° = +0.34V

The standard electrode potential for the M³⁺/M²⁺ half-cell gives the relative stability between M³⁺ and M²⁺. The reduction potential values are tabulated as below.

The negative values for titanium, vanadium and chromium indicate that the higher oxidation state is preferred. If we want to reduce such a stable Cr³⁺ ion, strong reducing agent which has high negative value for reduction potential like metallic zinc (E° = – 0.76 V) is required.

The high reduction potential of Mn³⁺/Mn²⁺ indicates Mn²⁺ is more stable than Mn³⁺. For Fe³⁺/Fe²⁺ the reduction potential is 0.77V, and this low value indicates that both Fe³⁺ and Fe²⁺ can exist under normal conditions. The drop from Mn to Fe is due to the electronic structure of the ions concerned.

Mn³⁺ has a 3d⁴ configuration while that of Mn²⁺ is 3d⁵. The extra stability associated with a half filled d sub shell makes the reduction of Mn³⁺ very feasible (E° = +1.51V).

Magnetic Properties

Most of the compounds of transition elements are paramagnetic. Magnetic properties are related to the electronic configuration of atoms. We have already learnt in XI STD that the electron is spinning around its own axis, in addition to its orbital motion around the nucleus. Due to these motions, a tiny magnetic field is generated and it is measured in terms of magnetic moment. On the basis of magnetic properties, materials can be broadly classified as

• Paramagnetic Materials
• Diamagnetic materials, besides these there are ferromagnetic and antiferromagnetic materials.

Materials with no elementary magnetic dipoles are diamagnetic, in other words a species with all paired electrons exhibits diamagnetism. This kind of materials are repelled by the magnetic field because the presence of external magnetic field, a magnetic induction is introduced to the material which generates weak magnetic field that oppose the applied field.

Paramagnetic solids having unpaired electrons possess magnetic dipoles which are isolated from one another. In the absence of external magnetic field, the dipoles are arranged at random and hence the solid shows no net magnetism. But in the presence of magnetic field, the dipoles are aligned parallel to the direction of the applied field and therefore, they are attracted by an external magnetic field.

Ferromagnetic materials have domain structure and in each domain the magnetic dipoles are arranged. But the spin dipoles of the adjacent domains are randomly oriented. Some transition elements or ions with unpaired d electrons show ferromagnetism.

3d transition metal ions in paramagnetic solids often have a magnetic dipole moments corresponding to the electron spin contribution only. The orbital moment L is said to be quenched. So the magnetic moment of the ion is given by µ = g\(\sqrt{S(S+1)}\) µ_B

Where S is the total spin quantum number of the unpaired electrons and is µ_B Bohr Magneton. For an ion with n unpaired electrons S = \(\frac{n}{2}\) and for an electron g = 2.

Therefore the spin only magnetic moment is given by

The magnetic moment calculated using the above equation is compared with the experimental values in the following table. In most of the cases, the agreement is good.

Catalytic Properties

The chemical industries manufacture a number of products such as polymers, flavours, drugs etc., Most of the manufacturing processes have adverse effect on the environment so there is an interest for eco friendly alternatives. In this context, catalyst based manufacturing processes are advantageous, as they require low energy, minimize waste production and enhance the conversion of reactants to products.

Many industrial processes use transition metals or their compounds as catalysts. Transition metal has energetically available d orbitals that can accept electrons from reactant molecule or metal can form bond with reactant molecule using its d electrons. For example, in the catalytic hydrogenation of an alkene, the alkene bonds to an active site by using its π electrons with an empty d orbital of the catalyst.

The σ bond in the hydrogen molecule breaks, and each hydrogen atom forms a bond with a d electron on an atom in the catalyst. The two hydrogen atoms then bond with the partially broken π – bond in the alkene to form an alkane.

In certain catalytic processes the variable oxidation states of transition metals find applications. For example, in the manufacture of sulphuric acid from SO₃, vanadium pentoxide (V₂O₅) is used as a catalyst to oxidise SO₂. In this reaction V₂O₅ is reduced to vanadium (IV) Oxide (VO₂).

Some more examples are discussed below,

(i) Hydroformylation of Olefins

(ii) Preparation acetic acid from acetaldehyde.

(iii) Zeigler – Natta Catalyst

A mixture of TiCl₄ and trialkyl aluminium is used for polymerization.

Alloy Formation

An alloy is formed by blending a metal with one or more other elements. The elements may be metals or non metals or both. The bulk metal is named as solvent, and the other elements in smaller portions are called solute. According to Hume-Rothery rule to form a substitute alloy the difference between the atomic radii of solvent and solute is less than 15%.

Both the solvent and solute must have the same crystal structure and valence and their electro negativity difference must be close to zero. Transition metals satisfying these mentioned conditions form a number of alloys among themselves, since their atomic sizes are similar and one metal atom can be easily replaced by another metal atom from its crystal lattice to form an alloy. The alloys so formed are hard and often have high melting points. Examples: Ferrous alloys, gold – copper alloy, chrome alloys etc.

Formation of Interstitial Compounds

An interstitial compound or alloy is a compound that is formed when small atoms like hydrogen, boron, carbon or nitrogen are trapped in the interstitial holes in a metal lattice. They are usually non-stoichiometric compounds. Transition metals form a number of interstitial compounds such as TiC, ZrH_1.92, Mn₄N etc. The elements that occupy the metal lattice provide them new properties.

• They are hard and show electrical and thermal conductivity
• They have high melting points higher than those of pure metals
• Transition metal hydrides are used as powerful reducing agents
• Metallic carbides are chemically inert

Formation of Complexes

Transition elements have a tendency to form coordination compounds with a species that has an ability to donate an electron pair to form a coordinate covalent bond. Transition metal ions are small and highly charged and they have vacant low energy orbitals to accept an electron pair donated by other groups. Due to these properties, transition metals form large number of complexes. Examples: [Fe(CN)₆]^4-, [Co(NH₃)₆]³⁺, etc. The chemistry of coordination compound is discussed in unit 5.

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Coordination Compounds and Double Salts

When two or more stable compounds in solution are mixed together and allowed to evaporate, in certain cases there is a possibility for the formation of double salts or coordination compounds. For example when an equimolar solution of ferrous sulphate and ammonium sulphate are mixed and allowed to crystallise, a double salt namely Mohr’s salt (Ferrous ammonium sulphate, FeSO₄.(NH₄)₂SO₄.6H₂O) is formed.

Let us recall the blood red colour formation in the inorganic qualitative analysis of ferric ion, the reaction between ferric chloride and potassium thiocyanate solution gives a blood red coloured coordination compound, potassium ferrithiocyanate K₃[Fe(SCN)₆].

If we perform a qualitative analysis to identify the constituent ions present in both the compounds, Mohr’s salt answers the presence of Fe²⁺, NH₄⁺ and SO₄^2- ions, whereas the potassium ferrithiocyanate will not answer Fe³⁺ and SCN^–ions. From this we can infer that the double salts lose their identity and dissociates into their constituent simple ions in solutions, whereas the complex ion in coordination compound, does not loose its identity and never dissociate to give simple ions.

Double salts and coordination compounds are complex compounds. The difference between double salt and coordination compound is that a double salt contains two salts with different crystal structures whereas a coordination compound contains a central metal ion surrounded by molecules or ions known as ligands.

A salt is essentially composed of an anion and a cation. But the main difference between a double salt and a complex salt is that a double salt is a combination of two salt compounds whereas a complex salt is a molecular structure that is composed of one or more complex ions.

A complex salt is a compound composed of a central metal atom having coordination bonds with ligands around it. Do not completely dissociate into its ions in water. It cannot be easily analyzed by determining the ions in the aqueous solution.

Both double salt as well sas complexes are formed by the combination of two or more stabel compounds in stoichiometic reatio. However they differ in the fact that double salt disssociate ito simple ions cmpletely with dissolved in water. However complex ions do not simple ions completely.

A double salt is a salt that contains more than one cation or more than one anion. Other examples include potassium sodium tartrate, ammonium iron (II) sulfate (Mohr’s salt), and bromlite. The fluorocarbonates contain fluoride and carbonate anions.

Double salts and coordination compounds are complex compounds. The difference between double salt and coordination compound is that a double salt contains two salts with different crystal structures whereas a coordination compound contains a central metal ion surrounded by molecules or ions known as ligands.

A double salt is formed from a three-component system, comprising two separate salts and water, and at a given temperature this may be represented by a triangular diagram. Phase diagram of a three-component system.

Double salts are addition compounds which lose their identity in aqueous solution whereas complexes which are also addition compounds do not lose their identity in aqueous solution.

In chemistry, a double bond is a covalent bond between two atoms involving four bonding electrons as opposed to two in a single bond. Double bonds occur most commonly between two carbon atoms, for example in alkenes. Other common double bonds are found in azo compounds (N=N), imines (C=N), and sulfoxides (S=O).

A complex salt is a salt that contains one or more complex ions – ions with metal centers and different molecules attached. Complex salts include potassium ferricyanide (used to create dyes and in blueprint paper) and potassium argentocyanide (used in silver plating).

The molecularity of the chemical reaction is equal to the sum of the stochiometric coefficients of the reactants in the chemical equation of the reaction. It is also defined as the number of reactant molecules taking part in a single step of the reaction.

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Electronic Configuration – Detailed Explanation with Examples

We have already learnt in XI STD to write the electronic configuration of the elements using Aufbu principle, Hund’s rule etc. According to Aufbau principle, the electron first fills the 4s orbital before 3d orbital. Therefore filling of 3d orbital starts from Sc, its electronic configuration is [Ar]3d¹4s² and the electrons of successive elements are progressively filled in 3d orbital and the filling of 3d orbital is complete in Zinc, whose electronic configuration is [Ar]3d¹⁰4s².

However, there are two exceptions in the above mentioned progressive filling of 3d orbitals; if there is a chance of acquiring half filled or fully filled 3d sub shell, it is given priority as they are the stable configuration, for example Cr and Cu. The electronic configurations of Cr and Cu are [Ar] 3d⁵4s¹ respectively.

The extra stability of half filled and fully filled d orbitals, as already explained in XI STD, is due to symmetrical distribution of electrons and exchange energy. Note: The extra stability due to symmetrical distribution can also be visualized as follows. When the d orbitals are considered together, they will constitute a sphere.

So the half filled and fully filled configuration leads to complete symmetrical distribution of electron density. On the other hand, an unsymmetrical distribution of electron density as in the case of partially filled configuration will result in building up of a potential difference.

To decrease this and to achieve a tension free state with lower energy, a symmetrical distribution is preferred. With these two exceptions and minor variation in certain individual cases, the general electronic configuration of d – block elements can be written as [Noble gas] (n – 1)d^1-10 ns^1-2

To calculate an electron configuration, divide the periodic table into sections to represent the atomic orbitals, the regions where electrons are contained. Groups one and two are the s-block, three through 12 represent the d-block, 13 to 18 are the p-block and the two rows at the bottom are the f-block.

Electronic configuration, also called electronic structure, the arrangement of electrons in energy levels around an atomic nucleus. In terms of a more refined, quantum-mechanical model, the K-Q shells are subdivided into a set of orbitals (see orbital), each of which can be occupied by no more than a pair of electrons.

There are different orbital shapes (s, p, d, f) Each orbital can only hold 2 electrons max. There is a hierarchy, i.e. s orbitals will be filled before p orbitals which will be filled before d orbitals and so on. (s<p<d<f) (note, this is a general rule but there are exceptions).

The electron configuration is the standard notation used to describe the electronic structure of an atom. When assigning electrons to orbitals, we must follow a set of three rules: the Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule.

If you are given with the atomic number of an element you can find it’s period number and group number. The period number is related to the number of electron occupied shells in the element and the period number is linked to its valence electrons.

There are two main exceptions to electron configuration: chromium and copper. In these cases, a completely full or half full d sub-level is more stable than a partially filled d sub-level, so an electron from the 4s orbital is excited and rises to a 3d orbital.

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Position of D – Block Elements in the Periodic Table

We have already learnt the periodic classification of elements in XI std. the transition metals occupy from group – 3 to group – 12 of the modern periodic table.

