CBSE Notes

By going through these CBSE Class 12 Chemistry Notes Chapter 8 The d-and f-Block Elements, students can recall all the concepts quickly.

The d-and f-Block Elements Notes Class 12 Chemistry Chapter 8

Transition Elements: d-Block. elements are called transition elements. They are placed in between the s-block on their left and the p-block on their right. They are called transition elements because they have properties intermediate between those of s- and p-block elements and represent a change from the most electropositive s-block elements to the least electropositive elements.

This element has partly filled (n -1) d-subshell in their atomic or ionic state. They are called d-block elements since in them 3d, 4d, 5d and 6d subshells are incomplete and the last electron enters (n -1) d subshell, i.e., penultimate (last but one) shell.

General Electronic Configuration:
The outermost shell: Their general electronic configuration is (n – 1) d^1-10 ns^0-2 where n is the outermost shell. Filling of 3d, 4d, 5d and 6d subshells leads to first transition series [4th period], 2nd transition series [5th period], 3rd transition series [6th period] and 4th incomplete transition series [7th period] respectively.

1st Transition Series or 3d series: Starts from Sc (At. No. 21) and ends at Zinc (Zn; atomic no. = 30).

Their configuration is 3d^1-10 ns^1-2.

• Scandium (Sc) = 21 = [Ar]¹⁸ 3d¹ 4s².
• Titanium (Ti) = 22 = [Ar]¹⁸ 3d² 4s²
• Vanadium (V). = 23 = [Ar]¹⁸ 3d³ 4s²
• Chromium (Cr) = 24 = [ Ar]¹⁸ 3d⁵ 4s¹
• Manganese (Mn) = 25 = [Ar]¹⁸ 3d⁶ 4s²
• Iron (Fe) = 26 = [Ar]¹⁸ 3d⁶ 4s²
• Cobalt (Co) = 27 = [Ar]¹⁸ 3d⁷ 4s²
• Nickel (Ni) = 28 = [Ar]¹⁸ 3d⁸ 4s²
• Copper (Cu) = 29 = [Ar]¹⁸ 3d¹⁰ 4s¹
• Zinc (Zn) == 30 = [Ar]¹⁸ 3d¹⁰ 4s²

It may be noted that the electronic configuration of chromium and copper given above are exceptional. They have a single electron in the 4s-orbital. It is due to the extra-stability of 3d⁵ and 3d¹⁰ (half-filled and completely filled) orbitals. They consist up of 10 elements.

2nd Transition Series or 4d Series: Starts from the element Ytterium (Y) [Z = 39] and ends of cadmium (Cd; Z = 46). Their electronic configuration is 4d^0-10 5s^1-2. Palladium (Pd; Z = 46) has exceptional configuration of 4d¹⁰ 5s°. They consists up of ten elements.

Third Transition Series or 5d series: Corresponds to filling of 5d sub-level. They consist up of ten elements La (Z = 57), H/(Z = 72), Ta, W, Re, Os, Ir, Pt, Au and Hg(Z = 80).

The electronic configuration of 3rd transition elements in the inner and valence shell is given below:

Fourth Transition Series or 6d Series: Corresponds to tire filling up of 6d Sub level and starts with Actinium (Ac; Z = 89), Rf (earlier Ku, Z = 104), Db, Sg, Bh, Hs, Mt, Ds, Rg and ends at Uub (Z = 112).

The fundamental difference in the electronic configuration of Representative Elements and Transition Elements: In the representative elements (s- and p-block elements), the valence electrons are present only in the outermost shell while in the transition elements, the valence electrons are present in the outermost shell as well as d- orbitals of the penultimate shell.

Zinc (Zn), Cadmium (Cd) and mercury (Hg) are misfits according to the definition of transition elements as they have 3d¹⁰, 4d¹⁰, 5d¹⁰ (completely filled inner d-orbitals) in the ground state of their atoms or in one of the common oxidation states (dipositive ions: Zn²⁺, Cd²⁺, Hg²⁺) respectively. They do not show properties of transition elements to any appreciable extent, except for their ability to form complexes.

The d-block of the periodic table contains the elements of Groups 3-12 in which the d-orbitals progressively filled in each of the four long periods.

General Properties of Transition Elements:
(a) Metallic character: All the elements of d block are metals. This is due to the fact that they have low ionisation energy values and have only one dr two s-electrons in the outermost shell. The transition elements exhibit good mechanical properties i.e., they are hard, malleable and ductile. Except for Zn, Cd and Hg, they have high m. and b.pts. They have thermal and electrical conductivity and metallic lustre.

(b) Variable oxidation states: The transition elements exhibit a variety of oxidation states in their compounds. This is due to the fact that (n -1) d orbitals are of comparable energy to ns orbitals and therefore some or all of the (n -1) d electrons can be used along with ns electrons in compounds formation. Some common oxidation states exhibited by elements of the first transition series are listed below:

(The value in parentheses are less common oxidation states.)

(c) Coloured compounds: Most of the compounds of transition elements are coloured. The d orbitals in transition metal compounds are not of equal energy. In transition elements, the d-orbital are partly filled and electron may be promoted from one d level to another d level by absorbing visible light. Consequently, the compound has the colour complement to the absorbed light. For example, Cu₂ absorbs red light and transmitted light contains an excess of the other colours of the spectrum and appears to be blue.

The ions Zn²⁺, Cu⁺, Ti⁴⁺, Sc³⁺ are white because they have either completely empty or completely filled d-subshell.

(d) Magnetic properties: Most of the transition metals and their compounds are paramagnetic i.e., they are attracted by the magnetic field. This, the property is due to the presence of unpaired electrons in the d orbitals of their atoms. The elements like iron, nickel and cobalt have appreciable paramagnetism and are called ferromagnetic substances. The magnetic moment of compounds of d-block elements are determined by spin only values which may be related to the number of impaired electrons (n) as

Spin magnetic moment (μ_s) = \(\sqrt{n(n+2)}\)

The agreement between the observed and calculated values is quite good for most of the elements. However, for the latter half of the series, the contribution from orbital magnetic moments is observed.

(e) Complex formation: The transition elements have a great tendency to form complexes.

This is because of:

• their small cation size
• high effective nuclear charge and
• presence of empty d-orbitals of appropriate energy for bonding to the ligands.

It has been noticed that

1. In each transition series, the stability of complexes increases with an increasing atomic number of the element and in a particular oxidation state with decreasing size of its atoms.
2. In the case of metal atoms that show more than one oxidation state, the complexes of greater stability are formed with the metal ions of the highest charge.

(f) Catalytic properties: Many transition metals and their compounds show catalytic properties e.g., Fe, Pt, V₂O₅, Ni etc. This property may be due to either the use of the d-orbitals or from the formation of transient intermediate compounds to absorb and activate the reacting substances.

(g) Alloy Formation: Transition metals mix freely with each other in the molten state and on cooling a solution of different transition metals to form alloys. For example, chromium dissolves in nickel to form Cr-Ni alloy; manganese dissolves in iron to form manganese steel. They take part in alloy formation because their atoms can exchange lattice sites of each other. Such alloys are formed by atoms with metallic radii that are within about 15 per cent of each other.

(h) Formation of Interstitial Compounds: Interstitial compounds are those which are formed when small atoms like H, C or N are trapped inside the crystal lattices of metals. They are usually non-stoichiometric and are neither typically ionic nor covalent. Many of the transition metals form interstitial compounds particularly with small non-metal atoms such as hydrogen, boron, carbon and nitrogen.

These small atoms enter into the voids sites between the packed atoms of the crystalline metal, e.g., TiC, Mn4N, Fe3M and TiH2 etc. The formulae quoted do not, of course, correspond to any normal oxidation state of the metal and often non-stoichiometric material is obtained with such composition as VH056 and TiH17. Because of the nature of their composition, these compounds are referred to as interstitial compounds.

The principal physical and chemical characteristics of these compounds are as follows:

1. They have high melting points, higher than those of pure metals.
2. They are very hard, some borides approach diamond in hardness.
3. They retain metallic conductivity.
4. They are chemically inert.

Some Important Compounds of Transition Metals:
Oxides: Transition metal oxides are formed by the action of oxygen with transition metals at high temperature. The general formulas of the oxides of transition metals are MO, M₂O₃, M₃O₄, MO₂, M₂O₅ and MO₃.

It has been observed that:

• the oxides in which the metals are in a lower oxidation state are basic in nature.
• the oxides in which the metal is in a higher oxidation state are acidic in nature.
• the oxides in which the metal exhibits intermediate oxidation states are amphoteric.

For example, among the oxides of manganese, MnO is basic, Mn₂O₃, Mn₃O₄ and MnO₂ are amphoteric while Mn₂O₇ is acidic.

Oxides Of 3d Metals:

Chemical Reactivity and E° Values: Transition metals vary widely in their chemical reactivity. Many of them are sufficiently electropositive to dissolve in mineral acids, although a few are ‘noble’: that is, they are unaffected by simple acids.

The metals of the first series with the exception of copper are relatively more reactive and are oxidised by 1M H⁺, though the actual rate at which these metals react with oxidising agents like hydrogen ion (H⁺) is sometimes slow. For example, titanium and vanadium, in practice, are passive to dilute non-oxidising acids at room temperature. The Ee values for M²⁺/M indicate a decreasing tendency to form divalent cations across the series.

This general trend towards less negative Ee values is related to the increase in the sum of the first and second ionisation enthalpies. It is interesting to note that the Ee values for Mn, Ni, and Zn are more negative than expected from the general trend. Whereas the stabilities of half-filled d subshell (d⁵) in Mn²⁺ and completely filled d subshell (d¹⁰) in zinc are related to their E^θ values; for nickel, E^θ value is related to the highest negative enthalpy of hydration.

An examination of the E^θ values for the redox couple M³⁺/M²⁺ shows that Mn³⁺ and CO³⁺ ions are the strongest oxidising agents in aqueous solutions. The ions Ti²⁺, V²⁺ and Cr²⁺ are strong reducing agents and will liberate hydrogen from a dilute acid. e.g.
2 Cr²⁺(aq) + 2 H⁺(aq) → 2 Cr³⁺(aq) + H₂(g)

Trends in Stability of Higher oxidation states as exhibited in the oxides and halides of 3d-Series:
The table below shows the stable halides of the 3d series of transition metals. The highest oxidation numbers are achieved in Ti tetrahalides, VF₅ and CrF₆. The +7 state for Mn is not represented but MnOF₃ is known and beyond Mn, no metal has a trihalide except Fe and COF₃.