Figure 4.1 – Position of d – block elements in the periodic table

D – Block elements composed of 3d series (4th period) Scandium to Zinc (10 elements), 4d series (5^th period) Yttrium to Cadmium (10 elements) and 5d series (6^th period) Lanthanum, Haffinium to mercury. As we know that the group-12 elements Zinc, Cadmium and Mercury do not have partially filled d-orbital either in their elemental state or in their normal oxidation states.

However they are treated as transition elements, because their properties are an extension of the properties of the respective transition elements. As per the IUPAC definition, the seventh period elements, starting from Ac, Rf to Cn also belong to transition metals. All of them are radioactive. Except Actinium; all the remaining elements are synthetically prepared and have very low half life periods.

The d-block elements are found in the middle of the period table. The d-block elements are called transition metals and have valence electrons in d orbital’s. The f-block elements,found in the two rows at the bottom of the periodic table, are called inner transition metals and have valence electrons in the f-orbital’s.

Transition elements are the elements that are found in Groups 3-12 (old groups IIA-IIB) on the periodic table (salmon-colored block in the middle of the table).

The Periodic Table provides a section for each of these groups of orbitals. The 10 electrons of the five d orbitals are filled by the elements found in the dropped central section of the table. This section is referred to as the d block elements, or the transition metals.

The d-block of the periodic table contains the elements of the groups 3-12 in which the d orbitals are progressively filled in each of the four long periods. The f-block consists of elements in which 4 f and 5 f orbitals are progressively filled. They are placed in a separate panel at the bottom of the periodic table.

The d-block elements are called transition elements because they exhibit transitional behaviour between s block and p-block elements. Their properties are transitional between highly reactive metallic elements of s block which are ionic in nature and the elements of p-block which are covalent in nature.

According to Aufbau principle , electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled. Therefore, 3d orbital is higher in energy than 4s. And hence electrons fill up in 4s before filling up in 3d.

The p sublevel has 3 orbitals, so can contain 6 electrons max. The d sublevel has 5 orbitals, so can contain 10 electrons max. And the 4 sublevel has 7 orbitals, so can contain 14 electrons max. In the picture below, the orbitals are represented by the boxes.

D-Block Elements:

Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate shells are called d-block elements. The Importance of d-block Transition Metals. Transition metals, for the most part, are good conductors. They are also malleable, ductile, lustrous, and sliver-white in color. The d-block metals, and some of it’s key alloys, shaped the Bronze Age, Iron Age, and most importantly the steel age.

The p-block elements are found on the right side of the periodic table. They include the boron, carbon, nitrogen, oxygen and flourine families in addition to the noble gases. The noble gases have full p-orbital’s and are nonreactive.

In chemistry and atomic physics, the main group is the group of elements (sometimes called the representative elements) whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements.

Chlorine is in group 17 of periodic table, also called the halogens, and is not found as the element in nature – only as a compound. The most common of these being salt, or sodium chloride, and the potassium compounds sylvite (or potassium chloride) and carnallite (potassium magnesium chloride hexahydrate).

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Group 18 (Inert gases) Elements

Occurrence:

All the noble gases occur in the atmosphere.

Physical Properties:

As we move along the noble gas elements, their atomic radius and boiling point increases from helium to radon. The first ionization energy decreases from helium to radon. Noble gases have the largest ionisation energy compared to any other elements in a given row as they have completely filled orbital in their outer most shell. They are extremely stable and have a small tendency to gain or lose electrons. The common physical properties of the group 18 elements are listed in the Table.

Physical Properties of group 18 Elements

Properties of Inert Gases:

Physical Properties:

Noble gases are monoatomic, odourless, colourless, tasteless, and non-inflammable. They are highly unreactive. They are non-metallic in nature.

Chemical Properties:

Only the xenon and krypton show some chemical reactivity. Xenon fluorides are prepared by direct reaction of xenon and fluorine under different conditions as shown below.

When XeF₆ is heated at 50°C in a sealed quartz vessel it forms XeOF₄.

When the reaction is continued the following reaction takes place.

2XeOF₄ + SiO₂ → 2XeO₂F₂ + SiF₄
2XeO₂F₂ + SiO₂ → 2XeO₃ + SiF₄

On hydrolysis with water vapour XeF₆ gives XeO₃

XeF₆ + 3H₂O → XeO₃ + 6HF

When XeF₆ reacts with 2.5 M NaOH, sodium per xenate is obtained.

2XeF₆ + 16NaOH → Na₄XeO₆ + Xe + O₂ + 12NaF + 8H₂O

Sodium per xenate is very much known for its strong oxidizing property. For example, it oxidises manganese (II) ion into permanganate ion even in the absence of the catalyst.

5XeO₆^4- + 2Mn²⁺ + 14H⁺ → 2MnO₄^– + 5XeO₃ + 7H₂O

Xenon reacts with PtF₆ and gave an orange yellow solid [XePtF₆] and this is insoluble in CCl₄.

Xenon difluoride forms addition compounds XeF₂.2SbF₅ and XeF₂.2TaF₅. Xenon
hexa fluorides forms compound with boron and alkali metals. Eg: XeF₆.BF₃, XeF₆MF, M-alkali metals.

There is some evidence for existence of xenon dichloride XeCl₂.

Krypton form krypton difluoride when an electric discharge is passed through Kr and flourine at 183°C or when gases are irradiated with SbF₅ it forms KrF₂.2SbF₃.

Structures of Compounds of Xenon:

Compound	Hybridaisation	Shape/Structure
XeF	sp3d	Linear
XeF4	sp3d2	Square planar
XeF6	sp3d3	Distorted octahedron
XeOF2	sp3d	T Shaped
XeOF4	sp3d2	Square pyramidal
XeO3	sp3	Pyramidal

Uses of Noble Gases:

The inertness of noble gases is an important feature of their practical uses.

Helium:

1. Helium and oxygen mixture is used by divers in place of air oxygen mixture. This prevents the painful dangerous condition called bends.
2. Helium is used to provide inert atmosphere in electric arc welding of metals.
3. Helium has lowest boiling point hence used in cryogenics (low temperature science)
4. It is much less denser than air and hence used for filing air balloons.

Neon:

Neon is used in advertisement as neon sign and the brilliant red glow is caused by passing electric current through neon gas under low pressure.

Argon:

Argon prevents the oxidation of hot filament and prolongs the life in filament bulbs

Krypton:

Krypton is used in florescent bulbs, flash bulbs etc. Lamps filed with krypton are used in airports as approaching lights as they can penetrate through dense fog.

Xenon:

Xenon is used in florescent bulbs, flash bulbs and lasers. Xenon emits an intense light in discharge tubes instantly. Due to this it is used in high speed electronic flash bulbs used by photographers.

Radon:

Radon is radioactive and used as a source of gamma rays. Radon gas is sealed as small capsules and implanted in the body to destroy malignant i.e. cancer growth.

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Group 17 (Halogen Group) Elements

Chlorine

Occurrence:

The halogens are present in combined form as they are highly reactive. The main source of flourine is fluorspar or fluorite. The other ores of flourine are cryolite, fluroapatite. The main source of chlorine is sodium chloride from sea water. Bromides and iodides also occur in sea water.

Physical Properties:

The common physical properties of the group 17 elements are listed in the table.

Physical Properties of Group 17 Elements

Properties:

Chlorine is highly reactive hence it doesn’t occur free in nature. It is usually distributed as various metal chlorides. The most important chloride is sodium chloride which occurs in sea water.

Preparation:

Chlorine is prepared by the action of conc. sulphuric acid on chlorides in presence of manganese dioxide.

4NaCl + MnO₂ + 4H₂SO₄ → Cl₂ + MnCl₂ + 4NaHSO₄ + 2H₂O

It can also be prepared by oxidising hydrochloric acid using various oxidising agents such as manganese dioxide, lead dioxide, potassium permanganate or dichromate.

PbO₂ + 4HCl → PbCl₂ + 2H₂O + Cl₂
MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂
2KMnO₄ + 16HCl → 2KCl + 2MnCl₂ + 8H₂O + 5Cl₂
K₂Cr₂O₇ + 14HCl → 2KCl + 2CrCl₃ + 7H₂O + 3Cl₂

When bleaching powder is treated with mineral acids chlorine is liberated

CaOCl₂ + 2HCl → CaCl₂ + H₂O + Cl₃
CaOCl₂ + H₂SO₄ → CaSO₄ + H₂O + Cl₂

Manufacture of Chlorine:

Chlorine is manufactured by the electrolysis of brine in electrolytic process or by oxidation of HCl by air in Deacon’s process.

Electrolytic Process:

When a solution of brine (NaCl) is electrolysed, Na⁺ and Cl^– ions are formed. Na⁺ ion reacts with
OH^– ions of water and forms sodium hydroxide. Hydrogen and chlorine are liberated as gases.

Deacon’s Process:

In this process a mixture of air and hydrochloric acid is passed up a chamber containing a number of shelves, pumice stones soaked in cuprous chloride are placed. Hot gases at about 723 K are passed through a jacket that surrounds the chamber.

The chlorine obtained by this method is dilute and is employed for the manufacture of bleaching powder. The catalysed reaction is given below,

Physical Properties:

Chlorine is a greenish yellow gas with a pungent irritating odour. It produces headache when inhaled even in small quantities whereas inhalation of large quantities could be fatal. It is 2.5 times heavier than air. Chlorine is soluble in water and its solution is referred as chlorine water. It deposits greenish yellow crystals of chlorine hydrate (Cl₂.8H₂O). It can be converted into liquid (Boiling point – 34.6°C) and yellow crystalline solid (Melting point – 102°C).

Chemical Properties:

Action with metals and non-metals:

It reacts with metals and non metals to give the corresponding chlorides.

2Na + Cl₂ → 2NaCl
2Fe + 3Cl₂ → 2FeCl₃
2Al + 3Cl₂ → 2AlCl₃
Cu + Cl₂ → CuCl₂
H₂ + Cl₂ → 2HCl; ∆H = – 44kCal
2B + 3Cl₂ → 2BCl₃

P₄ + 6Cl₂ → 4PCl₃
2As + 3Cl₂ → 2AsCl₃
2Sb + 3Cl₂ → 2SbCl₃

Affinity for Hydrogen:

When burnt with turpentine it forms carbon and hydrochloric acid.

C₁₀H₁₆ + 8Cl₂ → 10C + 16HCl

It forms dioxygen when reacting with water in presence of sunlight. When chlorine in water is exposed to sunlight it loses its colour and smell as the chlorine is converted into hydrochloric acid.

2Cl₂ + 2H₂O → O₂ + 4HCl

Chlorine reacts with ammonia to give ammonium chloride and other products as shown below:
With excess ammonia,

2NH₃ + 3Cl₂ → N₂ + 6HCl
6HCl + 6H₃ → 6NH₄Cl

Overall Reaction
8NH₃ + 3Cl₂ → N₂ + 6NH₄Cl

With excess chlorine,
NH₃ + 3Cl₂ → NCl₃ + 3HCl
3HCl + 3NH₃ → 3NH₄Cl

Overall Reaction
4NH₃ + 3Cl₂ → NCl₃ + 3HCl
3HCl + 3HCl₃ → 3NH₄Cl

Chlorine oxidises hydrogen sulphide to sulphur and liberates bromine and iodine from iodides and bromides. However, it doesn’t oxidise fluorides.

H₂S + Cl₂ → 2HCl + S
Cl₂ + 2KBr → 2KCl + Br₂
Cl₂ + 2KI → 2KCl + I₂

Reaction with Alkali:

Chlorine reacts with cold dilute alkali to give chloride and hypochlorite while with hot concentrated alkali chlorides and chlorates are formed.

Cl₂ + H₂O → HCl + HOCl
HCl + NaOH → NaCl + H₂O
HOCl + NaOH → NaOCl + H₂O

(Cl₂ + H₂O → HCl + HOCl) × 3
(HCl + NaOH → NaCl + H₂O) × 3
(HOCl + NaOH → NaOCl + H₂O) × 3
3NaOCl → NaClO₃ + 2NaCl

Oxidising and Bleaching Action:

Chlorine is a strong oxidising and bleaching agent because of the nascent oxygen.

Colouring matter + Nascent oxygen → Colourless oxidation product

The bleaching of chlorine is permanent. It oxidises ferrous salts to ferric, sulphites to sulphates and hydrogen sulphide to sulphur.

2FeCl₂ + Cl₂ → 2FeCl₃
Cl₂ + H₂O → HCl + HOCl
2FeSO₄ + H₂SO₄ + HOCl → Fe₂(SO₄)₃ + HCl + H₂O

Cl₂ + H₂O → HCl + HOCl
Na₂SO₃ + HOCl → Na₂SO₄ + HCl

Overall Reaction
Na₂SO₃ + H₂O + Cl₂ → CaOCl₂ + H₂O

Preparation of Bleaching Powder:

Bleaching powder is produced by passing chlorine gas through dry slaked lime (calcium hydroxide).
Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O

Displacement Redox Reactions:

Chlorine displaces bromine from bromides and iodine from iodide salts.

Cl₂ + 2KBr → 2KCl + Br₂
Cl₂ + 2KI → 2KCl + I₂

Formation of Addition Compounds:

Chlorine forms addition products with sulphur dioxide, carbon monoixde and ethylene. It forms substituted products with alkanes/arenes.

Uses of Chlorine:

It is used in

1. Purification of drinking water
2. Bleaching of cotton textiles, paper and rayon
3. Extraction of gold and platinum

Hydrochloric Acid:

Laboratory Preparation:

It is prepared by the action of sodium chloride and concentrated sulphuric acid.

NaCl + H₂SO₄ → NaHSO₄ + HCl
NaHSO₄ + NaCl → Na₂SO₄ + HCl

Dry hydrochloric acid is obtained by passing the gas through conc. sulphuric acid.

Properties:

Hydrogen chloride is a colourless, pungent smelling gas, easily liquefied to a colourless liquid (boiling point 189K) and frozen into a white crystalline solid (melting point 159K). It is extremely soluble in water.

HCl (g) + H₂O (l) → H₃O⁺ + Cl^–

Chemical Properties:

Like all acids it liberates hydrogen gas from metals and carbon dioxide from carbonate and bicarbonate salts.

Zn + 2HCl → ZnCl₂ + H₂
Mg + 2HCl → MgCl₂ + H₂
Na₂CO₃ + 2HCl → 2NaCl + CO₂ + H₂
CaCO₃ + 2HCl → CaCl₂ + CO₂ + H₂
NaHCO₃ + 2HCl → 2NaCl + CO₂ + H₂O

It liberates sulphur dioxide from sodium sulphite

Na₂SO₃ + 2HCl → 2NaCl + H₂O + SO₂

When three parts of concentrated hydrochloric acid and one part of concentrated nitric acid are mixed, Aquaregia (Royal water) is obtained. This is used for dissolving gold, platinum etc.

Au + 4H⁺ + NO₃^– + 4Cl^– → AuCl^–₄ + NO + 2H₂O
3Pt + 16H⁺ + 4NO₃^– + 18Cl^– → 3[PtCl₆]^2- + 4NO + 8H₂O

Uses of Hydrochloric Acid:

1. Hydrochloric acid is used for the manufacture of chlorine, ammonium chloride, glucose from corn starch etc.,
2. It is used in the extraction of glue from bone and also for purification of bone black

Trends in Physical and Chemical Properties of Hydrogen Halides:

Preparation:

Direct combination is a useful means of preparing hydrogen chloride. The reaction between hydrogen and flourine is violent while the reaction between hydrogen and bromine or hydrogen and iodine are reversible and don’t produce pure forms.

Displacement Reactions:

Concentrated sulphuric acid displaces hydrogen chloride from ionic chlorides. At higher temperatures the hydrogen sulphate formed react with further ionic chloride. Displacement can be used for the preparation of hydrogen flourides from ionic flourides. Hydrogen bromide and hydrogen iodide are oxidised by concentrated sulphuric acid and can’t be prepared in this method.

Hydrolysis of Phosphorus Trihalides:

Gaseous hydrogen halides are produced when water is added in drops to phosphorus tri halides except phosphorus triflouride.

PX₃ + 3H₂O → H₃PO₃ + 3HX

Hydrogen bromide may be obtained by adding bromine dropwise to a paste of red phosphorous and water while hydrogen iodide is conveniently produced by adding water dropwise to a mixture of red phosphorous and iodine.

2P + 3X₂ → 2PX₃
2PX₃ + 3H₂O → H₃PO₃ + 3HX (where X = Br or I)

Any halogen vapours which escapes with the hydrogen halide is removed by passing the gases through a column of moist red phosphorous.

From Covalent Hydrides:

Halogens are reduced to hydrogen halides by hydrogen sulphide.

H₂S + X₂ → 2HX + S

Hydrogen chloride is obtained as a by-product of the reactions between hydrocarbon of halogens.

General Properties:

In line with the decreasing bond dissociation enthalpy, the thermal stability of hydrogen halides decreases from flouride to iodide. For example, Hydrogen iodide decomposes at 400° C while hydrogen flouride and hydrogen chloride are stable at this temperature.

At room temperature, hydrogen halides are gases but hydrogen flouride can be readily liquefied. The gases are colourless but, with moist air gives white fumes due to the production of droplets of hydrohalic acid. In HF, due to the presence of strong hydrogen bond it has high melting and boiling points. This effect is absent in other hydrogen halides.

	HF	HCl	HBr	HI
Bond dissociation enthalphy (kJ mol-1)	+562	+431	+366	+299
% of ionic character	43	17	13	7

Acidic Properties:

The hydrogen halides are extremely soluble in water due to the ionisation.

HX + H₂O → H₃O⁺ + X^–
(X – F, Cl, Br, or I)

Solutions of hydrogen halides are therefore acidic and known as hydrohalic acids. Hydrochloric, hydrobromic and hydroiodic acids are almost completely ionised and are therefore strong acids but HF is a weak acid i.e. 0.1mM solution is only 10% ionised, but in 5M and 15M solution HF is stronger acid due to the equilibrium.

HF + H₂O ⇄ H₃O⁺ + F^–
HF + F^– ⇄ HF^–₂

At high concentration, the equilibrium involves the removal of flouride ions is important. Since it affects the dissociation of hydrogen flouride and increases and hydrogen ion concentration Several stable salts NaHF₂, KHF₂ and NH₄HF₂ are known. The other hydrogen halides do not form hydrogen dihalides.

Hydrohalic acid shows typical acidic properties. They form salts with acids, bases and reacts with metals to give hydrogen. Moist hydroflouric acid (not dry) rapidly react with silica and glass.

SiO₂ + 4HF → SiF₄ + 2H₂O
Na₂SiO₃ + 6HF → Na₂SiF₆ + 3H₂O

Oxidation:

Hydrogen iodide is readily oxidised to iodine hence it is a reducing agent.

2HI ⇄ 2H⁺ + I₂ + 2e^–

Acidic solution of iodides is readily oxidised. A positive result is shown by liberation of iodine which gives a blue-black colouration with starch.

Hydrogen bromide is more difficult to oxidise than HI. HBr reduces slowly H₂SO₄ into SO₂

2HBr + H₂SO₄ → 2H₂O + Br₂ + SO₂

But hydrogen iodide and ionic iodides are rapidly reduced by H₂SO₄ into H₂S and not into SO₂.

8HI + H₂SO₄ → 4H₂O + 4I₂ + H₂S

Reducing property of hydrogen iodide can be also explained by using its reaction with alcohols into ethane. It converts nitric acid into nitrous acid and dinitrogen dioxide into ammonium.

Hydrogen chloride is unaffected by concentrated sulphuric acid but affected by only strong oxidising agents like MnO₂, potassium permanganate or potassium chloride.

To summarize the trend,

Property	Order
Reactivity of hydrogen	Decreases from fluorine to iodine
Stability	Decreases from HF to HI
Volatility of the hydrides	HF < HI < HBr < HCl
Thermal Stability	HF > HI > HBr > HCl
Boiling Point	HCl < HBr < HI
Acid strength	Increases from HF to HI

Inter Halogen Compounds:

Each halogen combines with other halogens to form a series of compounds called inter halogen compounds. In the given table of inter halogen compounds a given compound A is less electronegative than B.

Properties of Inter Halogen Compounds:

1. The central atom will be the larger one
2. It can be formed only between two halogen and not more than two halogens
3. Fluorine can’t act as a central metal atom being the smallest one
4. Due to high electronegativity with small size florine helps the central atom to attain high coordination number
5. They can undergo the auto ionization.
6. They are strong oxidising agents

2 ICI ⇄ I⁺ + ICI^–₂
2 ICI ⇄ ICI⁺₂ + ICI₄^–

Reaction with Alkali:

When heated with the alkalis, larger halogen form oxyhalogens and the smaller forms halide.

Structure of Inter Halogen Compounds:

The structures of different type of interhalogen compunds can be easily explained using VSEPR theory. The details are given below.

Oxides of Halogen

Fluorine reacts readily with oxygen and forms difluorine oxide (F₂O) and diflourine dioxide (F₂O₂) where
it has a – 1 oxidation state. Other halogens do not react with oxygen readily. But the following oxides can be prepared by some indirect methods. Except flourine all the other halogens have positive oxidation states.

Oxoacids of Halogens:

Chlorine forms four types of oxoacids namely hypochlorus acid, chlorous acid, chloric acid and perchloric acid. Bromine and iodine forms the similar acids except halous acid. However, flourine only forms hypofulric acid. The oxidizing power oxo acids follows the order:

HOX > HXO₃ > HXO₃ > HXO₄

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Oxygen Preparation and its Properties

Preparation

The atmosphere and water contain 23% and 83% by mass of oxygen respectively. Most of the world’s rock contain combined oxygen. Industrially oxygen is obtained by fractional distillation of liquefied air. In the laboratory, oxygen is prepared by one of the following methods.

The decomposition of hydrogen peroxide in the presence of catalyst (MnO₂) or by oxidation with potassium permanganate.

2H₂O₂ ⇄ 2H₂O + O₂
5H₂O₂ + 2MnO₄^– + 6H⁺ → 5O₂ + 8H₂O + 2Mn²⁺

The thermal decomposition of certain metallic oxides or oxoanions gives oxygen.

Properties

Under ordinary condition oxygen exists as a diatomic gas. Oxygen is paramagnetic. Like nitrogen and flourine, oxygen form strong hydrogen bonds. Oxygen exists in two allotropic forms namely dioxygen (O₂) and ozone or trioxygen (O₃).

Although negligible amounts of ozone occurs at sea level it is formed in the upper atmosphere by the action of ultraviolet light. In the laboratory ozone is prepared by passing electrical discharge through oxygen. At a potential of 20,000 V about 10% of oxygen is converted into ozone it gives a mixture known as ozonised oxygen. Pure ozone is obtained as a pale blue gas by the fractional distillation of liquefied ozonised oxygen.

The ozone molecule has a bent shape and symmetrical with delocalised bonding between the oxygen atoms.

Chemical Properties

The chemical properties of oxygen and ozone differ vastly. Oxygen combines with many metals and non-metals to form oxides. With some elements such as s-block elements combination of oxygen occurs at room temperature.

Some of less reactive metals react when powdered finely and made to react exothermically with oxygen at room temperature but a lump of metal is unaffected under same condition. These finely divided metals are known as pyrophoric and when set the powder on fie, heat is liberated during a reaction.

On the other hand ozone is a powerful oxidising agent and it reacts with many substances under conditions where oxygen will not react. For example, it oxidises potassium iodide to iodine. This reaction is quantitative and can be used for estimation of ozone.

O₃ + 2KI + H₂O → 2KOH + O₂I₂

Ozone is commonly used for oxidation of organic compounds. In acidic solution ozone exceeds the oxidising power of flourine and atomic oxygen. The rate of decomposition of ozone drops sharply in alkaline solution.

Uses:

1. Oxygen is one of the essential component for the survival of living organisms.
2. It is used in welding (oxyacetylene welding)
3. Liquid oxygen is used as fuel in rockets etc

Allotrophic forms of Sulphur

Sulphur exists in crystalline as well as amorphous allotrophic forms. The crystalline form includes rhombic sulphur (α sulphur) and monoclinic sulphur (β sulphur). Amorphous allotropic form includes plastic sulphur (γ sulphur), milk of sulphur and colloidal sulphur.