The ability of F to stabilize the higher oxidation state is due to either higher lattice energy of the higher compound as in the case of COF₃ or higher bond energy terms for the higher covalent compounds e.g. VF₅ and CrF₆.

Table: Formulae of halides of groups 4 to 12:

Halides: Transition metals react with halogens at high temperature to form transition metal halides. The order of reactivity of halogens is F > Cl > Br > I. The halides of higher oxidation states are the fluorides because metals can be oxidised to higher oxidation states with only fluorine which is most reactive. Bonding in fluorides is mainly ionic. In chlorides, bromides and iodides, the ionic character decrease with the atomic mass of halogen.

The ability of O to stabilize the higher oxidation state is also demonstrated in the oxohalides. Although V^v is represented only by VF₅, the oxohalides VOX₃ are known where X = F, Cl or Br. Another feature of fluorides is their instability in the low oxidation states e.g., VX₂ (X = Cl, Br or I) and the same applies to CuX. On the other hand, all Cu¹¹ halides are known except the iodide, in this case, Cu²⁺ oxidises I^– to I₂:
Cu²⁺ + 2I^–→ CuI(S) + \(\frac{1}{2}\) I₂

However Cul is unstable in solution and undergoes disproportionation.
2Cu⁺ → Cu²⁺ + Cu

The stability of Cu²⁺ (aq) rather than Cu⁺ (aq) is due to the much more negative Ah d H ° of Cu²⁺ (aq) than Cu⁺ (aq), which more than compensates for the 2nd IE of Cu.

The highest oxidation number in the oxides (Table) coincides with the group number and is attained in Sc₂O₃, to Mn₂O₇. Beyond Group 7, no higher oxides of Fe above Fe₂O₃, are known, although ferrates (VI):(FeO₄)^2- are formed in alkaline media they readily decompose to Fe₂O₃ and O₂. Besides the oxides, oxidations stabilise V^v as VO²⁺, Vlv VO²⁺ and TiIV as TiO²⁺.

The ability of O to stabilise these high oxidation states exceeds F in this regard. Thus the highest Mn fluoride is MnF₄ whereas the highest oxide is Mn₂O₇. The ability of oxygen to form multiple bonds to metals explains its superiority. In the covalent oil Mn₂O₇, each Mn is tetrahedrally surrounded by O’s including an Mn-O-Mn bridge. The tetrahedral [MO₄]^n- ions are known for V^v, Cr^vI, Mn^v, Mn^vI and Mn^vII.

Comparison of the First Row Transition Metals Through The d-EIectron Configuration:
In the d° configuration of the simple ions, only Sc³⁺ is known to have this configuration. This configuration then occurs for those metals in which the formal oxidation states equal the total no. of 3d and 4s electron. This is true for Ti (IV), V (V), Cr (VI) and Mn (VII).

→ The d¹ Configuration: Except Vanadium (IV), all others with this configuration are either reducing or undergo disproportionation. For .example, disproportionation occurs for Cr (V) and Mn (VI)
3 CrO₄^3- + 8H⁺ → 2 CrO₄^– + Cr³⁺ + 4 H₂O
3 MnO₄^2- + 4H⁺ → 2 MnO₄^– + MnO₃  + 2H₂0

→ The d² Configuration: This configuration ranges from Ti¹¹ which is very strongly reducing, to FeVI which is very strongly oxidizing. Vanadium (III) is also reducing.

→ The d³ configuration is shown by Chromium (III). It is quite stable and takes part in complex formation.

→ The d⁴ Configuration: There are really no stable species with the configuration Cr (II) is strongly reducing. ,

→ The d⁵ Configuration: The two important species with this configuration are Mn²⁺ and Fe³⁺, the latter may, however, be reduced to Fe²⁺.

→ The d⁶ Configuration: Iron (II) and Cobalt (III) are important species with this configuration. Iron (II) is quite stable although a mild reducing agent and cobalt (III) are stable in the presence of strong complexing reagents.

→ The d⁷ Configuration: The species with this configuration is cobalt (II) which is stable in aqueous solutions but gets oxidized to form CO (III) complexes in the presence of strong ligands.

→ The d⁸ Configuration: Nickel (II) is the most important species with this species.

→ The d⁹ Configuration: This configuration is found in Cu²⁺ compounds. It is by far the most important in the chemistry of copper.

→ The d¹⁰ Configuration: The two species Cu⁺ and Zn²⁺ are important with this configuration. Whereas Copper (I) is easily oxidized to copper (II), zinc (II) is the only state known for zinc.

General Group Trends in the Chemistry of the d-Block Metals:
Group 4: The titanium group of transition metals consists up of the elements titanium, zirconium and hafnium. The most important member of this group is titanium. It is extremely strong, has a high melting point and is resistant to corrosion. It is in great demand as a structural material. Zirconium and hafnium are silvery-white metals. The most stable oxidation state for the elements of this group is +4.

The titanium also possesses a + 3 oxidation state. The typical compounds of these elements are chlorides, TiCl₄, ZrCl₄, HfCl₄ and oxides TiO₂, ZrO₂ and HfO₂, ZrO₂ is a refractory material. Zirconium and hafnium occur together and exhibit similar properties.

Their atomic radii (Zr = 160 pm, Hf = 159 pm) are almost equal. This is due to the reason that usually increases in size down the group is cancelled by the lanthanide contraction. Zirconium and hafnium are both important for the generation of nuclear energy.

The vanadium group (Group 5): consists of vanadium, niobium and tantalum. The most stable oxidation state for this group is + 5. Vanadium is used as an additive to steel. The most important compound of vanadium is its pentoxide, V₂O₅, which is used as a catalyst in many reactions. Niobium alloys are used in jet engines, Tantalum is very resistant to corrosion and is used for making apparatus in chemical plants. It is also used in surgery, as for bone pins.

The chromium group (Group 6): contains the elements, chromium, molybdenum and tungsten. The most important oxidation states for chromium are + 3 and + 6, and for molybdenum and tungsten are + 5 and + 6. The zero oxidation states for these elements arise in metal carbonyls such as Cr (CO)₆.

These elements have very high melting and boiling points. Tungsten is the metal with the highest melting point. These metals are also very hard. Chromium is unreactive or passive at low temperatures because of the formation of a surface coating of oxide. It is due to this passive behaviour that chromium is used for electroplating iron to prevent rusting.

The metals of this group are very useful. Chromium is used in many ferrous alloys such as stainless steel, chrome steel, etc. It is also used for electroplating iron or some other metals to prevent corrosion. K₂Cr₂O₇ is a very important and useful compound of chromium.

Molybdenum and tungsten form very hard alloys with steel and are used in making cutting tools. Tungsten is also used as filament in electric bulbs. Molybdenum disulphide, MoS2 acts as a lubricant because it has a layer lattice.

Group 7: the manganese group: consist of the elements manganese, technetium and rhenium. These elements exhibit all the oxidation states from 0 to + 7, the most important being + 2, + 4 and + 7 for manganese and + 4 and + 7 for technetium and rhenium.

The elements of this group have quite a high melting and boiling points.

Manganese is obtained from its oxide ores by reduction with carbon or aluminium. Manganese metal does not have any use as such but is used in the manufacture of alloys such as ferromanganese (Fe + Mn) and manganese bronze (Mn + Cu + Zn). KMnO₄ is an important compound of Mn which is used as an oxidizing agent and finds so many other applications. Manganese dioxide is used as a catalyst. Rhenium is used in electronic filaments, high-temperature thermocouples and in flashbulbs.

The metals of group 8,9 and 10 are known as iron group metals: These elements are:

The first triad comprising of iron, cobalt and nickel is known as ferrous metals. These metals are ferromagnetic. Iron and cobalt exhibit oxidation states of + 3 and + 2 in their compounds while nickel compounds are generally in the + 2 oxidation state.

The elements of the second and third triad, ruthenium, rhodium, palladium, osmium, iridium and platinum are collectively known as platinum metals. These elements are relatively less abundant and exhibit a wider range of oxidation states. They are inert and serve as good catalysts.

The copper group (group 11): includes the elements copper, silver and gold. These metals are known as coinage metals. They form alloys with many metals. The most stable oxidation state for copper is + 2. For silver and gold, the oxidation state of + 1 is relatively more stable. The metals of this group have the highest electrical and thermal conductivities.

The zinc group (group 12): consists of the elements zinc, cadmium and mercury. The elements of this group show none of the characteristic properties of transition metals. The metals of this group are moderately electropositive and exhibit an oxidation state of + 2. Since their ionization energy is very high none of these metals possesses an oxidation state higher than + 2. Mercury, due to metal-metal bond also shows a formal oxidation state of +1. The elements of this group are diamagnetic. Mercury is the only metal that exists as a liquid.

Potassium Dichromate (K2Cr2O7):
Preparation: Chromite ore is fused with sodium or potassium carbonate in free access of air when the following reaction occurs.

The yellow solution of sodium chromate (Na₂CrO₄) is filtered and acidified with a calculated amount of sulphuric acid from which orange sodium dichromate Na₂Cr₂O₇. 2H₂O can be crystallised.
2 Na₂CrO₄ + 2H⁺ → Na₂Cr₂O₇ + 2Na⁺ + H₂O

Potassium dichromate is prepared by treating the solution of sodium dichromate with a calculated amount of potassium chloride.
Na₂Cr₂O₇ + 2 KCl → K₂Cr₂O₇ + 2 NaCl.

Orange crystals of potassium dichromate crystallise out. The dichromate and chromate are interconvertible in an aqueous solution depending upon the pH of the solution.
2 CrO₄^2- + 2 H⁺ → Cr₂O₇^2- + H₂O.
Cr2O₇^2- + 2 OH^– → 2 CrO₄^2- + H₂O.

The structures of chromate ionCrO₄^2-, and dichromate ion, Cr₂O₇^2- and shown below. Whereas CrO₄^2- is tetrahedral sharing one comer with Cr-O-Cr bond angle of 126°.

In acidic solution, both sodium and potassium dichromates are strong oxidizing agents, the former has a greater solubility in water.
Cr₂O₇^2- + 14 H⁺ + 6e^– → 2 Cr³⁺ + 7 H₂O [E° = 1.33 V]

→ Acidified solution oxidizes
1. iodides to iodine
6I^– → 3I₂ + 6e^–

2. sulphides to sulphur
3 H₂S → 6 H+ + 3 S + 6e^–

3. tin (II) to tin (IV)
3 Sn²⁺ → 3 Sn⁴⁺ + 6e^–

4. iron (II) salts to iron (III)
6 Fe²⁺ → 6 Fe³⁺ + 6e^–
The full ionic equation may be obtained by adding the two half-reactions. e.g.,
Cr₂07^2- + 14 H⁺ + 6 Fe²⁺ → 2 Cr³⁺ + 6 Fe³⁺ + 7 H₂O.