Rhombic sulphur also known as α sulphur, is the only thermodynamically stable allotropic form at ordinary temperature and pressure. The crystals have a characteristic yellow colour and composed of S₈ molecules. When heated slowly above 96°C, it converts into monoclinic sulphur.

Upon cooling below 96°C the β form converts back to α form. Monoclinic sulphur also contains S₈ molecules in addition to small amount of S₆ molecules. It exists as a long needle like prism and is also called as prismatic sulphur. It is stable between 96°C – 119°C and slowly changes into rhombic sulphur.

When molten sulphur is poured into cold water a yellow rubbery ribbon of plastic sulphur is produced. They are very soft and can be stretched easily. On standing (cooling slowly) it slowly becomes hard and changes to stable rhombic sulphur.

Sulphur also exists in liquid and gaseous states. At around 140°C the monoclinic sulphur melts to form mobile pale yellow liquid called λ sulphur. The vapour over the liquid sulphur consists of 90% of S₈, S₇ & S₆ and small amount of mixture of S₂, S₃, S₄, S₅ molecules.

Sulphur Dioxide

Preparation

From Sulphur

A large-scale production of sulphur dioxide is done by burning sulphur in air. About 6-8% of sulphur is oxidised to SO₃.

S + O₂ → SO₂
2S + 3O₂ → 2SO₃

From Sulphides

When sulphide ores such as galena (PbS), zinc blende (ZnS) are roasted in air, sulphur dioxide is liberated. Large amounts of sulphur dioxide required for manufacturing of sulphuric acid and other industrial purpose is prepared by this method.

Laboratory Preparation

Sulphur dioxide is prepared in the laboratory treating a metal or metal sulphite with sulphuric acid

Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O
SO₃^2- + 2H⁺ → H₂O + SO₂

Properties

Sulphur dioxide gas is found in volcanic eruptions. A large amount of sulphur dioxide gas is released into atmosphere from power plants using coal and oil and copper melting plants. It is a colourless gas with a suffocating odour. It is highly soluble in water and it is 2.2 times heavier than air. Sulphur dioxide can be liquefied (boiling point 263 K) at 2.5 atmospheric pressure and 288 K.

Chemical Properties

Sulphur dioxide is an acidic oxide. It dissolves in water to give sulphurous acid.

Reaction with Sodium Hydroxide and Sodium Carbonate

Sulphur dioxide reacts with sodium hydroxide and sodium carbonate to form sodium bisulphite and sodium sulphite respectively.

Oxidising Property

Sulphur dioxide, oxidises hydrogen sulphide to sulphur and magnesium to magnesium oxide.

2H₂S + SO₂ → 3S + 2H₂O
2Mg + SO₂ → 2MgO + S

Reducing Property

As it can readily be oxidised, it acts as a reducing agent. It reduces chlorine into hydrochloric acid.
SO₂ + 2H₂O + Cl₂ → H₂SO₄ + 2HCl

It also reduces potassium permanganate and dichromate to Mn²⁺ and Cr³⁺ respectively.

2KMnO₄ + 5SO₂ + 2H₂O → K₂SO₄ + 2MnSO₄ + 2H₂SO₄
K₂Cr₂O₇ + 3SO₂ + H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + H₂O

Reaction with Oxygen

Sulphur dioxide is oxidised to sulphur trioxide upon heating with oxygen at high temperature. This reaction is used for the manufacture of sulphuric acid by contact process.

Bleaching Action of Sulphur Dioxide

In presence of water, sulphur dioxide bleaches coloured wool, silk, sponges and straw into colourless due to its reducing property.

However, the bleached product (colourless) is allowed to stand in air, it is reoxidised by atmospheric oxygen to its original colour. Hence bleaching action of sulphur dioxide is temporary.

Uses:

1. Sulphur dioxide is used in bleaching hair, silk, wool etc.
2. It can be used for disinfecting crops and plants in agriculture.

Structure of Sulphur Dioxide

In sulphur dioxide, sulphur atom undergoes sp² hybridisation. A double bond arises between S and O is due to pπ – dπ overlapping.

Sulphuric Acid: (H₂SO₄)

Preparation:

Sulphuric acid can be manufactured by lead chamber process, cascade process or contact process. Here we discuss the contact process.

Manufacture of Sulphuric Acid by Contact Process:

The contact process involves the following steps.

(i) Initially sulphur dioxide is produced by burning sulphur or iron pyrites in oxygen/air.

S + O₂ → SO₂
4FeS₂ + 11O₂ → 2Fe₂O₃ + 8SO₂

(ii) Sulphur dioxide formed is oxidised to sulphur trioxide by air in the presence of a catalyst such as V₂O₅ or
platinised asbestos.

(iii) The sulphur trioxide is absorbed in concentrated sulphuric acid and produces oleum (H₂S₂O₇). The oleum is converted into sulphuric acid by diluting it with water.

To maximise the yield the plant is operated at 2 bar pressure and 720 K. The sulphuric acid obtained in this process is over 96% pure.

Physical Properties:

Pure sulphuric acid is a colourless, viscous liquid (Density: 1.84 g/mL at 298 K). High boiling point and viscosity of sulphuric acid is due to the association of molecules together through hydrogen bonding.

The acid freezes at 283.4 K and boils at 590 K. It is highly soluble in water and has strong affinity towards water and hence it can be used as a dehydrating agent. When dissolved in water, it forms mono (H₂SO₄.H₂O) and dihydrates (H₂SO₄.2H₂O) and the reaction is exothermic.

The dehydrating property can also be illustrated by its reaction with organic compounds such as sugar, oxalic acid and formic acid.

Chemical Properties:

Sulphuric acid is highly reactive. It can act as strong acid and an oxidising agent.

Decomposition:

Sulphuric acid is stable, however, it decomposes at high temperatures to sulphur trioxide.

H₂SO₄ → H₂O + SO₃

Acidic Nature:

It is a strong dibasic acid. Hence it forms two types of salts namely sulphates and bisulphates.

Oxidising Property:

Sulphuric acid is an oxidising agent as it produces nascent oxygen as shown below.

Sulphuric acid oxidises elements such as carbon, sulphur and phosphorus. It also oxidises bromide and iodide to bromine and iodine respectively.

C + 2H₂SO₄ → 2SO₂ + 2H₂O + CO₂
S + 2H₂SO₄ → 3SO₂ + 2H₂O
P₄ + 10H₂SO₄ → 4H₃PO₄ + 10SO₂ + 4H₂O
H₂S + H₂SO₄ → SO₂ + 2H₂O + S
H₂SO₄ + 2HI → SO₂ + 2H₂O + I₂
H₂SO₄ + 2HBr → SO₂ + 2H₂O + Br₂

Reaction with Metals:

Sulphuric acid reacts with metals and gives different product depending on the reactants and reacting condition.

Dilute sulphuric acid reacts with metals like tin, aluminium, zinc to give corresponding sulphates.

Zn + H₂SO₄ → ZnSO₄ + H₂↑
2Al + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂↑

Hot concentrated sulphuric acid reacts with copper and lead to give the respective sulphates as shown below.

Cu + 2H₂SO₄ → CuSO₄ + 2H₂O + SO₂↑
Pb + 2H₂SO₄ → PbSO₄ + 2H₂O + SO₂↑

Sulphuric acid doesn’t react with noble metals like gold, silver and platinum.

Reaction with Salts:

It reacts with different metal salts to give metal sulphates and bisulphates.

KCl + H₂SO₄ → KHSO₄ + HCl
KNO₃ + H₂SO₄ → KHSO₄ + HNO₃
Na₂CO₃ + H₂SO₄ → Na₂SO₄ + H₂O + CO₂
2NaBr + 3H₂SO₄ → 2NaHSO₄ + 2H₂O + Br₂ + SO₂

Reaction with Organic Compounds:

It reacts organic compounds such as benzene to give sulphonic acids.

Uses of Sulphuric Acid:

1. Sulphuric acid is used in the manufacture of fertilisers, ammonium sulphate and super phosphates and other chemicals such as hydrochloric acid, nitric acid etc.
2. It is used as a drying agent and also used in the preparation of pigments, explosives etc.

Test for Sulphate/Sulphuric Acid:

Dilute solution of sulphuric acid/aqueous solution of sulphates gives white precipitate (barium sulphate) with barium chloride solution. It can also be detected using lead acetate solution. Here a white precipitate of lead sulphate is obtained.

Structure of Oxoacids of Sulphur:

Sulphur forms many oxoacids. The most important one is sulphuric acid. Some acids like sulphurous and dithionic acids are known in the form of their salts only since the free acids are unstable and cannot be isolated. Various oxo acids of sulphur with their structures are given below.

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Group 15 (Nitrogen Group) Elements

Occurrence:

About 78% of earth atmosphere contains dinitorgen (N₂) gas. It is also present in earth crust as sodium nitrate (Chile saltpetre) and potassium nitrate (Indian saltpetre). The 11^th most abundant element phosphorus, exists as phosphate (fluroapatite, chloroapatite and hydroxyapatite). The other elements arsenic, antimony and bismuth are present as sulphides and are not very abundant.

Physical Properties:

Some of the physical properties of the group 15 elements are listed below

Table 3.1 Physical Properties of Group 15 Elements

Nitrogen:

Preparation:

Nitrogen, the principal gas of atmosphere (78% by volume) is separated industrially from liquid air by fractional distillation

Pure nitrogen gas can be obtained by the thermal decomposition of sodium azide about 575 K

2NaN₃ → 2 Na + 3Na₂

It can also be obtained by oxidising ammonia using bromine water

8NH₃ + 3Br₂ → 6NH₄Br + N₂

Properties

Nitrogen gas is rather inert. Terrestrial nitrogen contains 14.5% and 0.4% of nitrogen-14 and nitrogen-15 respectively. The later is used for isotopic labelling. The chemically inert character of nitrogen is largely due to high bonding energy of the molecules 225 cal mol^-1 (946 kJ mol^-1).

Interestingly the triply bonded species is notable for its less reactivity in comparison with other iso-electronic triply bonded systems such as -C≡C-, C≡O, X-C≡N, X-N≡C, -C≡C-, and -C≡N. These groups can act as donor where as dinitrogen cannot. However, it can form complexes with metal (M ← N≡N) like CO to a less extent.

The only reaction of nitrogen at room temperature is with lithium forming Li₃N. With other elements, nitrogen combines only at elevated temperatures. Group 2 metals and T forms ionic nitrides.

6Li + N₂ → 2Li₃N

Direct reaction with hydrogen gives ammonia. This reaction is favoured by high pressures and at optimum temperature in presence of iron catalyst. This reaction is the basis of Haber’s process for the synthesis of ammonia.

\(\frac{1}{2}\)N₂ + \(\frac{3}{2}\)H₂ ⇄ NH₃
∆H_f = -46.2 kJ mol^-1

With oxygen, nitrogen produces nitrous oxide at high temperatures. Even at 3473 K nitrous oxide yield is only 4.4%.

2N₂ + O₂ → 2N₂O

Uses of Nitrogen

1. Nitrogen is used for the manufacture of ammonia, nitric acid and calcium cyanamide etc.
2. Liquid nitrogen is used for producing low temperature required in cryosurgery, and so in biological preservation.

Ammonia (NH₃)

Preparation:

Ammonia is formed by the hydrolysis of urea.
NH₂CONH₂ + H₂O → 2NH₃ + CO₂

Ammonia is prepared in the laboratory by heating an ammonium salt with a base.

2NH₄⁺ + OH^– → 2NH₃ + H₂O
2NH₄Cl + CaO → CaCl₂ + 2NH₃ + H₂O

It can also be prepared by heating a metal nitrides such as magnesium nitride with water.

Mg₃N₂ + 6H₂O → 3Mg(OH)₂ + 2NH₃

It is industrially manufactured by passing nitrogen and hydrogen over iron catalyst a small amount of K₂O and Al₂O₃ is also used to increase the rate of attainment of equilibrium) at 750 K at 200 atm pressure. In the actual process the hydrogen required is obtained from water gas and nitrogen from fractional distillation of liquid air.