Potassium Permanganate, KMn04:
Preparation:
1. KMnO₄ is prepared by fusion of MnO, and KOH with an oxidizing agent like KNO₃.

Potassium manganate (K₂MnO₄) gives KMnO₄ in a neutral/acidic solution.
3 MnO₄^2- + 4H) → 2 MnO₄^– + MnO₂ + 2H₂O

2. Comfnercially, KMnO₄ is prepared from MnO₂ by electrolytic oxidation.

KMnO₄ forms dark purple crystals. It is not very soluble in water. When heated it decomposes at 240°C.

KMnO₄ shows weak temperature-dependent paramagnetism. It arises due to the coupling of the diamagnetic ground state of MnO₄ ion with paramagnetic excited states under the influence of the magnetic field.

The manganate and permanganate ions are tetrahedral. The green MnO₄^2- is paramagnetic with one unpaired electron, but MnO₄^– is diamagnetic.

Permanganic acid (HMnO₄) can be obtained by low-temperature evaporation of its aqueous solution. It is a strong oxidizing agent and in a pure state, it is explosive above 0°C.

→ Acidified KMnO₄ is a strong oxidizing agent:
MnO₄^– + e^– → MnO₄^2-
[Reduction half reaction]; E° = + 0.56 V

MnO₄^– + 4H⁺ + 3e^– → MnO₂ + 2H₂O
[Reduction half reaction]; E° = + 1.69 V.

MnO₄^– + 8H⁺ + 5e^– → Mn²⁺ + 4 H₂O
[Reduction half reaction]; E° = + 1.52 V.

A few important oxidizing reactions of KMnO₄ are given below.
1. In acid solutions:
(a) Iodine is liberated from potassium iodide:
10I^– + 2MnO₄^– + 16H⁺ → 2Mn²⁺ + 8H₂0 + 5I₂

(b) Fe²⁺ ion is converted to Fe³⁺:
5Fe²⁺ + MnO₄^– + 8H⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
The green iron (II) solution changes to yellow iron (III) solution.

(c) Oxalate ion or oxalic acid is oxidized at 60°C:
5 C2O₄^2- + 2MnO₄^–+ 16H⁺ → 2Mn²⁺ + 8H₂O + 10CO₂

(d) Hydrogen sulphide is oxidized, sulphur being precipitated:
H₂S + 2H⁺ → + S^2-
5S^2- + 2MnO₄^– + 16H⁺ → 2Mn²⁺ + 8H₂O + 5S

(e) Sulphurous acid or sulphite is oxidized to a sulphate or sulphuric acid:
5 SO₃^2- + 2MnO₄^– + 6H⁺ → 2Mn²⁺ + 3H₂O + 5SO₄^2-

(f) Nitrite is oxidized to nitrate:
5NO₂^– + 2MnO₄^– + 6H⁺ → 2Mn²⁺ + 5NO₃^– + 3H₂O

2. In neutral or faintly alkaline solution:
(a) A notable reaction is the oxidation of iodide to iodate:
2MnO₄^– + H₂O + I^– → 2MnO₂ + 2OH^– + IO₃^–

(b) Thiosulphate: s oxidised almost quantitatively to sulphate, a trace of thionate being formed
8MnO₄^– + 3S₂O₃^2- + H₂O → 8MnO₂ + 6SO₄^2- + 2OH^–

(c) Manganous salt is oxidised to MnOz; the presence of zinc sulphate or zinc oxide catalyses the oxidation:
2MnO₄^– + 3Mn²⁺ + 2H₂O → 5MnO₂ + 4H⁺

Note: Permanganate titrations in presence of hydrochloric acid are unsatisfactory since hydrochloric acid is oxidised to chlorine.

Uses: Besides its use in analytical chemistry, potassium permanganate is used as a favourite oxidant in preparative organic chemistry. Its uses for the bleaching of wool, cotton, silk and other textile fibres and for the decolourisation of oils are also dependent on its strong oxidising power.

→ The Inner Transition Elements (f-Block): It consists up of two series of elements.
1. Lanthanoids (58-71) which follow lanthanum (57) in the periodic table. are fourteen in number. Because La (57) closely resembles the lanthanoids, it is usually included in the study of lanthanoids and the general symbol Ln is often used.

2. Actinoids [the fourteen elements (90-103)] which follow actinium (89). A discussion of actinoids also includes actinium (Ac), besides the 14 elements.

In the case of lanthanoids, 4f orbitals/are filled up and in the case of actinoids, 5f orbitals are filled up. That is why they are called Inner Transition Elements as the last electron (also called differentiating electron) enters the antepenultimate energy level, ie.., (n – 2) f-orbitals: inner to the penultimate energy level and they form a transition series within the transition series (d-block elements)

Lanthanides or Landhanoids or Lanthanoness: In these elements, the last electron enters the 4f-orbitals and is also referred to as the first inner transition series. Earlier they were called rare earth.

Actinides or Actinoids or Actions: In these elements, the last electron enters one of the 5f-orbitals and is also required as the second inner transition series.

The study of lanthanoids is easier because they show only one stable oxidation state (+ 3). On the other hand, the chemistry of actinoids is much more complicated partly because they show a wide range of oxidation states and partly because they are radioactive.

→ Electronic Configuration: General E.C. of lanthanoids and actinoids is (n – 2) f^1-14 (n – 1) d^0-1 ns². Thus they have three incomplete shells, viz., (n – 2), (n – 1) and nth. Their electronic config. is given below.

Trends in Ionic Radii of Trivalent Lanthanoids

Table: Electronic configurations and radii of lanthanum and lanthanoids:

Only electrons outside [Xe] core are indicated.

The overall decrease in atomic and ionic radii from lanthanum to lutetium (the lanthanoid contraction) is a unique feature in the tire chemistry of lanthanoids. It is due to the imperfect shielding of one 4f electron from another 4f electron due to the highly diffused shape of f-orbitals. With the increase in nuclear charge, there is a fairly regular decrease in their sizes.

The cumulative effect of the contraction of the lanthanoid series called lanthanoid contraction causes the radii of the third transition series to be very similar to those of the correspond ing members of the 2nd series. As a.result Zr (160 pm) and Hf (159 pm) have identical radii and so exist together.

Thus Zr and Hf face difficulty in their separation.
Colour and Paramagnetism: Many trivalent lanthanoids are coloured both in the solid-state and in aqueous solutions. It is due to the presence of f-electrons. Neither Lu³⁺ ion (f°) nor Lu³⁺ (f¹⁴) shows any colour. All others (having f¹ to f¹³ arrangement) show colour. The lanthanoid ions other than f° (La³⁺, Ce⁴⁺) type and the f¹⁴ type (Yb²⁺, Lu³⁺) are all paramagnetic due to the presence of unpaired electrons. Paramagnetism is maximum in neodymium.

→ Ionisation Enthalpies: The IE₁ and IE₂ are around 600 kJ mol^-1 and 1200 kJ moH comparable with those of calcium. The values of IE₃ indicate that the exchange energy considerations impart a degree of stability to empty, half-filled and completely filled f-level (f°, f⁷, f¹⁴). It is evident from the abnormally low values of IE₃ of lanthanum, gadolinium and lutetium.

→ Oxidation States: All the lanthanoids show a + 3 oxidation state which is most significant: Some of them show + 2 and + 4 oxidation states also as ions in solution or in solid compounds. This irregularity arises mainly from the extra stability of empty, half-filled or completely filled f-subshell. Ce^IV [₁₆Rn-noble gas] is formed easily, but it is a strong oxidant reverting to the common oxidation state of + 3.
Ce⁴⁺ + e^– → Ce³⁺] Reduction half reaction
E° for the above is = + 1.74 V

If can oxidise water, but the reaction rate is slow. Pr, Nd, Tb and Dy also exhibit + 4 states in oxides MOr Eu²⁺ is formed by losing the two s-electrons (F). However, Eu²⁺ is a strong reducing agent changing to the common oxidation state + 3.
EU²⁺ → EU³⁺ + e^–

Similarly Yb²⁺ [f¹⁴] is a reductant. Tb (IV) [f⁷] is an oxidant.
Tb4+ + e- → Tb3+

Samarium also shows oxidation states of + 2 and + 3 like Eu.

Properties and Uses:
1. All the lanthanides are silvery-white soft metals and tarnish rapidly in the air. Hardness increases with increasing atomic number, Samarium is steel hard.

2. M.Pts are in the range of 1000 to 1200 K, but Sm melts at 1623 K

3. They have a metallic structure.

4. They are good conductor of heat and electricity.

5. In their chemical behaviour, lanthanoids are generally similar to calcium (Ca = 20), but with increasing atomic number, they behave more like aluminium.

6. Values for E° for the reduction half-reaction.
Ln³⁺ (aq) + 3e^– → Ln (s) are in the range of – 2.2 to – 2.4 V except for Ln = Eu, for which it is – 2.0 V,

7. The metals combine with hydrogen when heated gently.
2 Ln + 3 H₂ → 2 Ln H₃

8. When heated in carbon, carbides of the type Ln₃C, Ln₂C₃ and LnC₂ are formed.

9. They liberate H2 gas from dilute acids.
2 Ln + 6HCl (dil.) → 2 LnCl₃ + 3H₂I

10. They bum in halogens to form halides.
2 Ln + 3X₂ → 2 LnX₃

11. They form oxides and hydroxides of the type M₂O₃ and M (OH)₃. They are basically like alkaline earth metal oxides and hydroxides. A summary of the chemical reactions of lanthanoids are given below:

The best use of the lanthanoids is for the production of alloy steels for plates and pipes. A well-known alloy is Misch Metall which consists of a lanthanoid metal (~ 95%) and iron (~ 5%) and traces of S, C, Ca and Al. It is used to produce bullets, shell and lighter flint. Mixed oxides are used as catalysts in cracking. Some Ln oxides are used as phosphors in TV screens.

→ The Actinoids: The elements which follow actinium (89) in the periodic table are called actinoids [Starting from Th – 90 to Lr – 103.]

→ Actinoids are radioactive elements. Earlier members have relatively long half-lives and later ones from day to 3 minutes for Lr. This makes their study difficult.

→ Electronic Configuration: All the actinoids have a 7s² electron configuration and variable occupancy of the 5f and 6d subshells. The irregularities in the electronic configuration of actinoids like those in lanthanoids are related to the stabilities of 5f°, 5f⁷ and 5f¹⁴. 5f orbitals participate in bonding to a greater extent.

→ Oxidation states: There is a greater range of oxidation states shown by actinoids due to the comparable energies of 5f, 6d and 7s levels.

The electronic configurations and oxidation states of actinoids are in the following tables.

Table: Oxidation states of actinium and actinoids:

Unlike 4f orbitals of lanthanoids which are buried, 5f orbitals of actinoids participate in bonding to a far greater extent. In addition to showing an oxidation state of + 3 like lanthanoids, actinoids also show oxidation states of + 4, + 5, + 6 and + 7. Actinoids which show several oxidation states [Np, Pu, Am] make it difficult to review their chemistry.

Physical and Chemical reactivity of actinoids

• Actinoids are all silvery-white metals.
• They are highly reactive metals, especially when finely divided.
• Boiling water reacts with them to give a mixture of oxides, hydrides.

Some applications of d-Block elements

1. Iron and Steel are the most important construction materials.
2. Compounds like TiO are prepared for use in the pigment industry.
3. Mn02 is used in dry battery cells. The battery industry also requires Zn and Ni/Cd.
4. Cu, Ag and Au are coinage metals.
5. Many of the metals and their compounds find use as catalysts like finely divided Ni, V₂O₅, TiCl₄, Fe etc.
6. PdCl₂ is used in the Wacker process for the oxidation of ethyne to ethanol
7. Ni-complexes are used in the polymerisation of alkynes and other organic compounds like benzene.
8. AgBr is used in the photographic industry.

By going through these CBSE Class 12 Chemistry Notes Chapter 7 The p-Block Elements, students can recall all the concepts quickly.

The p-Block Elements Notes Class 12 Chemistry Chapter 7

The p-block elements are placed in groups 13 to 18 of the periodic table. Their valence shell electronic configuration is ns²np⁶ to ns²np⁶ (except He which has 1s² electronic configuration.)

Group 15 Elements: Group 15 elements include Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi).

Table 7.1: Atomic and Physical Properties of Group 15 Elements

^αE₁₁₁ single bond (E = element); ^bE^3-; ^cE³⁺; ^dWhite phosphorus; ^eGrey α-format 38.6 atm of Sublimation temperature; g At 63 K h Grey α-form; Molecular N₂

As we go down the group, there is a shift from non-metallic to the metallic character through metalloids character. N and P are non-metals, AS and Sb are metalloids and Bi is a typical metal.

→ Occurrence: N₂ comprises 78% by volume of the atmosphere. In the earth’s crust, it occurs as sodium nitrate (NaNO₃) called Chile saltpetre and potassium nitrate (KNO₃) called Indian saltpetre. It is found in the form of proteins in plants and animals. Phosphorus (P) occurs in minerals of the apatite family Ca₉ (PO₄)₆. CaX₂ [X = F, Cl, or OH] e.g. fluorapatite Ca₉ (PO₄)₆. CaF₂ which are the main components of the phosphate rocks. P is an important constituent of animal and plant matter and is present in bones and living cells.

→ Electronic Configuration: The valence shell electronic configuration of these elements is ns2np3. The s orbital is completely filled and p orbitals are exactly half-filled (p_x¹ p_y¹ P_z¹). Thus their electronic configuration is extra stable.

→ Atomic and Ionic Radii: Covalent and ionic radii increase in size down the group due to the addition of the orbit each time. There is a considerable increase from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and /or f orbitals in heavier members.

→ Ionization Enthalpy: I.E. values of Group 15 elements is much greater than that of Group 14 elements due to the extra-stability of half-filled p-orbitals and smaller size. It decreases from top to bottom within the group due to a gradual increase in atomic size. The order of successive ionisation enthalpies is Δ_iH₁ < Δ_iH₂ < Δ_iH₃.

→ Electronegativity: Electronegativity decrease down the group. The difference is not much pronounced amongst the heavier elements.

→ Physical Properties: All the elements are polyatomic. Dinitrogen (N₂) is a diatomic gas while others like P₄ etc. are solids. The b. pts; in general, increase from top to bottom in the group, but them. pts increase up to As and then decrease up to Bi.

Except for Nitrogen, all the elements show allotropy.

Chemical Properties:
Oxidation states and trends in chemical reactivity: Tire common oxidation states of these elements are – 3, + 3 and + 5. The tendency to exhibit – 3 oxidation state decreases down the group due to an increase in size and metallic character.

N does not show + 5 oxidation state due to the absence of d-orbitals and Bi does not show + 5 oxidation state due to the Inert pair Effect. The only well characterised Bi (V) compound is BiF₅. N exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen. P also exhibits + 1 and + 4 oxidation states in some oxoacids.

In the case of N, all oxidation states from + 1 to + 4 tend to disproportionate into acid solution1. For example
3 HNO₂ → HNO₃ + 2 NO + H₂O

Similarly, in the case of P nearly all intermediate oxidation states disproportionate into + 5 and – 3 both in alkali and acid,

Anomalous properties of Nitrogen: N differs from its congeners due to

1. its smaller size.
2. high electronegativity and high ionisation enthalpy.
3. non-availability of d-orbitals.

Nitrogen has a unique ability to form an-pit multiple bonds with itself and with other elements having small size and having high electronegativity (e.g., C, O)

N₂ has a triple bond: N ≡ N [one a and two K bonds]
∴ Its bond enthalpy (941.4 kJ mol^-1) is very high.
P and As can form dπ-pπ bonds with transition metals.

1. Reactivity towards hydrogen: All the elements of Group 15 form hydrides of the type EH₃ where E = N, P, As, Sb, Bi. The stability of hydrides decreases from NH₃ to BiH₃ and hence reducing character increases from NH₃ to BiH₃. Basicity of hydrides also decreases in the order NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃.

Table 7.2: Properties of Hydrides of Group 15 Elements

2. Reactivity towards oxygen: All these elements form oxides: E₂O₃ and E₂O₅. The oxide in the higher oxidation state is more acidic.
Their acidic character decreases down in a group. The oxides of the type EO₃ of N and Pare purely acidic, that of As, Sb amphoteric and those of Bi predominantly basic.

3. Reactivity towards halogens:
2E + 3X₂ → EX₃
2E + 5X₂ → 2EX₅
N does not form pentahalide due to absence of d-orbitais. All trihalides are stable except that of N.

Pentahalides are more covalent than trihalides.

4. Reactivity towards metals: These elements combine with metals showing – 3 oxidation state in compounds like Ca₃N₂, Ca₃P₂, Mg₃Bi₂.

Dinitrogen, N₂:
Preparation:
1. Commercially dinitrogen is produced by the liquefaction and fractional distillation of air when N₂ distils the first.

2. Lab. Method:
NH₄Cl (aq) + NaNO₂ (aq) → N₂ (g) + 2H₂O (l) + NaCl (aq)

3. Pure N₂ can be obtained from sodium or barium azide.

→ Properties: Dinitrogen is a colourless, odourless, tasteless and non-toxic gas. It has two stable isotopes: ¹⁴N, ¹⁵N. It has very low solubility in water.

Because of the high bond enthalpy of N ≡ N, dinitrogen is rather inert at room temperature.

At high temperature, it directly combines with some metals to form largely ionic nitrides and with non-metals, covalent nitrides.

Haber’s process:

Uses:

1. Manufacture of NH₃ and Calcium Cyanamide.
2. To create an inert atmosphere.
3. Liquid N₂ is used as a refrigerant,

Ammonia:
Preparation:
1. From Urea:
NH₂CONH₂ + 2H₂O → (NH₄)₂ CO₃ → 2NH₃ + H₂O + CO₂

2. From ammonium salts and caustic soda/lime.
2NH₄Cl + Ca(OH)₂ → 2NH₃ + 2H₂O + CaCl₂
(NH₄)₂ SO₄ + 2 NaOH → 2NH₃ + 2H₂O + Na₂SO₄

3. On a large scale, it is prepared by Haber’s process.

The pressure is 200 × 10⁵ Pa and iron oxide with small amounts of K₂O and Al₂O₃ is used as a catalyst,

Properties of Ammonia:

1. It is a colourless gas with a pungent smell.
2. It shows hydrogen bonding in solid and liquid states.

Structure of NH3:

3. It is highly soluble in water. Its aqueous solution is weakly basic.
NH₃ (g) + H₂O (l) ⇌ NH₄⁺ (aq) + OH^– (aq)

4. It forms ammonium salts with acids.
NH₃ + HCl → NH₄Cl

5. As a weak base it precipitates the hydroxides of many metals
from their salt solutions, e.g.,

6. Due to its tendency to donate a pair of electrons on N, it acts as a Lewis-Base, forms complexes with Cu²⁺, Ag⁺.


Uses:

1. Production of nitrogenous fertilizers.
2. Production of nitric acid.
3. Liquid ammonia is a non-aqueous solvent.
4. Liquid NH₃ is a referigerant.

Oxides of Nitrogen: Nitrogen forms a number of oxides in different oxidation state. The details of names, formulae, ox. state, preparation, colour etc. are given:

Lewis dot main resonance structures and bond parameters of oxides are given on the next page.

Table 7.4 Structures of Oxides of Nitrogen

Nitric Acid, HNO₃
Preparation: It is prepared in the laboratory by heating NaNO₃ or KNO₃ and cone. H₂SO₄.

Method of manufacture of Nitric acid by Ostwald’s process.

Step:
1. Catalytical oxidation of NH₃ by O, or air.

2. NO thus formed combines with 02 giving NO₂
2NO (g) + O₂ (g) ⇌ 2NO₂(g).

3. N02 SO formed dissolves in water to give HNO₃.
3NO₂ + H₂O(l) → 2 HNO₃ + NO (g)

Structure of HNO₃: In the gaseous state, HNO₃ exists as a planar molecule.

In an aqueous solution, it behaves as a strong acid.
HNO₃ (aq) + H₂O (Z) → (aq) + NO₃ (aq).
Concentrated HNO₃ is a strong oxidizing agent.