Properties

Ammonia is a pungent smelling gas and is lighter than air. It can readily liquefied by at about 9 atmospheric pressure. The liquid boils at – 38.4°C and freezes at – 77° C. Liquid ammonia resembles water in its physical properties. i.e. it is highly associated through strong hydrogen bonding. Ammonia is extremely soluble in water (702 Volume in 1 Volume of water) at 20°C and 760mm pressure.

At low temperatures two soluble hydrate NH₃.H₂O and 2NH₃.H₂O are isolated. In these molecules ammonia and water are linked by hydrogen bonds. In aqueous solutions also ammonia may be hydrated in a similar manner and we call the same as (NH₃.H₂O)

NH₃ + H₂O ⇄ NH₄⁺ + OH^–

The dielectric constant of ammonia is considerably high to make it a fairly good ionising solvent like water.

Chemical Properties

Action of Heat:

Above 500°C ammonia decomposes into its elements. The decomposition may be accelerated by metallic catalysts like Nickel, Iron. Almost complete dissociation occurs on continuous sparking.

Reaction with Air/Oxygen:

Ammonia does not burn in air but burns freely in free oxygen with a yellowish flame to give nitrogen and steam.

4NH₃ + 3O₂ ⇄ N₂ + 6H₂O

In presence of catalyst like platinum, it burns to produce nitric oxide. This process is used for the manufacture of nitric acid and is known as ostwalds process.

4NH₃ + 5O₂ ⇄ 4NO + 6H₂O

Reducing Property:

Ammonia acts as a reducing agent. It reduces the metal oxides to metal when passed over heated metallic oxide.

3PbO + 2NH₃ → 3Pb + N₂ + 3H₂O

Reaction with Acids:

When treated with acids it forms ammonium salts. This reaction shows that the affinity of ammonia for proton is greater than that of water.

Reaction with Chlorine and Chlorides:

Ammonia reacts with chlorine and chlorides to give ammonium chloride as a final product. The reactions are different under different conditions as given below.

With excess ammonia

2 NH₃ + 3 Cl₂ → N₂ + 6 HCl
6 HCl + 6 NH₃ → 6 NH₄Cl

With excess of chlorine ammonia reacts to give nitrogen trichloride, an explosive substance.

2NH₃ + 6Cl₂ → 2NCl₃ + 6HCl
2NH₃(g) + HCl(g) → NH₄Cl(s)

Formation of Amides and Nitrides:

With strong electro positive metals such as sodium, ammonia forms amides while it forms nitrides with metals like magnesium.

2Na + 2NH₃ → 2NaNH₂ + H₂
3Mg + 2NH₃ → Mg₃N₂ + 3H₂

With Metallic Salts:

Ammonia reacts with metallic salts to give metal hydroxides (in case of Fe) or forming complexes (in case Cu)

Formation of Amines:

Ammonia forms ammonated compounds by ion dipole attraction. Eg. [CaCl₂.8NH₃]. In this, the negative ends of ammonia dipole is attracted to Ca²⁺ ion.

It can also act as a ligand and form coordination compounds such as [Co(NH₃)₆]³⁺, [Ag(NH₃)₂]⁺. For example when excess ammonia is added to aqueous solution copper sulphate a deep blue colour compound [Cu(NH₃)₄]²⁺ is formed.

Structure of Ammonia

Ammonia molecule is pyramidal in shape N-H bond distance is 1.016 Å and H-H bond distance is 1.645 Å with a bond angle 107°. The structure of ammonia may be regarded as a tetrahedral with one lone pair of electrons in one tetrahedral position hence it has a pyramidal shape as shown in the figure.

Nitric Acid

Preparation

Nitric acid is prepared by heating equal amounts of potassium or sodium nitrate with concentrated sulphuric acid.

KNO₃ + H₂SO₄ → KHSO₄ + HNO₃

The temperature is kept as low as possible to avoid decomposition of nitric acid. The acid condenses to a fuming liquid which is coloured brown by the presence of a little nitrogen dioxide which is formed due to the decomposition of nitric acid.

4HNO₃ → 4NO₂ + 2H₂O + O₂

Commercial Method of Preparation

Nitric acid prepared in large scales using Ostwald’s process. In this method ammonia from Haber’s process is mixed about 10 times of air. This mixture is preheated and passed into the catalyst chamber where they come in contact with platinum gauze.

The temperature rises to about 1275 K and the metallic gauze brings about the rapid catalytic oxidation of ammonia resulting in the formation of NO, which then oxidised to nitrogen dioxide.

4NH₃ + 5O₂ → 4NO + 6H₂O + 120kJ
2NO + O₂ → 2NO₂

The nitrogen dioxide produced is passed through a series of adsorption towers. It reacts with water to give nitric acid. Nitric oxide formed is bleached by blowing air.

3NO₂ + H₂O → 2HNO₃ + NO

Properties

Pure nitric acid is colourless. It boils at 86°C. The acid is completely miscible with water forming a constant boiling mixture (98% HNO₃, Boiling point 120.5 °C). Fuming nitric acid contains oxides of nitrogen. It decomposes on exposure to sunlight or on being heated, into nitrogen dioxide, water and oxygen.

4HNO₃ → 4NO₂ + 2H₂O + O₂

Due to this reaction pure acid or its concentrated solution becomes yellow on standing. In most of the reactions, nitric acid acts as an oxidising agent. Hence the oxidation state changes from + 5 to a lower one. It doesn’t yield hydrogen in its reaction with metals. Nitric acid can act as an acid, an oxidizing agent and an nitrating agent.

As an Acid:

Like other acids it reacts with bases and basic oxides to form salts and water

ZnO + 2HNO₃ → Zn(NO₃)₂ + H₂O
3FeO + 10HNO₃ → 3Fe(NO₃)₃ + NO + 5H₂O

As an Oxidising Agent:

The nonmetals like carbon, sulphur, phosphorus and iodine are oxidised by nitric acid.

C + 4HNO₃ → 2H₂O + 4NO₂ + CO₂
S + 2HNO₃ → H₂SO₄ + 2NO
P₄ + 20HNO₃ → 4H₃PO₄ + 4H₂O + 20NO₂
3I₂ + 10HNO₃ → 6HIO₃ + 10NO + 2H₂O
HNO₃ + F₂ → HF + NO₃F
3H₂S + 2HNO₃ → 3S + 2NO + 4H₂O

As an Nitrating Agent:

In organic compounds replacement of a – H atom with – NO₂ is often referred as nitration. For example.

Nitration takes place due to the formation of nitronium ion

HNO₃ + H₂SO₄ → NO₂⁺ + H₂O + HSO₄^–

Action of Nitric Acid on Metals

All metals with the exception of gold, platinum, rhodium, iridium and tantalum reacts with nitric acid. Nitric acid oxidises the metals. Some metals such as aluminium, iron, cobalt, nickel and chromium are rendered passive in concentrated acid due to the formation of a layer of their oxides on the metal surface, which prevents the nitric acid from reacting with pure metal.

With weak electropositive metals like tin, arsenic, antimony, tungsten and molybdenum, nitric acid gives metal oxides in which the metal is in the higher oxidation state and the acid is reduced to a lower oxidation state. The most common products evolved when nitric acid reacts with a metal are gases NO₂, NO and H₂O. Occasionally N₂, NH₂OH and NH₃ are also formed.

The reactions of metals with nitric acid are explained in 3 steps as follows:

Primary Reaction:
Metal nitrate is formed with the release of nascent hydrogen

M + HNO₃ → MNO₃ + (H)

Secondary Reaction:

Nascent hydrogen produces the reduction products of nitric acid.

Tertiary Reaction:

The secondary products either decompose or react to give final products

Decomposition of the Secondary:

Reaction of Secondary Products:

HNO₂ + NH₃ → N₂ + 2H₂O
HNO₂ + NH₂OH → N₂O + 2H₂O
HNO₂ + HNO₃ → 2NO₂ + H₂O

Examples:

Copper reacts with nitric acid in the following manner

3Cu + 6HNO₃ → 3Cu(NO₃)₂ + 6(H)
6(H) + 3HNO₃ → 3HNO₂ + 3H₂O
3HNO₂ → HNO₃ + 2NO + H₂O

Overall Reaction
3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O

The concentrated acid has a tendency to form nitrogen dioxide

Cu + 4HNO₃ → Cu(NO₃₎₂ + 2NO₂ + 2H₂O

Magnesium reacts with nitric acid in the following way

4Mg + 8HNO₃ → 4Mg(NO₃)₂ + 8[H]
HNO₃ + 8H → NH₃ + 3H₂O
HNO₃ + NH₃ → NH₄NO₃

Overall Reaction
4Mg + 10HNO₃  → 4Mg(NO₃)₂ + N₂O + 5H₂O

If the acid is diluted we get N₂O

4Mg + 10HNO₃ → 4Mg(NO₃)₂ + N₂O + 5H₂O

Uses of Nitric Acid:

1. Nitric acid is used as a oxidising agent and in the preparation of aquaregia.
2. Salts of nitric acid are used in photography (AgNO₃) and gunpowder for firearms. (NaNO₃)

Oxides and Oxoacids of Nitrogen

Structures of Oxides of Nitrogen:

Structures of Oxoacids of Nitrogen:

Preparation of Oxoacids of Nitrogen:

Name Formula Oxidation state Preparation

Allotropic Forms of Phosphorus:

Phosphorus has several allotropic modification of which the three forms namely white, red and black phosphorus are most common. The freshly prepared white phosphorus is colourless but becomes pale yellow due to formation of a layer of red phosphorus upon standing.

Hence it is also known as yellow phosphorus. It is poisonous in nature and has a characteristic garlic smell. It glows in the dark due to oxidation which is called phosphorescence. Its ignition temperature is very low and hence it undergoes spontaneous combustion in air at room temperature to give P₂O₅.

The white phosphorus can be changed into red phosphorus by heating it to 420°C in the absence of air and light. Unlike white phosphorus it is not poisonous and does not show Phosphorescence. It also does not ignite at low temperatures. The red phosphorus can be converted back into white phosphorus by boiling it in an inert atmosphere and condensing the vapour under water.

The phosphorus has a layer structure and also acts as a semiconductor. The four atoms in phosphorus have polymeric structure with chains of P₄ linked tetrahedrally. Unlike nitrogen P≡P is less stable than P-P single bonds. Hence, phosphorus atoms are linked through single bonds rather than triple bonds. In addition to the above two more allotropes namely scarlet and violet phosphorus are also known for phosphorus.

Properties of Phosphorus:

Phosphorus is highly reactive and has the following important chemical properties

Reaction with Oxygen:

Yellow phosphorus readily catches fire in air giving dense white fumes of phosphorus pentoxide. Red phosphorus also reacts with oxygen on heating to give phosphorus trioxide or phosphorus pentoxide.

Reaction with Chlorine:

Phosphorus reacts with chlorine to form tri and penta chloride. Yellow phosphorus reacts violently at room temperature, while red phosphorus reacts on heating.

Reaction with Alkali:

Yellow phosphorus reacts with alkali on boiling in an inert atmosphere liberating phosphine. Here phosphorus act as reducing agent.

Reaction with Nitric Acid:

When phosphorus is treated with conc. nitric acid it is oxidised to phosphoric acid. This reaction is catalysed by iodine crystals.

Reaction with Metals:

Phosphorus reacts with metals like Ca and Mg to give phosphides. Metals like sodium and potassium react with phosphorus vigorously.

Uses of Phosphorus:

1. The red phosphorus is used in the match boxes
2. It is also used for the production of certain alloys such as phosphor bronze

Phosphine (PH₃)

Phosphine is the most important hydride of phosphorus

Preparation:

Phosphine is prepared by action of sodium hydroxide with white phosphorus in an inert atmosphere of carbon dioxide or hydrogen.

Phosphine is freed from phosphine dihydride (P₂H₄) by passing through a freezing mixture. The dihydride
condenses while phosphine does not.

Phosphine can also prepared by the hydrolysis of metallic phosphides with water or dilute mineral acids.

Phosphine is prepared in pure form by heating phosphorous acid.

A pure sample of phosphine is prepared by heating phosphonium iodide with caustic soda solution.

Physical Properties:

It is colourless, poisonous gas with rotten fish smell. It is slightly soluble in water and is neutral to litmus test. It condenses to a colourless liquid at 188 K and freezes to a solid at 139.5 K.