→ Reaction with Cu:
3Cu + 8 HNO₃ (dil) → 3 Cu (NO₃)₂ + 2 NO + 4 H₂O
Cu + 4 HNO₃ (cone.) → Cu (NO₃)₂ + 2NO₂ + 2H₂O

→ Reaction With Zinc
4 Zn + 10 HNO₃ (dil.) → 4 Zn (NO₃)₂ + 5 H₂O + N₂O
Zn + 4HNO₃ (cone.) → Zn (NO₃)₂ + 2H₂O + 2NO₂

→ Reaction with I₂

→ Reaction with Carbon
C + 4 HNO₃ → CO₂ + 2H₂O + 4NO₂

→ Reaction with Sulphur
S₈ + 48 HNO₃ (Cone.) → 8H₂SO₄ + 48 NO₂ + 12 H₂O

→ Reaction with Phosphorus
P₄ + 20 HNO₃ → 4H₃PO₄ + 20 NO₂ + 4H₂O

→ Brown Ring Test for nitrates (NO₃^–)
NO₃– + 3 Fe²⁺ + 4 H⁺ → NO + 3 Fe³⁺ + 2H₂O.

Uses of Nitric acid:

1. Oxidant in the laboratory.
2. manufacture of ammonium nitrate for fertilizers.
3. manufacture of explosive like TNT, nitroglycerine. Phosphorus: Phosphorus exists in three important allotropic forms: white, red and black.

White Phosphorus: White Phosphorus consists of discrete tetrahedral molecules as shown:

White phosphorus P₄

Red Phosphorus: Red Phosphorus is obtained by heating white phosphorus at 573 K in an inert atmosphere for several days. It is much less reactive than white phosphorus. It is polymeric, consisting of P4 tetrahedral linked together as shown in Fig. below.

Red Phosphorus

Black Phosphorus has two forms α-black phosphorus and β-black phosphorus. α-Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803 K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in the air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not bum in the air up to 673 K.

Phosphine (PH₃) Preparation:
1. From calcium phosphide and H₂O/dil. HCl
Ca₃P₂ + 6H₂O → 3Ca(OH)₂ + 2PH₃.
Ca₃P₂ + 6HCl → 3CaCl₂ + 2PH₃.

2. Lab. Method: By heating white phosphorus with cone. NaOH in an inert atom of CO₂.
P₄ + 3NaOH +- 3H₂O → PH₃ + NaH₂PO₂ (Sod. hypophosphite)

Properties:

1. Colourless gas with rotten fish smell, highly poisonous.
2. Explodes in contact with traces of oxidizing agents like HNO₃ Cl₂ etc.,
3. Slightly soluble in water.
4. When absorbed in CuSO₄ or HgCl₂ solutions, phosphides are obtained.
   3 CuSO4 + 2 PH₃ → Cu₃P₂ + 3H₂SO₄
   3HgCl₂ + 2PH₂  → Hg₂ P₂ + 6HCl.
5. PH₃ is weakly basic like NH₃.
   PH₃ + HBr → PH₄Br.

Phosphorus halides: P forms two types of halides, PX₃ [X = F, Cl, Br, I]and PX₅[X = F,Cl,Br]
Phosphorus trichloride (PCl₃)

Preparation:
P₄ + 6Cl₂ → 4PCl₃.

Properties:

1. Colourless oily liquid.
2. Hydrolysis in the presence of mixture.
   PCl₃ + 3H₂O → H₃PO₃ + 3HCl

Shape: It has a pyramidal shape as shown in which phosphorus is sp³ hybridized.

Shape of PCl₃

Phosphorus pentachloride (PCl5)
Preparation:

Properties:

1. It is a yellowish-white powder.
2. In moist air, it hydrolyses first to PoCl₃, and then H₃PO₄.
   PCl₅ + H₂O → POCl₃ + 2 HCl
   POCl₃ + 3H₂O → 4 H₃PO₄ + 3HCl
3. On heating, it sublimes but decomposes on stronger heating.
   

In gaseous and liquid phases, it has a trigonal bipyramidal structure shown below. The three equatorial P-Cl bonds are equivalent, while the two axial bonds are longer than equatorial bonds. This is due to the fact that the axial bond pairs suffer more repulsion as compared to equatorial bond pairs.

In the solid-state, it exits as an ionic solid, [PCl₄]⁺[PCl₆]^– in which the cation, [PCl₄]⁺ is tetrahedral and the anion, [PCl₆]⁺ octahedral.

Oxo-Acids of Phosphorus: P forms a number of oxo-acids like

1. Hypophosphorous acid (H₃PO₂) (ox. state + 1)
2. Orthophosphoric acid (H₃PO₃) (ox. state + 3)
3. Pyrophosphrus acid (H₄P₂O₅) (ox. state + 3)
4. Hypophosphoric acid (H₄P₂O₆) (ox. state +4)
5. Orthophosph’oric acid (H₃PO₄) (ox. state + 5)
6. Pyrophosphoric acid (H₄P₂O₇) (ox. state +5)
7. Metaphosphoric acid (HPO₃) (ox. state + 5)

The compositions of the oxo-acids are interrelated in terms of loss or gain of H₂O molecule or oxygen atom.
These acids in the + 3 oxidation state of P tend to disproportionate as follows:

The acids which contain the P-H bond are strong reducing agents.

Group 16 Elements: Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po) constitute Group 16 of the periodic table. They are also called Chalcogens (ore-forming).

→ Occurrence: Oxygen is the most abundant of all the elements of the earth. Dry air contains 20.95% oxygen by volume. However Sulphur is present in the earth’s crust to 0.03 – 0.1% only. Combined Sulphur is present in gypsum CaSO₄. 2H₂O, Epsom salt MgSO₄.7H₂O. Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores.

→ Electronic Configuration: The elements of Group 16 have six electrons in the outermost shell and have ns2np4 general electronic configuration.

→ Atomic and Ionic Radii: Due to the increase in the no. of shells, atomic and ionic radii increase from top to bottom, The size of the oxygen atom is, however, exceptionally small.

→ Ionisation Enthalpy: Ionisation enthalpy decreases down the group.

Table: Some Physical Properties of Group 16 Elements

^aSingle bond; ^bApproximate value; ^cAt the melting point; ^dRhombic sulphur; ^eHexagonal grey; ^fMonoclinic form, 673 K.
Oxygen shows oxidation states of + 2 and +1 in oxygen fluorides OF₂ and O₂F₂ respectively.

→ Electron Gain Enthalpy: Because of the compact nature of the oxygen atom, it has less negative electron gain enthalpy than sulphur.

→ Electronegativity: Next to F, oxygen has the maximum value of electronegativity value amongst the elements. Within the group, its value decreases from top to bottom implying that metallic character increases from O to Po.

→ Physical Properties: Oxygen and Sulphur are non-metals, Selenium and Tellurium metalloids, whereas Polonium is a metal. Pollonium is radioactive [t_1/2 = 13.8 days]

All these elements exhibit allotropy.
M.Pts and B.Pts. increase with the increase in atomic no. down the group. The large difference between the M.Pts and B.Pts. of oxygen and sulphur may be explained on the basis of their atomicity: Oxygen exists as diatomic molecules (O₂) whereas Sulphur exists as polyatomic (S₈).

Chemical Properties:
Oxidation states and trends in chemical reactivity:
O = – 2, -1, +1, (in O₂F₂), + 2 (in OF₂)
S = – 2, + 2, +4, + 6 (in SF₆)
Se = – 2, (+ 2), +4, + 6 (SeF₆)
Te=-2,(+2), +4, + 6(TeF₆)
Po = + 2, + 4
[Within brackets less stable.]

The stability of + 4 oxidation state increases and that of + 6 decreases due to the Inert-pair Effect. Bonding in + 4, + 6 oxidation states are primarily covalent.

The anomalous behaviour of oxygen is due to its

1. small size,
2. high electronegativity,
3. absence of d-orbitals.

1. Reactivity with Hydrogen.
(a) There is strong hydrogen bonding in H₂O which is not found in H₂S.
(b) All the elements of Group 16 form hydrides of the type H₂E[E = S, Se, Te, Po]
(c) H₂O is neutral. Other hydrides are acidic and acidic character increases.
H₂O < H₂S < H₂Se < H₂Te < H₂Po
(d) Thermal stability decreases:
H₂O > H₂S > H₂Se > H₂Te > H₂Po
(e) H₂O does do not have reducing character. Reducing character increases from H₂S < H₂Se < H₂Te.

2. Reactivity with Oxygen: All these elements form oxides of EO₂ and EO₃ types [E = S, Se, Te, or Po]. Both types are acidic in nature.

3. Reactivity towards halogens: Elements of Group 16 form a large no. of halides of the type, EX₆, EX₄, EX₂ where X = halogen. The stability of the halides decreases in the order F^– > Cl > Br > I^–

Dioxygen, O₂:

Industrially, O₂ is obtained from air by first removing CO₂ and water vapour and then, the remaining gases are liquified and fractionally distilled to give N₂ and O₂.

Properties of O₂:

1. Colourless, odourless gas.
2. 30.8 cm³3 of O₂ soluble per litre of water at 293 K.
3. It has three isotopes ¹⁶O, ¹⁷O, ¹⁸O.
4. Molecular oxygen (O₂) is paramagnetic.
5. Reactions with metals, non-metals and other compounds are as follows:
   
   

Uses:

• It is used in oxyacetylene welding.
• Manufacture of many metals, particularly steel.
• Oxygen cylinders are widely used in hospitals, high altitude flying and mountaineering.
• The contraction of fuels, e.g., hydrazines in liquid oxygen, provides tremendous thrust in rockets.

→ Simple Oxides: A binary compound of oxygen with another element is called oxide.

Oxides can be simple like MgO, Al₂O₃ or mixed like Pb₃O₄, Pe₃O₄.
In general metallic oxides like CaO, BaO, Na₂O are basic. Non-metal oxides like SO₂, CO₂ are acidic.

Some oxides like Al₂O₃ are amphoteric. They react with acids as well as bases.
Al₂O₃ + 6 HCl → 2AlCl₃ + 3H₂O.
Al₂O₃ + 2NaOH → 2Na AlO₂ + H₂O

Ozone, O₃: Ozone is an allotropic form of oxygen. At a height of 20 km from sear level, it is formed from atmospheric oxygen in the presence of sunlight.

The ozone layer protects the earth’s surface from UV rays.

Preparation:

ΔH° = + 142 kJ mol^-1

Properties:

1. It is a pale blue gas, dark blue liquid and violet-black solid.
2. O₃ is thermodynamically unstable w.r.t. O₂.
   ∴ It acts as a powerful oxidizing agent.