Chemical Properties:

Thermal Stability:

Phosphine decomposes into its elements when heated in absence of air at 317 K or when electric current is passed through it.

Combustion:

When phosphine is heated with air or oxygen it burns to give meta phosphoric acid.

Basic Nature:

Phosphine is weakly basic and forms phosphonium salts with halogen acids.

It reacts with halogens to give phosphorus penta halides.

PH₃ + 4Cl₂ → PCl₅ + 3HCl

Reducing Property:

Phosphine precipitates some metal from their salt solutions.

3AgNO₃ + PH₃ → Ag₃P + 3HNO₃

It forms coordination compounds with lewis acids such as boron trichloride.

Structure:

In phosphine, phosphorus shows sp³ hybridisation. Three orbitals are occupied by bond pair and fourth corner is occupied by lone pair of electrons. Hence, bond angle is reduced to 93.5°. Phosphine has a pyramidal shape.

Uses of Phosphine:

Phosphine is used for producing smoke screen as it gives large smoke. In a ship, a pierced container with a mixture of calcium carbide and calcium phosphide, liberates phosphine and acetylene when thrown into sea. The liberated phosphine catches fire and ignites acetylene. These burning gases serves as a signal to the approaching ships. This is known as Holmes signal.

Phosphorous Trichloride and Pentachloride:

Phosphorous Trichloride:

Preparation:

When a slow stream of chlorine is passed over white phosphorus, phosphorous trichloride is formed. It can also be obtained by treating white phosphorus with thionyl chloride.

P₄ + 8SOCl₂ → 4PCl₃ + 4SO₂ + 2S₂Cl₂

Properties

When phosphorous trichloride is hydrolysed with cold water it gives phosphorous acid.

PCl₃ + 3H₂O → H₃PO₃ + 3HCl

This reaction involves the coordination of a water molecule using a vacant 3d orbital on the phosphorous atom following by elimination of HCl which is similar to hydrolysis of SiCl₄.

PCl₃ + H₂O → PCl₃.H₂O → P(OH)Cl₂ + HCl

This reaction is followed by two more steps to give P(OH)₃ or H₃PO₃.

HPOCl₂ + H₂O → H₂PO₂Cl + HCl
H₂PO₂Cl + H₂O → H₂PHO₃ + HCl

Similar reactions occurs with other molecules that contains alcohols and carboxylic acids.

3C₂H₅OH + PCl₃ → 3C₂H₅Cl + H₃PO₃
3C₂H₅COOH + PCl₃ → 3C₂H₅COCl + H₃PO₄

Uses of Phosphorus Trichloride:

Phosphorus trichloride is used as a chlorinating agent and for the preparation of H₃PO₃.

Phosphorous Pentachloride:

Preparation

When PCl₃ is treated with excess chlorine, phosphorous pentachloride is obtained.

PCl₃ + Cl₂ → PCl₅

Chemical Properties

On heating phosphorous pentachloride, it decomposes into phosphorus trichloride and chlorine.

Phosphorous pentachloride reacts with water to give phosphoryl chloride and orthophosphoric acid.

PCl₅ + H₂O → POCl₃ + 2HCl
POCl₃ + 3H₂O → H₃PO₄ + 3HCl

Overall Reaction
PCl₅ + 4H₂O → H₃PO₄ + 5HCl

Phosphorous pentachloride reacts with metal to give metal chlorides. It also chlorinates organic compounds similar to phosphorus trichloride.

2Ag + PCl₅ → 2 AgCl + PCl₃
Sn + 2PCl₅ → SnCl₄ + 2PCl₃
C₂H₅OH + PCl₅ → C₂H₅Cl + HCl + POCl₃
C₂H₅COOH + PCl₅ → C₂H₅COCl + HCl + POCl₃

Uses of Phosphorus Pentachloride

Phosphorous pentachloride is a chlorinating agent and is useful for replacing hydroxyl groups by chlorine atom.

Structure of Oxides and Oxoacids of Phosphorus

Phosphorous forms phosphorous trioxide, phosphorous tetra oxide and phosphorous pentaoxides. In phosphorous trioxide four phosphorous atoms lie at the corners of a tetrahedron and six oxygen atoms along the edges. The P-O bond distance is 165.6 pm which is shorter than the single bond distance of P-O (184 pm) due to pπ-dπ bonding and results in considerable double bond character.

In P₄O₁₀ each P atoms form three bonds to oxygen atom and also an additional coordinate bond with an
oxygen atom.

Terminal P-O bond length is 143 pm, which is less than the expected single bond distance. This may be due to lateral overlap of filled p orbitals of an oxygen atom with empty d orbital on phosphorous.

Oxoacids of Phosphorous-Structure:

Oxoacids of Phosphorus-Preparation:

Group 16 (Oxygen group) Elements:

Occurrence:

Elements belonging group 16 are called chalgogens or ore forming elements as most of the ores are oxides or sulphides. First element oxygen, the most abundant element, exists in both as dioxygen in air (above 20% by weight as well as volume) and in combined form as oxides.

Oxygen and sulphur makes up about 46.6% & 0.034 & of earth crust by weight respectively. Sulphur exists as sulphates (gypsum, epsom etc…) and sulphide (galena, Zinc blende etc…). It is also present in the volcanic ashes. The other elements of this groups are scarce and are often found as selenides, tellurides etc along with
sulphide ores.

Physical Properties:

The common physical properties of the group 16 elements are listed in the Table.

Table 3.2 Physical Properties of Group 16 Elements

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Group 14 (Carbon Group) Elements

Occurrence:

Carbon is found in the native form as graphite. Coal, crude oil and carbonate rocks such as calcite, magnesite etc contains large quantities of carbon in its combined form with other elements. Silicon occurs as silica (sand and quartz crystal). Silicate minerals and clay are other important sources for silicon.

Physical Properties:

Some of the physical properties of the group 14 elements are listed below

Table 2.4 Physical properties of group 14 elements

Tendency for Catenation

Catenation is an ability of an element to form chain of atoms. The following conditions are necessary for catenation.

• The valency of element is greater than or equal to two
• Element should have an ability to bond with itself
• The self bond must be as strong as its bond with other elements
• Kinetic inertness of catenated compound towards other molecules.

Carbon possesses all the above properties and forms a wide range of compounds with itself and with other elements such as H, O, N, S and halogens.

Allotropes of Carbon

Carbon exists in many allotropic forms. Graphite and diamond are the most common allotropes. Other important allotropes are graphene, fullerenes and carbon nanotubes.

Graphite is the most stable allotropic form of carbon at normal temperature and pressure. It is soft and conducts electricity. It is composed of flat two dimensional sheets of carbon atoms. Each sheet is a hexagonal net of sp² hybridised carbon atoms with a C-C bond length of 1.41 Å which is close to the C-C bond distance in benzene (1.40 Å). Each carbon atom forms three σ bonds with three neighbouring carbon atoms using three of its valence electrons and the fourth electron present in the unhybridised p orbital forms a π-bond.

These π electrons are delocalised over the entire sheet which is responsible for its electrical conductivity. The successive carbon sheets are held together by weak vander Waals forces. The distance between successive sheet is 3.40 Å. It is used as a lubricant either on its own or as a graphited oil.

Unlike graphite the other allotrope diamond is very hard. The carbon atoms in diamond are sp³ hybridised and bonded to four neighbouring carbon atoms by σ bonds with a C-C bond length of 1.54 Å. This results in a tetrahedral arrangement around each carbon atom that extends to the entire lattice as shown in figure 2.5. Since all four valance electrons of carbon are involved in bonding there is no free electrons for conductivity. Being the hardest element, it used for sharpening hard tools, cutting glasses, making bores and rock drilling.

Fullerenes are newly synthesised allotropes of carbon. Unlike graphite and diamond, these allotropes are discrete molecules such as C₃₂, C₅₀, C₆₀, C₇₀, C₇₆ etc. These molecules have cage like structures as shown in the figure. The C₆₀ molecules have a soccer ball like structure and is called buckminster fullerene or buckyballs.

It has a fused ring structure consists of 20 six membered rings and 12 five membered rings. Each carbon atom is sp² hybridised and forms three σ bonds & a delocalised π bond giving aromatic character to these molecules. The C-C bond distance is 1.44 Å and C = C distance 1.38 Å.

Carbon nanotubes, another recently discovered allotropes, have graphite like tubes with fullerene ends. Along the axis, these nanotubes are stronger than steel and conduct electricity. These have many applications in nanoscale electronics, catalysis, polymers and medicine.

Another allotrophic form of carbon is graphene. It has a single planar sheet of sp² hybridised carbon atoms that are densely packed in a honeycomb crystal lattice.

Carbon Monoxide [CO]:

Preparation:

Carbon monoxide can be prepared by the reaction of carbon with limited amount of oxygen.

2C + O₂ → 2CO

On industrial scale carbon monoxide is produced by the reaction of carbon with air. The carbon monoxide formed will contain nitrogen gas also and the mixture of nitrogen and carbon monoxide is called producer gas.

The producer gas is then passed through a solution of copper(I)chloride under pressure which results in the formation of CuCl(CO).2H₂O. At reduced pressures this solution releases the pure carbon monoxide. Pure carbon monoxide is prepared by warming methanoic acid with concentrated sulphuric acid which acts as a dehydrating agent.

HCOOH + H₂SO₄ → CO + H₂SO₄.H₂O

Properties

• It is a colourless, odourless, and poisonous gas. It is slightly soluble in water.
• It burns in air with a blue flame forming carbon dioxide.

2CO + O₂ → 2CO₂

When carbon monoxide is treated with chlorine in presence of light or charcoal, it forms a poisonous gas carbonyl chloride, which is also known as phosgene. It is used in the synthesis of isocyanates.

CO + Cl₂ → COCl₂

Carbon monoxide acts as a strong reducing agent.

3CO + Fe₂ O₃ → 2Fe + 3CO₂

Under high temperature and pressure a mixture of carbon monoxide and hydrogen (synthetic gas or syn gas) gives methanol.

CO + 2H₂ → CH₃OH

In oxo process, ethene is mixed with carbon monoxide and hydrogen gas to produce propanal.

CO + C₂H₄ + H₄ → CH₃CH₂CHO

Fischer Tropsch Synthesis:

The reaction of carbon monoxide with hydrogen at a pressure of less than 50 atm using metal catalysts at 500 – 700 K yields saturated and unsaturated hydrocarbons.

nCO + (2n+1)H₂ → C_nH(_2n+2) + nH₂O
nCO + 2nH₂ → C_nH_2n + nH₂O

Carbon monoxide forms numerous complex compounds with transition metals in which the transition meal is in zero oxidation state. These compounds are obtained by heating the metal with carbon monoxide.

Eg. Nickel tetracarbonyl [Ni(CO)₄], Iron pentacarbonyl [Fe(CO)₅], Chromium hexacarbonyl [Cr(CO)₆].

Structure:

It has a linear structure. In carbon monoxide, three electron pairs are shared between carbon and oxygen. The bonding can be explained using molecular orbital theory as discussed in XI standard. The C-O bond distance is 1.128Å. The structure can be considered as the resonance hybrid of the following two canonical forms.

Uses of Carbon Monoxide:

1. Equimolar mixture of hydrogen and carbon monoxide – water gas and the mixture of carbon monoxide and nitrogen – producer gas are important industrial fuels.
2. Carbon monoxide is a good reducing agent and can reduce many metal oxides to metals.
3. Carbon monoixde is an important ligand and forms carbonyl compound with transition metals.

Carbon Dioxide:

Carbon dioxide occurs in nature in free state as well as in the combined state. It is a constituent of air (0.03%). It occurs in rock as calcium carbonate and magnesium carbonate.

Production

On industrial scale it is produced by burning coke in excess of air.
C + O₂ → CO₂ ∆H = – 394 kJ mol^-1

Calcination of lime produces carbon dioxide as by product.
CaCO₃ → CaO + CO₂

Carbon dioxide is prepared in laboratory by the action of dilute hydrochloric acid on metal carbonates.
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

Properties

It is a colourless, nonflammable gas and is heavier than air. Its critical temperature is 31^° C and can be readily liquefied.

Carbon dioxide is a very stable compound. Even at 3100 K only 76% decomposes to form carbon monoxide and oxygen. At still higher temperature it decomposes into carbon and oxygen.

Oxidising Behaviour:

At elevated temperatures, it acts as a strong oxidising agent. For example,
CO₂ + 2Mg → 2MgO + C

Water Gas Equilibrium:

The equilibrium involved in the reaction between carbon dioxide and hydrogen, has many industrial applications and is called water gas equilibrium.

Acidic Behaviour:

The aqueous solution of carbon dioxide is slightly acidic as it forms carbonic acid.

CO₂ + H₂O ⇄ H₂CO₃ ⇄ H⁺ + HCO₃^–

Structure of Carbon Dioxide

Carbon dioxide has a linear structure with equal bond distance for the both C-O bonds. In this molecule there is two C-O sigma bond. In addition there is 3c-4e bond covering all the three atoms.

Uses of Carbon Dioxide

1. Carbon dioxide is used to produce an inert atomosphere for chemical processing.
2. Biologically, it is important for photosynthesis.
3. It is also used as fie extinguisher and as a propellent gas.
4. It is used in the production of carbonated beverages and in the production of foam.

Silicon Tetrachloride:

Preparation:

Silicon tetrachloride can be prepared by passing dry chlorine over an intimate mixture of silica and carbon by heating to 1675 K in a porcelain tube

SiO₂ + 2C + 2Cl₂ → SiCl₄ + 2CO

On commercial scale, reaction of silicon with hydrogen chloride gas occurs above 600 K

Si + 4HCl → SiCl₄ + 2H₂

Properties:

Silicon tetrachloride is a colourless fuming liquid and it freezes at – 70°C. In moist air, silicon tetrachloride is hydrolysed with water to give silica and hydrochloric acid.

SiCl₄ + 2H₂O → 4HCl + SiO₂

When silicon tetrachloride is hydrolysed with moist ether, linear perchloro siloxanes are formed [Cl-(Si Cl₂O)_nSiCl₃ where n = 1 – 6.

Alcoholysis

The chloride ion in silicon tetrachloride can be substituted by nucleophile such as OH, OR, etc using suitable reagents. For example, it forms silicic esters with alcohols.

Ammonialysis

Similarly silicon tetrachloride undergoes ammonialysis to form chlorosilazanes.

Uses:

1. Silicon tetrachloride is used in the production of semiconducting silicon.
2. It is used as a starting material in the synthesis of silica gel, silicic esters, a binder for ceramic materials.

Silcones:

Silicones or poly siloxanes are organo silicon polymers with general empirical formula (R₂SiO). Since their empirical formula is similar to that of ketone (R₂CO), they were named “silicones”. These silicones may be linear or cross linked. Because of their very high thermal stability they are called high – temperature polymers.

Preparation:

Generally silicones are prepared by the hydrolysis of dialkyldichlorosilanes (R₂SiCl₂) or diaryldichlorosilanes
Ar₂SiCl₂, which are prepared by passing vapours of RCl or ArCl over silicon at 570 K with copper as a catalyst.

The hydrolysis of dialkylchloro silanes R₂SiCl₂ yields to a straight chain polymer which grown from both the sides

The hydrolysis of monoalkylchloro silanes RSiCl₃ yields to a very complex cross linked polymer. Linear silicones can be converted into cyclic or ring silicones when water molecules is removed from the terminal – OH groups.

Types of Silicones:

(i) Liner Silicones:

They are obtained by the hydrolysis and subsequent condensation of dialkyl or diaryl silicon chlorides.

(a) Silicone Rubbers:

These silicones are bridged together by methylene or similar groups.

(b) Silicone Resins:

They are obtained by blending silicones with organic resins such as acrylic esters.

(ii) Cyclic Silicones

These are obtained by the hydrolysis of R₂SiCl₂.

(iii) Cross linked Silicones

They are obtained by hydrolysis of RSiCl₃

Properties

The extent of cross linking and nature of alkyl group determine the nature of polymer. They range from oily liquids to rubber like solids. All silicones are water repellent. This property arises due to the presence of organic side groups that surrounds the silicon which makes the molecule looks like an alkane.

They are also thermal and electrical insulators. Chemically they are inert. Lower silicones are oily liquids whereas higher silicones with long chain structure are waxy solids. The viscosity of silicon oil remains constant and doesn’t change with temperature and they don’t thicken during winter.

Uses:

1. Silicones are used for low temperature lubrication and in vacuum pumps, high temperature oil baths etc
2. They are used for making water proofing clothes
3. They are used as insulting material in electrical motor and other appliances
4. They are mixed with paints and enamels to make them resistant towards high temperature, sunlight, dampness and chemicals.

Silicates

The mineral which contains silicon and oxygen in tetrahedral [SiO₄]^2- units linked together in different patterns are called silicates. Nearly 95% of the earth crust is composed of silicate minerals and silica. The glass and ceramic industries are based on the chemistry silicates.

Types of Silicates:

Silicates are classified into various types based on the way in which the tetrahedral units, [SiO₄]^4- are
linked together.

Ortho Silicates (Neso Silicates):

The simplest silicates which contain discrete [SiO₄]^4- tetrahedral units are called ortho silicates or
neso silicates.

Examples:

Phenacite – Be₂SiO₄ (Be²⁺ ions are tetrahedrally surrounded by O^2- ions), Olivine – (Fe/Mg)₂SiO₄ (Fe²⁺ and Mg²⁺ cations are octahedrally surrounded by O^2- ions).

Pyro Silicate (or) Soro Silicates:

Silicates which contain [Si₂O₇]^6- ions are called pyro silicates (or) Soro silicates. They are formed by joining two [SiO₄]^4- tetrahedral units by sharing one oxygen atom at one corner (one oxygen is removed while joining). Example: Thortveitite – Sc₂Si₂O₇

Cyclic Silicates (or Ring Silicates)

Silicates which contain (SiO₃)_n^2n- which are formed by linking three or more tetrahedral SiO₄^2- units cyclically are called cyclic silicates. Each silicate unit shares two of its oxygen atoms with other units.

Example: Beryl [Be₃Al₂ (SiO₃)₆] (an aluminosilicate with each aluminium is surrounded by 6 oxygen atoms octahedrally)

Inosilicates:

Silicates which contain ‘n’ number of silicate units linked by sharing two or more oxygen atoms are called inosilicates. They are further classified as chain silicates and double chain silicates.

Chain Silicates (or Pyroxenes):

These silicates contain [(SiO₃)_n]^2n- ions formed by linking ‘n’ number of tetrahedral [SiO₄]^4- units linearly. Each silicate unit shares two of its oxygen atoms with other units. Example: Spodumene – LiAl(SiO₃)₂.

Double Chain Silicates (or Amphiboles):

These silicates contains [Si₄O₁₁]_n^6n- ions. In these silicates there are two different types of tetrahedra:

• These sharing 3 vertices
• Those sharing only 2 vertices.

Examples:

1. Asbestos:

These are fibrous and noncombustible silicates. Therefore they are used for thermal insulation material, brake linings, construction material and filters. Asbestos being carcinogenic silicates, their applications are restricted.

Sheet or Phyllo Silicates

Silicates which contain (Si₂O₅)_n^2n- are called sheet or phyllo silicates. In these, Each [SiO₄]^4- tetrahedron unit shares three oxygen atoms with others and thus by forming twodimensional sheets. These sheets silicates form layered structures in which silicate sheets are stacked over each other. The attractive forces between these layers are very weak, hence they can be cleaved easily just like graphite.

Example: Talc, Mica etc.

Three Dimensional Silicates (or Tecto Silicates):

Silicates in which all the oxygen atoms of [SiO₄]^2- tetrahedra are shared with other tetrahedra to form three-dimensional network are called three dimensional or tecto silicates. They have general formula (SiO₂)_n.

Examples: Quartz

These tecto silicates can be converted into three dimensional aluminosilicates by replacing [SiO₄]^4- units
by [AlO₄]^5- units. Eg. Feldspar, Zeolites etc.,

Zeolites:

Zeolites are three-dimensional crystalline solids containing aluminium, silicon, and oxygen in their regular three dimensional framework. They are hydrated sodium alumino silicates with general formula Na₂O.(Al₂O₃).x(SiO₂).yH₂O (x=2 to 10; y=2 to 6).

Zeolites have porous structure in which the monovalent sodium ions and water molecules are loosely held. The Si and Al atoms are tetrahedrally coordinated with each other through shared oxygen atoms. Zeolites are similar to clay minerals but they differ in their crystalline structure.

Zeolites have a three dimensional crystalline structure looks like a honeycomb consisting of a network of interconnected tunnels and cages. Water molecules moves freely in and out of these pores but the zeolite framework remains rigid.

Another special aspect of this structure is that the pore/channel sizes are nearly uniform, allowing the crystal to act as a molecular sieve. We have already discussed in XI standard, the removal of permanent hardness of water using zeolites.

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Group 13 (Boron Group) Elements

The boron occurs mostly as borates and its important ores are borax – Na₂[B₄O₅(OH)₄].8H₂O and kernite – Na₂[B₄O₅(OH)₄].2H₂O. Aluminium is the most abundant metal and occurs as oxides and also found in aluminosilicate rocks. Commercially it is extracted from its chief ore, bauxite (Al₂O₃.2H₂O). The other elements of this group occur only in trace amounts. The other elements Ga, In and Tl occur as their sulphides.

Physical Properties:

Some of the physical properties of the group 13 elements are listed below

Table 2.3 Physical properties of group 13 elements

Element	Most Common Allotropes
Boron	Amorphous boron, α-rhombohedral boron, β-rhombohedral boron, γ-orthorhombic boron, α-tetragonal boron, β-tetragonal boron
Carbon	Diamond, Graphite, Graphene, Fullerenes, Carbon nanotubes
Silicon	Amorphous Silicon, Crystalline Silicon
Germanium	α-germanium, β-germanium
Tin	Grey tin, white tin, rhombic tin, sigma tin
Phosphorus	White phosphorus, Red phosphorus, Scarlet phosphorus, Violet phosphorus, Black phosphorus.
Arsenic	Yellow arsenic, gray arsenic & Black arsenic
Anitimony	Blue-white antimony, Yellow, Black
Oxygen	Dioxygen, ozone
Sulphur	Rhombus sulphur, monoclinic sulphur
Selenium	Red selenium, Gray selenium, Black selenium, Monoclinic selenium
Tellurium	Amorphous & Crystalline

Chemical Properties of Boron:

Boron is the only nonmetal in this group and is less reactive. However, it shows reactivity at higher temperatures. Many of its compounds are electron deficient and has unusual type of covalent bonding which is due to its small size, high ionisation energy and similarity in electronegativity with carbon and hydrogen.

Formation of Metal Borides:

Many metals except alkali metals form borides with a general formula M_xB_y (x ranging upto 11 and y ranging upto 66 or higher) Direct combination of metals with boron:

Direct Combination of Metals With Boron:

Reduction of Boron Trihalides:

Reduction of borontrichloride with a metal assisted by dihydrogen gives metal borides.

Formation of Hydrides:

Boron does not react directly with hydrogen. However, it forms a variety of hydrides called boranes. The simplest borane is diborane – B₂H₆. Other larger boranes can be prepared from diborane. Treatment of gaseous boron triflouride with sodium hydride around 450 K gives diborane. To prevent subsequent pyrolysis, the product diborane is trapped immediately.

Formation of Boron Trihalides:

Boron combines with halogen to form boron trihalides at high temperatures.

Formation of Boron Nitride:

Boron burns with dinitrogen at high temperatures to form boron nitride.

Formation of Oxides:

When boron is heated with oxygen around 900 K, it forms its oxide.

Reaction with Acids and Alkali:

Halo acids have no reaction with boron. However, boron reacts with oxidising acids such as sulphuric acid and nitric acids and forms boric acid.

2B + 3H₂SO₄ → 2H₃BO₃ + 3SO₂
B + 3HNO₃ → H₃BO₃ + 3NO₂

Boron reacts with fused sodium hydroxide and forms sodium borate.