O₃ → O₂ + O
1.  It oxidizes lead sulphide to lead sulphate.
PbS + 4O₃(g) → PbSO₄ (s) + 4O₂

2. It oxidizes I^– ions to iodine
2I^– (aq) + H₂O (l) + O₃ (g) → 2 OH^– (aq) + I₂ (s) + O₂ (g)

3. It oxidizes NO to NO₂
NO (g) + O₃ (g) → NO₂ (g) + O₂ (g)

Structure: The two oxygen-oxygen bond lengths in the ozone molecule are identical (128 pm) and the molecule is angular as expected with a bond angle of about 117°. It is a resonance hybrid of two main forms:

Uses:

1. It is used as a germicide, disinfectant and for sterilising water.
2. It is also used for bleaching oils, ivory, flour, starch.
3. It is an oxidizing agent.

Sulphur-Allotropic Forms: Out of numerous allotropes, the two most important forms are:

• Yellow rhombic (α-sulphur),
• Monoclinic (β-Sulphur)
  

The temperature 369 K is called Transition temperature.

The structure of (a) S₈ ring in rhombic sulphur and (b) S₆ form:

Both rhombic and monoClinic Sulphur have S₈ molecules. These S₈ molecules are packed to give different crystal structures. The S₈ ring in both forms is puckered and has a crown shape.

In cyclo-S₆ the ring adopts the chair form. At elevated temperatures (~ 1000 K) S₂ is the dominant species and is paramagnetic like O₂

Sulphur Dioxide (SO₂):
Preparation:
1. By burning S in air or O₂.
S(s) + O₂(g) → SO₂(g)

2. Lab. Method: By treating a sulphite with dilute H₂SO₄4.
SO₃^2- (aq) + 2H⁺ (aq) → H₂O (l) + SO₂(g)

3. Industrially, it is produced by roasting of sulphide ores:

Properties:
1. It is a colourless gas with a pungent smell.

2. It is highly soluble in water.

3. It liquefies at room temp, under a pressure of 2 atmospheres.

4. Liquified SO₂ boils at 263 K.

5, SO₂ when dissolved in water, forms sulphurous acid.
SO₂(g) + H₂O (l) → H₂SO₃ (aq)

6. Reaction with NaOH:


7. Reaction with Cl₂:

8. Reaction with O₂:

9. Moist SO₂ behaves as a reducing agent. It reduces iron (III) ions to iron (II) ions and decolourised acidified potassium permanganate(VII) solution. It is a test for SO₂ gas.
2Fe³⁺ (aq) +SO₂ + 2H₂O → 2 Fe²⁺ (aq) + SO₄^2- + 4H⁺
5 SO₂ + 2 MnO₄^– + 2H₂O → 5 SO₄ ^2- + 4H⁺ + 2 Mn²⁺

Structure: The molecule of SO₂ is angular. It is a resonance hybrid of the two canonical structures

Uses:

• Refining of petroleum, sugar
• Bleaching wool and silk.
• Antichlor, disinfectant and preservative.
• Manufacture of H2S04 and other industrial chemicals.
• Liquid S02 is used as a solvent.

Oxo-Acids of Sulphur: Sulphur forms a number of oxo-acids such as H₂SO₂, H₂S₂O₃, H₂S₂O₄, H₂S₂O₅, H₂S_xO₆ (X = 2 to 5), H₂SO₄, H₂S₂O₇, H₂SO₅, H₂S₂O₈.

Structures of Sulpiwrous and Sulphuric acids

Sulphuric Acid (H₂SO₄): It is called the king of chemicals and is one of the most widely used industrial chemicals.

→ Manufacture: Sulphuric acid is manufactured by Contact Process which involves 3 steps:

1. Burning of Sulphur or Sulphide ores in the air to produce SO₂.
2. Conversion of SO₂ into SO₃ by reaction with air in the presence of catalyst V₂O₅.
3. Absorption of SO₃, in H₂SO_4, to give Oleum (H₂S₂O₇).
   The key step is the catalytical oxidation of SO₂ with O₂ to give SO₃ in the presence of V₂O₅.
   
   The reaction is exothermic, reversible and proceeds with a decrease in volume.
   ∴ Low temperature (720 K) and high pressure (2 bar) are favourable conditions for maximum yield.

The tempi should not be very low otherwise the rate of reaction will below.

The SO₃ gas from the catalytical converter is absorbed in concentrated H₂SO₄ to produced Oleum (H₂S₂O₇) which on dilution with water gives H₂SO₄ of the desired concentration.
SO₃ + H₂SO₄ → H₂S₂O₇ (oleum)
H₂S₂O₇ + H₂O → 2 H₂SO₄

The sulphuric acid obtained by the contact process is 96-98% pure.

Flow diagram the iiai1ufat true of sulphuric acid

Properties:

1. H₂SO₄ is a colourless, dense, oily liquid.
2. It dissolves in water producing a large amount of heat. Hence it should be slowly added to water (and never water in acid)

The chemical properties of H₂SO₄ are due to its
(a) low volatility,
(b) strong acidic character,
(c) strong affinity for water,
(d) ability to act as an oxidizing agent.

In water, it ionises in 2 steps:
1. H₂SO₄ (aq) + H₂O (l) → H₃O+ (aq) + HSO₄^– (aq)
K_a1 = very large (> 10)

2. HSO₄^–(aq) + H₂O (l) → H₃O⁺ (aq) + SO₄ ^2- (aq);
K_a2 = 1.2 × 10^-2.

(a) It is used to prepare more volatile acids:
NaNO₃ + H₂SO4 → NaHSO₄ + HNO₃.
2 NaCl + H₂SO₄ → Na₂SO₄ + 2 HCl

(b) Cone. H₂SO₄ is a strong dehydrating agent

(c) Hot cone. H₂SO₄ is a strong oxidising agent
H₂SO₄ → H₂O + SO₂ + O

(d) Reaction with metals and non-metals:
Cu + 2 H₂SO₄ → CuSO₄ + SO₂ + 2H₂O
C + 2H₂SO₄ (cone.) → CO₂ + 2SO₂ + 2H₂O

Uses of H₂SO₂ :

• It is used in the manufacture of hundreds of other compounds and also in many industrial processes.
• Manufacture of fertilizers.
• Petroleum refining and manufacture of pigments, paints and dyestuff intermediates.
• Detergent industry.
• Metallurgical applications.
• Storage batteries and so on.

Group 17 Elements: Fluorine, chlorine, .bromine, iodine and astatine are the members of Group 17. They are collectively called halogens (sea-salt producers)

→ Occurrence: Fluorine is present as CaF₂ (fluorspar) and Na₃ AlF₆ (cryolite). Seawater contains chlorides, bromides and iodides of sodium, potassium, magnesium and calcium. The deposits of dried up seas contain NaCl and carnallite, KCl. MgCl₂,6Fl₂O. Seaweeds contain 0.5% of iodine and Chile saltpetre contains up to 0.2% of sodium iodate.

→ Electronic Configuration: All these elements have seven electrons in their valence shell [ns²np³] which is 1 electron less than the next noble gas.

→ Atomic and ionic radii: The halogens have the smallest atomic radii in their respective periods due to the maximum effective nuclear charge. The size of the anions is greater than the corresponding atoms. The atomic and ionic radii increase from top to bottom.

→ Ionisation Enthalpy: They have little tendency to lose electrons. Thus they have very high ionisation enthalpy. I.E. decreases from top to bottom.

→ Electronegativity: They have high electronegativities. It decreases from top to bottom. Fluorine is the most electronegative element in the periodic table. [F = 4.0]

→ Electron Gain Enthalpy: Halogens have maximum negative- electron gain enthalpy in their corresponding periods. It becomes less negative down the group. Electron gain enthalpy of F is less than that of Cl due to the compact size of F.

→ Physical Properties: F₂ and Cl₂ are gases. Br₂ is a liquid and I₂ is solid. They are all diatomic. M.pts and B.pts increase from top to bottom. F₂ is yellow, Cl₂ is greenish-yellow, Br₂ is red and I₂ violet in colour.

F₂ has a smaller enthalpy of dissociation as compared to Cl₂.

→ Chemical Properties:
Oxidation States and Trends in Chemical Reactivity: All the halogens exhibit a -1 oxidation state. However, chlorine, bromine and iodine exhibit +1, + 3, +5 and +7 oxidation states also as explained below:

Table: Atomic and Physical Properties of Halogens (Group 17)

^aRadioactive; ^bPauling scale; ^cFor the liquid at temperatures (K) given in the parentheses; ^dsolid; ^eThe half-cell reaction is X₂(g) + 2e^– → 2X^–(aq).

The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms, e.g., in interhalogens, oxides and oxoacids.

The oxidation states of +4 and +6 occur in the oxides and oxoacids of chlorine and bromine. The fluorine atom has no d-orbitals in its valence shell and therefore cannot expand its octet. Being the most electronegative, it exhibits only a -1 oxidation state.

All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F2 is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the solid phase.

In general, a halogen oxidises halide ions of higher atomic number.
F₂ + 2X^– → 2F + X₂ (X = Cl, Br or I)
Cl^– + 2X^– → 2CT + X₂ (X = Br or I)
Br₂ + 2I^– → 2Br + I₂

The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials which are dependent on the parameters indicated below:

The relative oxidising power of halogens can further be illustrated by their reactions with water. Fluorine oxidises water to oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids.

The reaction of iodine with water is non-spontaneous. In fact, I^– can be oxidised by oxygen in an acidic medium; just the reverse of the reaction observed with fluorine.
2F₂ (g) + 2 H₂O (l) → 4H⁺ (aq) + 4 F (aq) + O₂ (g)
X₂ (g) + H₂O (l) → HX (aq) + HOX (aq) (where X = Cl or Br)
4I^– (aq) + 4H⁺ (aq) + O₂ (g) → 2I₂ (s) + 2 H₂O (l)

Anomalous behaviour of Fluorine: It is due to

1. Highest electronegativity and Ionisation enthalpy in Group 17.
2. Small size.
3. Non-availability of d-orbitals.
4. The low bond dissociation energy of the F-F bond.

1. Reactivity towards hydrogen: They all react with hydrogen to give hydrogen halides, but an affinity for hydrogen decreases from F to I.

The acidic strength is in order HF < HCl < HBr < HI and stability is under HF > HCl > HBr > HI.

2. Reactivity towards oxygen: Halogens form many oxides with oxygen but most of them are unstable. F forms OF₂ and O₂F₂, but only OF₂ is stable. Oxides of Cl, Br, I am powerful oxidizing agents.