2B + 6NaOH → 2Na₂BO₃ + 3H₂

Uses of Boron:

1. Boron has the capacity to absorb neutrons. Hence, its isotope ¹⁰B₅ is used as moderator in nuclear reactors.
2. Amorphous boron is used as a rocket fuel igniter.
3. Boron is essential for the cell walls of plants.
4. Compounds of boron have many applications. For example eye drops, antiseptics, washing powders etc contains boric acid and borax. In the manufacture of Pyrex glass, boric oxide is used.

Borax [Na₂B₄O₇.10H₂O]:

Preparation:

Borax is a sodium salt of tetraboric acid. It is obtained from colemanite ore by boiling its solution with sodium carbonate.

Borax is normally formulated as Na₂B₄O₇.10H₂O. But it contains, tetranuclear units [B₄O₅. (OH)₄]^2-. This form is known as prismatic form. Borax also exists two other forms namely, jeweller or octahderal borax (Na₂B₄O₇.5H₂O) and borax glass (Na₂B₄O₇).

Properties

Borax is basic in nature and its solution in hot-water is alkaline as it dissociates into boric acid and sodium hydroxide.

Na₂B₄O₇ + 7H₂O → 4H₃BO₃ + 2NaOH

On heating it forms a transparent borax beads.

Borax reacts with acids to form sparingly soluble boric acid.

Na₂B₄O₇ + 2HCl + 5H₂O → 4H₃BO₃ + 2NaCl
Na₂B₄O₇ + H₂SO₄ + 5H₂O → 4H₃BO₃ + Na₂SO₄

When treated with ammonium chloride it forms boron nitride.

Na₂B₄O₇ + 2NH₄Cl → 2NaCl + 2BN + B₂O₃ + 4H₂O

Uses of Borax:

1. Borax is used for the identifiation of coloured metal ions
2. In the manufacture optical and borosilicate glass, enamels and glazes for pottery
3. It is also used as a flx in metallurgy and also acts as a preservative

Boric Acid [H₃BO₃ or B(OH)₃]:

Preparation:

Boric acid can be extracted from borax and colemanite.

Na₂B₄O₇ + H₂SO₄ → 4H₃BO₃ + Na₂SO₄
Ca₂B₆O₁₁ + 11H₂O + 4SO₂ → 2Ca(HSO₃)₂ + 6H₃BO₃

Properties:

Boric acid is a colourless transparent crystal. It is a very weak monobasic acid and, it accepts hydroxyl ion rather than donating proton.

B(OH)₃ + 2H₂O ⇄ H₃O⁺ + [B(OH)₄]^–

It reacts with sodium hydroxide to form sodium metaborate and sodium tetraborate.

H₃BO₃ + NaOH → NaBO₂ + 2H₂O
4H₃BO₃ + 2NaOH → Na₂B₄O₇ + 7H₂O

Action of Heat:

Boric acid when heated at 373 K gives metaboric acid and at 413 K, it gives tetraboric acid. When heated at red hot, it gives boric anhydride which is a glassy mass.

Action of Ammonia

Fusion of urea with B(OH)₃, in an atmosphere of ammonia at 800 – 1200 K gives boron nitride.

Ethyl Borate Test

When boric acid or borate salt is heated with ethyl alcohol in presence of conc. sulphuric acid, an ester, triethylborate is formed. The vapour of this ester burns with a green edged flame and this reaction is used to identify the presence of borate.

Note:

The trialkyl borate on reaction with sodium hydride in tetrahydrofuran to form a coordination compound Na[BH(OR)₃], which acts as a powerful reducing agent.

Formation of Boron Triflouride:

Boric acid reacts with calcium flouride in presence of conc. sulphuric acid and gives boron triflouride.

3CaF₂ + 3H₂SO₄ + 2 B(OH₃) → 3CaSO₄ + 2BF₃ + 6H₂O

Boric acid when heated with soda ash it gives borax

Na₂CO₃ + 4B(OH)₃ → Na₂B₄O₇ + CO₂ + 6H₂O

Structure of Boric Acid:

Boric acid has a two dimensional layered structure. It consists of [BO₃]^3- unit and these are linked to each other by hydrogen bonds as shown in the Figure 2.2.

Uses of Boric Acid:

1. Boric acid is used in the manufacture of pottery glases, enamels and pigments.
2. It is used as an antiseptic and as an eye lotion.
3. It is also used as a food preservative.

Diborane

Preparation:

As discussed earlier diborane can be prepared by the action of metal hydride with boron. This method is used for the industrial production. Diborane can also be obtained in small quantities by the reaction of iodine with sodium borohydride in diglyme.

2NaBH₄ + I₂ → B₂H₆ + 2NaI + H₂

On heating magnesium boride with HCl a mixture of volatile boranes are obtained.

2Mg₃B₂ + 12HCl → 6MgCl₂ + B₄H₁₀ + H₂
B₄H₁₀ + H₂ → 2B₂H₆

Properties:

Boranes are colourless diamagnetic compounds with low thermal stability. Diborane is a gas at room temperature with sweet smell and it is extremely toxic. It is also highly reactive.

At high temperatures it forms higher boranes liberating hydrogen.

Diboranes reacts with water and alkali to give boric acid and metaborates respectively.

B₂H₆ + 6H₂O → 2H₃BO₃ + 6H₂
B₂H₆ + 2NaOH + 2H₂O → 2NaBO₂ + 6H₂

Action of Air:

At room temperature pure diborane does not react with air or oxygen but in impure form it gives B₂O₃ along
with large amount of heat.

B₂H₆ + 3O₂ → B₂O₃ + 3H₂O
∆H = – 2165 KJ mol^2-

Diborane reacts with methyl alcohol to give trimethyl Borate.

B₂H₆ + 6CH₃OH → 2B(OCH₃)₃ + 6H₂

Hydroboration:

Diborane adds on to alkenes and alkynes in ether solvent at room temperature. This reaction is called hydroboration and is highly used in synthetic organic chemistry, especially for anti Markovnikov addition.

B₂H₆ + 6RCH = CHR → 2(RCH₂ – CHR)₃B

Reaction with Ionic Hydrides

When treated with metal hydrides it forms metal borohydrides

Reaction with Ammonia:

When treated with excess ammonia at low temperatures diborane gives diboranediammonate. On heating at higher temperatures it gives borazole.

Structure of Diborane:

In diborane two BH₂ units are linked by two bridged hydrogens. Therefore, it has eight B-H bonds. However, diborane has only 12 valance electrons and are not sufficient to form normal covalent bonds. The four terminal B-H bonds are normal covalent bonds (two centre – two electron bond or 2c-2e bond).

The remaining four electrons have to be used for the bridged bonds. i.e. two three centred B-H-B bonds utilise two electrons each. Hence, these bonds are three centre-two electron bonds (3c-2e). The bridging hydrogen atoms are in a plane as shown in the figure 2.3. In diborane, the boron is sp³ hybridised.

Three of the four sp³ hybridised orbitals contains single electron and the fourth orbital is empty. Two of the half filled hybridised orbitals of each boron overlap with the 1s orbitals of two hydrogens to form four terminal 2c-2e bonds, leaving one empty and one half filled hybridised orbitals on each boron. The Three centre – two electron bonds), B-H-B bond formation involves overlapping the half filled hybridised orbital of one boron, the empty hybridised orbital of the other boron and the half filled 1s orbital of hydrogen.

Uses of Diborane:

1. Diborane is used as a high energy fuel for propellant
2. It is used as a reducing agent in organic chemistry
3. It is used in welding torches

Boron Triflouride:

Preparation:

Boron triflouride is obtained by the treatment of calcium fluoride with boron trioxide in presence of conc. sulphuric acid.

It can also be obtained by treating boron trioxide with carbon and fluorine.

B₂O₃ + 3C + 3F₂ → 2BF₃ + 3CO

In the laboratory pure BF₃ is prepared by the thermal decomposition of benzene diazonium tetrafloro borate.

Properties:

Boron triflouride has a planar geometry. It is an electron deficient compound and accepts electron pairs to form coordinate covalent bonds. They form complex of the type [BX₄]^2-.

BF₃ + NH₃ → F₃B ← NH₃
BF₃ + H₂O → F₃B ← OH₃

On hydrolysis, boric acid is obtained. This then gets converted into Hydro floroboric acid.

Uses of Boron Triflouride:

1. Boron trifloride is used for preparing HBF₄, a catalyst in organic chemistry
2. It is also used as a flourinating reagent.

Aluminium Chloride:

Preparation:

When aluminium metal or aluminium hydroxide is treated with hydrochloric acid, aluminium trichloride is formed. The reaction mixture is evaporated to obtain hydrated aluminium chloride.

2Al + 6HCl → 2AlCl₃ + 3H₂
Al(OH)₃ + 3HCl → AlCl₃ + 3H₂O

McAfee Process:

Aluminium chloride is obtained by heating a mixture of alumina and coke in a current of chlorine.

2Al₂O₃ + 3C + 6Cl₂ → 4AlCl₃ + 3CO₂

On industrial scale it is prepared by chlorinating aluminium around 1000 K

Properties:

Anhydrous aluminium chloride is a colourless, hygroscopic substance. An aqueous solution of aluminium chloride is acidic in nature. It also produces hydrogen chloride fumes in moist air.

AlCl₃ + 3H₂O → Al(OH)₃ + 3HCl

With ammonium hydroxide it forms aluminium hydroxide.

AlCl₃ + 3NH₄OH → Al(OH)₃ + 3NH₄Cl

With excess of sodium hydroxide it produces metal aluminate

AlCl₃ + 4NaOH → NaAlO₂ + 2H₂O + 3NaCl

It behaves like a Lewis acid and forms addition compounds with ammonia, phosphine and carbonylchloride etc… Eg. AlCl₃.6NH₃.

Uses of Aluminium Chloride:

1. Anhydrous aluminium chloride is used as a catalyst in Friedels Craft reactions
2. It is used for the manufacture of petrol by cracking the mineral oils.
3. It is used as a catalyst in the manufacture on dyes, drugs and perfumes.

Alums:

The name alum is given to the double salt of potassium aluminium sulphate [K₂SO₄.Al₂(SO₄)₃.24.H₂O]. Now a days it is used for all the double salts with M^‘SO₄.M^”₂(SO₄)₃. 24H₂O, where M’ is univalent metal ion or [NH₄]⁺ and M^” is trivalent metal ion.

Examples:

Potash alum [K₂SO₄.Al₂(SO₄)₃.24.H₂O]; Sodium alum [Na₂SO₄.Al₂(SO₄)₃. 24.H₂O], Ammonium alum [(NH₄)₂(SO₄)₃.24.H₂O], Chrome alum [K₂SO₄.Cr₂(SO₄)₃.24.H₂O]. Alums in general are more soluble in hot water than in cold water and in solutions they exhibit the properties of constituent ions.

Preparation:

The alunite the alum stone is the naturally occurring form and it is K₂SO₄. Al₂(SO₄)₃. 4Al(OH)₃. When alum stone is treated with excess of sulphuric acid, the aluminium hydroxide is converted to aluminium sulphate. A calculated quantity of potassium sulphate is added and the solution is crystallised to generate potash alum. It is purified by recrystallisation.

K₂SO₄.Al₂(SO₄)₃.4Al(OH)₃ + 6H₂SO₄ → K₂SO₄ + 3Al₂(SO₄)₃ + 12H₂O
K₂SO₄ + Al₂(SO₄)₃ + 24 H₂O → K₂SO₄.Al₂(SO₄)₃. 24 H₂O

Properties

Potash alum is a white crystalline solid it is soluble in water and insoluble in alcohol. The aqueous solution is acidic due to the hydrolysis of aluminium sulphate it melts at 365 K on heating. At 475 K loses water of hydration and swells up. The swollen mass is known as burnt alum. Heating to red hot it decomposes into potassium sulphate, alumina and sulphur trioxide.

Potash alum forms aluminium hydroxide when treated with ammonium hydroxide.

K₂SO₄.Al₂(SO₄)₃.24 H₂O + 6NH₄OH → K₂SO₄ + 3(NH₄)₂SO₄ + 24 H₂O + 2Al(OH)₃

Uses of Alum:

1. It is used for purification of water
2. It is also used for water proofing and textiles
3. It is used in dyeing, paper and leather tanning industries
4. It is employed as a styptic agent to arrest bleeding.