3. Reactivity towards metals
Mg + Br₂ (l) → MgBr₂(s)
The ionic character of halids decreases in the order MF > MCI > MBr > MI.

4. Formation of interhalogen compounds: They form interhalogen compounds among themselves of the type AB, AB₃, AB₅, AB₇ where A = larger size halogen and B = smaller size halogen.

Chlorine, Cl₂:
Preparation:
1. By heating MnO₂ with cone. HCl

However a mixture of common salt and cone. H₂SO₄ is used in place of HCl.
2 NaCl + MnO₂ + 3 H₂SO₄ → 2 NaHSO₄ + MnSO₄ + 2H₂O + Cl₂

2. By the action of HC1 on KMn04
2 KMnO₄ + 16HCl → 2 KCl + 2 MnCl₂ + 5 Cl₂ + 8 H₂O

Manufacture of Cl₂:
1. Deacon’s process:

2. Cl₂ is obtained by the electrolysis of blue solution.

Properties:
1. Greenish-yellow gas with a pungent smell.

2. 2.5 times heavier than air.

3. Can be liquified easily into greenish-yellow liquid.

4. Soluble in water.

5. It reacts with a number of metals and non-metals and other compounds.
2Al + 3Cl₂ → 2Al Cl₃
2 Na + Cl₂ → 2NaCl
2 Fe +3 Cl₂ → 2FeCl₃
P₄ +6 Cl₂ → 4PCl
S₃ +4 Cl₂ → 4 S₂Cl₂
H₂ + Cl₂ → 2HCl
H₂S + Cl₂ → 2HCl + S
C₁₀H₁₆ + 8 Cl₂ → 16HCl + 10 C.
8NH₃ (excess) + Cl₂ → 6NH₄Cl + N₂
NH₂  + 3Cl₂ (excess) → NCl₃ + 3HCl

6. Chlorine dissolved in water on standing loses its yellow colour due to the formation of HCl and HOCl. Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties: Its oxidising properties are:

(a) 2 FeSO4 + H₂SO₄ + Cl₂ → Fe₂(SO₄)₃ + 2 HCl
Na₂SO₃ + Cl₂ + H₂O → 4 Na₂SO₄ + 2 HCl
SO₂ + 2H₂O + Cl₂ → 4 H₂SO₄ + 2HCl
I₂ + 6 H₂O + 5 Cl₂ → 2 HIO₃ + 10 HCl.

(b) It is a powerful bleaching agent due to oxidation. It is permanent.

Uses:

1. For bleaching wood pulp, cotton and textiles.
2. In the extraction of gold and platinum.
3. Manufacture of dyes, drugs, DDT, refrigerants etc.
4. In sterilising water.

Hydrogen Chloride, HCl:
Preparation-Lab.
Method:
2 NaCl + H₂SO₂  → Na₂SO₄ + 2 HCl cone.

Properties:

1. It is a colourless gas with a pungent smell.
2. Highly soluble in water.
3. Can be liquified easily to a colourless liquid and solidified to a white crystalline solid.
4. It ionises in an aqueous solution.
   HCl + H₂O (l) → 4 H₃O+ (aq) + Cl^– (aq); Ka = 107
   The high value of Ka shows it is a strong acid in water.
5. Reaction with NH₃
   NH₃ + HCl → NH₄Cl (White dense fumes)
6. 3 parts of cone. HCl and one part of cone. HNO₃ on mixing give aqua regia which is used to dissolve, Au, Pt.
   Au + 4H⁺ + NO₃ + 4 Cl^– → AUCl₄^– + NO + 2H₂O
   3Pt + 16H⁺ + 4NO₃^– + 18Cl^– → 3PtCl₆^2- + 4NO + 8H₂O .
7. Reaction with salts:
   Na₂CO₃ + 2HCl → 2 NaCl + H₂O + CO₂ ↑
   NaHCO₃ + HCl → NaCl + H₂O + CO₂ ↑
   Na₂SO₃ + 2HCl → 2 NaCl + H₂O + SO₂ ↑

Uses of HCl:

• Manufacture of Cl₂, NH₄Cl and glucose from com starch.
• Extracting glue from bones and purifying bone black.
• In medicine and as a lab. reagent

Table: Oxo-Acids of Halogens

Structures of Oxoacids:

Inter Halogen Compounds:

Some Properties of Interhalogen Compounds:

Uses:

• Interhafogen compounds are very useful fluorinating agents.
• They can be used as non-aqueous solvents.
• ClF₃ and BrF₃ are used for the production of UF₆ in the enrichment of ²³⁶U.
  U (s) + 3 ClF₃ (l) → UF₆(g) + 3 Cl F (g) argon (Ar)

Group 18 Elements: Group 18 Elements are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn). Because of ns2nph structure [1s² for He], They are chemically unreactive. They are called noble gases.

→ Occurrence: Except radon (Rn) which is radioactive, all noble gases occur in the atmosphere in traces. Noble gases constitute 1 % (by volume) of dry air. He is present in pitch blende, monazite, cleveite. Radon is obtained as a decay product of ²²⁶Ra.
²²⁶₈₈Ra → ²²²₈₆Ra + ⁴₂He

→ Electronic Configuration: All noble gases have ns²np⁶ electronic configuration except He which is 1s².

→ Ionisation Enthalpy: They have very high I.E. which decreases down the group.

→ Atomic radii: They increase down the group with an increase in at. number.

→ Electron Gain Enthalpy: They have large positive values because of their stable electronic configuration. They have no tendency to accept electrons.

Physical Properties:

1. These gases are monoatomic.
2. They are colourless, odourless and tasteless.
3. They are sparingly soluble in water.
4. Due to weak dispersion forces, they have very low m.pts and b. pts.
5. He has the lowest B.pt (4.2 K) of any known substance.

→ Chemical Properties: In general, noble gases are least reactive. It is due to

1. Stable electronic configuration of ns2np6 (or 1s² for He).
2. They have high Ionisation enthalpy and more positive electron gain enthalpy,

Neil Bartlett prepared the first compound of Xe: red coloured Xe⁺ Pt F₆^– by mixing PtF₆^– and Xenon. The compounds of Kr are fewer. He, Ne and Ar are not found to form any true compounds. Only KrF₂ has been studied. Compound involving Rn viz. RnF₂ has been identified but not isolated.

Xenon-fluorine Compounds: Xenon forms three binary compounds with fluorine: XeF₂, XeF₄, XeF₆ by the direct combination of elements under appropriate experimental conditions.

Xenon-oxygen compounds: Xenon and oxygen form XeO₃. Partial hydrolysis of XeF₄ and XeF₆ gives oxyfluorides XeOF₄ and XeO₂F₂.


structures of. (a) XeF₂, (b) XeF₄, (c) XeF₆, (d) XeOF₄ and (e) XeO₃
XeO₃ is a colourless explosive solid and has a pyramidal molecular structure. XeOF₄ is a colourless volatile liquid and has a square pyramidal molecular structure.

Uses: He is used in filling balloons because He is a light and non-inflammable gas. It is used in gas-cooled nuclear reactors. Liquid He is used in low temp, physics. It is used in the driving apparatus.

Neon is used in discharge tubes and fluorescent bulbs for display purposes. Argon is mainly used to provide an inert atmosphere in high-temperature metallurgical processes.

There are no significant uses of Xenon and Krypton. They are used in light bulbs designed for special purposes.

By going through these CBSE Class 12 Chemistry Notes Chapter 6 General Principles and Processes of Isolation of Elements, students can recall all the concepts quickly.

General Principles and Processes of Isolation of Elements Notes Class 12 Chemistry Chapter 6

The extraction and isolation of an element from its combined form involves various principles of chemistry. Metallurgy is the scientific and technological process of extracting a metal from its ore. The natural materials in which the metals or their compounds occur in the earth are called minerals. The mineral from which the metal is extracted conveniently and economically is called an ore. Ores are usually contaminated with earthly or undesired materials known as Gangue.

The extraction and isolation of metals from their respective ores involve the following major steps :

1. The concentration of the ore
2. Isolation of the metal from its concentrated ore, and
3. Purification of the metal.

Principle Ores of Some Important Metals:

Among metals, aluminium is the most abundant. For the purpose of extraction, bauxite is chosen for aluminium. For iron, usually oxide ores [Haematite Fe203] are taken. Before proceeding for concentration, ores are graded and crushed to a reasonable size.

Concentration Of Ores: Removal of the unwanted materials (e.g. sand, clays etc.) from the ore is called concentration, dressing, or benefaction. Unwanted impurities present in the ore are called gangue.

The nature of the impurities, the type of the metal and the environmental factors are taken into consideration.
(a) Hydraulic Washing: it is based on the difference in gravities of the ore and the gangue particles. It is, therefore, a type of gravity separation. When a stream of water is run through the powdered ore, the lighter gangue particles are washed away and heavier ore particles settle down.

(b) Magnetic Separation: This procedure is based on the difference in the magnetic properties of the ore and impurities present in it. When passed over the conveyer belt of a magnetic roller, magnetic particles settle in a heap nearer and non-magnetic impurities a bit away, as shown below:

Magnetic Separation

Froth Floatation Process: This process is used for removing gangue from sulphide ores only. Powdered sulphide ore is mixed with water to which Collectors (e.g. pine oil, fatty acid, xanthates etc.) enhance the non-wettability of the mineral particles and froth stabilizers (e.g. cresols, aniline) which stabilize the froth ore added. The mineral particles become wet by oils while the gangue particles by water.

A rotating paddle agitates the mixture and draws air in it. As a result, froth is formed which carries the mineral particles. The froth is light and is skimmed off from where ore particles are recovered after drying it

Froth Flotation Process

Leaching: This process consists of treating the powdered ore with a suitable reagent (such as acids bases or other chemicals) which can n selectively dissolve the ore but not the impurities.

In Baeyer’s process, pure aluminium oxide is obtained from the bauxite ore (which contains impurities of Fe2Os and silicates) by treating the powdered ore with a 45% solution of NaOH when alumina dissolves leaving behind impurities like Fe203 which are filtered off.

Na [Al(OH)4] or NaAlO., is neutralized by CO2 when Al(OH)3 gets precipitated.

Al(OH)3, obtained above is filtered, washed and finally heated to about 1473 K to get pure alumina Al₂O₃

Extraction of Crude Metal From Concentrated Ore: The concentrated ore must be converted into a form that is suitable for a reduction. Sulphide ore is usually converted into its oxide before reduction. Oxides are easier to reduce.

Two steps are involved
(a) Conversion of the concentrated ore into oxide ore
(b) Reduction of the metal oxide into metal

(a) Conversion To Oxide:
1. Calcination: It is the process of converting ore into its oxide by heating it strongly below its melting point either in the absence or limited supply of air. Volatile matter is driven off.

2. Roasting: It is the process of converting ore into its metallic oxide by heating strongly at a temperature insufficient to melt in excess of air.
2 ZnS + 3O₂ → 2 ZnO + 2 SO₂
2 PbS + 3O₂ → 2 PbO + 2SO₂
2 Cu₂S + 3O₂ → 2 Cu₂O + 2 SO₂

(b) Reduction of oxide to the metal: It is done with a suitable reducing agent (C or CO or even another metal)
M_xO + yC → xM + yCO

Thermodynamic Principles Of Metallurgy: The change in Gibbs energy ΔG is given by
ΔG = ΔH – TΔS …(1)
where ΔH = Enthalpy change at temperature T
ΔS = Entropy change at temperature T

Also ΔG° = -RT ln K. ……….(2)
where R = Gas constant;
K = Equilibrium constant.
The reaction will proceed if ΔG° is negative K in that case will be positive.

This happens only when the reaction proceeds towards products.
1. When the value of ΔG° is negative in equation (1) then only the reaction will proceed. If ΔS is positive, on increasing the temperature (T), the value of TΔS would increase (ΔH < TΔS) and then ΔG will become -ve.

2. If reactants and products of two reactions are put together in a system and the net ΔG of the possible reactions is – ve, the overall reaction will occur. So the process of interpretation involves the coupling of the two reactions, getting the sum of their ΔG and looking for its magnitude and sign. Such coupling is easily understood through Gibbs energy (ΔG°) vs – T plots for the formation of the oxides.

H.J. T Ellingham gave a graphical representation of Gibbs energy. It provides a sound basis for considering the choice of reducing agent in the reduction of oxides. Such a diagram helps us in predicting the feasibility of the thermal reduction of ore.

The reducing agent forms its oxide when the metal oxide is reduced. The role of the reducing agent is to provide ΔG° negative and large enough to make the sum of ΔG° of the two reactions (oxidation of the reducing agent and reduction of the metal oxide) negative.

As we know, during reduction, the oxide of a metal decomposes.
M_xO (s) → xM (Solid or Liquid) + \(\frac{1}{2}\) O₂(g) ……..(3)

The reducing agent takes away the oxygen.
xM (s or l) + \(\frac{1}{2}\) O₂ → MxO(s) ……[ΔG°(M, MxO) ………(4)

If reduction is carried out through equa lion (3), the oxidadion of the reducing agent (e.g. C or CO) will be as:
C(s) + \(\frac{1}{2}\) Oz(g) → CO(g) …………[ΔG(C, CO)]…..(5)
CO(g) + \(\frac{1}{2}\) Oz(g) → CO₂(g)……[ΔG(C, CO)]…(6)

If carbon is taken, there may be complete oxidation to CO2
\(\frac{1}{2}\) C(s)+ \(\frac{1}{2}\)O₂ (g) → \(\frac{1}{2}\)CO₂(g) ….[2 ΔG(C, CO)]…(7)

On subtracting equation (4), we get
M_xO(s) + C(s) → xM(s or l) +CO(g) …(8)
M_xO (s) + C (s) → xM (s or l) + CO₂ (g) …(9)
M_xO (s) + \(\frac{1}{2}\)C(s) → xM (s or l) + \(\frac{1}{2}\)CO₂ (g). …..(10)

These reactions describe the actual reduction of the metal oxide M_xO that is to be accomplished.
Increasing T (Heating) favours a negative value of Δ_rG°. Therefore, the temperature is chosen such that the sum of Δ_rG° in the two combined redox process is negative.

Applications:
(a) Extraction of Iron from its Oxides: Concentrated oxide ores of iron are mixed with limestone and coke and fed into a Blast furnace from the top. Here coke reduces oxide to the metal as follows:
FeO(s) + C(s) → Fe(s or l) + CO(g) ……….(11)

In two steps:
1. FeO(s) → Fe(s) + \(\frac{1}{2}\)O₂ (g) [ΔG(FeO, Fe)] …….(2)
2. C(s) + \(\frac{1}{2}\) O₂ (g) → CO(g) [ΔG(C, CO)] ……..(3)

From (12) and (13), The ne.t Gibbs energy change becomes
ΔG(C, CO) + ΔG(FeO, Fe) = Δ_rG ……….(14)

∴ The resultant reaction will take place when the r.h.s. of equation (14) becomes negative.

The reactions occurring in the Blast furnace at different temperatures are as follows:
At 500 – 800 K (lower temp, range)
3 Fe₂O₃ + CO → 2 Fe₃O₄ + CO₂
Fe₃O₄ + 4CO → 3 Fe + 4 CO₂
Fe₂O₃ + C0 → 2Fe + CO₂

At 900 -1500 K (higher temp, in the blast furnace)
C + CO₂ → 2CO
FeO + CO → Fe + CO₂

Limestone is also decomposed to CaO which removes silicate impurity of the ore as slag. The slag is in a molten state and separates out from iron.
CaCO₃ (s) → CaO (s) + CO₂ (g)
CaO (s) + SiO₂ (s) → CaSiO₂  (fusible slag)

Iron obtained from Blast furnace contains 4% carbon and many impurities in smaller amount (e.g. S, P, Si, Mn) is called pig iron. Cast iron is different from pig iron. It has a slightly lower carbon content (3%) and is extremely hard and brittle.

Blast Furnace

Further Reductions: Wrought iron or malleable iron is the purest form of commercial iron and is prepared from cast iron by oxidising

impurities in a reverberatory furnace lined with haematite. This haematite oxidises carbon to carbon monoxide :
Fe₂O₃ + 3 C → 2 Fe + 3CO

Limestone is added as a flux and sulphur, silicon and phosphorus are oxidised and passed into the slag. The metal is removed and freed from the slag by passing through rollers.

(b) Extraction of Copper from Cuprous Oxide [Copper (I) Oxide]: Most of the copper ores are sulphide ores, concentrated sulphide ores are roasted/smelted to give oxides.
2 Cu₂S + 3 O₂ → 2 Cu₂O + S

The oxide can then be easily reduced to give Cu metal.
Cu₂0 + C_coke → 2 Cu + CO

In the actual process, the ore is heated in a reverberatory furnace after mixing with silica.
In the furnace, iron oxide slags off’ as iron silicate and copper is produced in the form of copper matte which contains Cu₂S and FeS.
FeO + SiO₂ → Fe Si O₃(slag)

In the silica-lined convertor, copper matte is charged. Some silica is added and a hot air blast is blown when the following reactions take place.
2 FeS + 3O₂ → 2 FeO + 2SO₂
FeO + SiO₂ → FeSiO₃
2 Cu₂S + 3O₂ → 2 Cu₂O + 2 SO₂
2 Cu₂O + Cu₂S → 6 Cu + SO₂
The solidified copper obtained has a blistered appearance due to the evolution of SO₂ and so it is called blister copper.

(C) Extraction of Zinc from Zinc Oxide: Zinc oxide is reduced using coke

The metal is distilled off and collected by rapid chilling.

Electrochemical Principles Of Metallurgy: In electrolysis, metal ions in solution or molten form are reduced or by adding some reducing element. Here
ΔG° =. – n FE° ….(i)
where n = no. of electrons transferred
E° = Standard e.m.f. of the cell [redox couple]
F = Faraday = 96,500 C
ΔG° = Change in standard Gibbs energy.

More reactive metals have large negative values of the electrode potential. So their reduction is difficult. If the difference of two E° values corresponds to a positive E° and consequently negative ΔG° in the equation
1. above, then the less reactive metal will come out of the solution and more reactive metal will go into the solution.
Cu²⁺ (aq) + Fe (S) → Cu (S) + Fe²⁺ (aq)

Examples:
Metallurgy of Aluminium: Purified Al₂O₃ is mixed with Na₃ AlF₆ or CaF₂ to, lower the melting point of the mix and to make the solution conductive. The fused matrix is electrolysed.
2 Al₂O₃ + 3C → 4 Al + 3 CO₂

This process of electrolysis is widely known as the Hall-Fieroult process.

Steel cathode and graphite anode are used. The graphite anode is useful here for the reduction of oxide to the metal.

Electrolytic cell for the extraction of aluminium

The electrolysis of the molten mass is carried out in an electrolytic cell using carbon electrodes. The oxygen liberated at the anode reacts with the carbon of the anode producing CO and CO₂.This way for each kg of aluminium – produced, about 0.5 kg of carbon anode is burnt away.

The electrolytic reactions are:
Cathode: Al³⁺ (melt) + 3e^– → Al (l)
Anode: C(s) + O^2- (melt) → CO(g) + 2e^–
C(s) + 2 O^2- (melt) → CO₂ (g) + 4e^–

Copper from low-grade ores and scraps: Copper is extracted from low grades ores by Hydrometallurgy. It is leached out using acid or bacteria. The solution containing Cu²⁺ (aq) is treated with iron or H₂.
Cu²⁺ (aq) + H₂ (g) → Cu (s) + 2H⁺(aq)

Oxidation-Reduction: Some non-metals are extracted based on oxidation. Chlorine from prime solution is oxidized to Cl₂.
2 Cl^– (aq) + 2 H₂O (l) → 2 OH^– (aq) + H₂ (g) + Cl₂ (g)

In the extraction of gold and silver, metal is leached with CN^–.
Ag → Ag⁺ + e^– oxidation
Au → Au⁺ + e^– oxidation

The metal is later recovered by the displacement method.
4 Au (s) + 8 CN^– (aq) + 2H₂O (aq) + O₂ → 4 [Au (CN)₂]^– (aq) + 4 OH^– (aq)
2 [Au (CN)₂]^– (aq) + Zn (s) → 2 Au (s) + [Zn (CN)₄]^2- (aq)
In this reaction zinc acts as a reducing agent.

Refining: To obtain metal of high purity and to remove the last traces of impurities from the extracted metal, they are subjected to refining.

It is based upon the difference in properties of the metal and the impurity.
Several techniques are listed below:

1. Distillation,
2. Liquation,
3. Electrolysis,
4. Zone refining,
5. Vapour phase refining,
6. Chromatographic methods.

1. Distillation: This is useful for low boiling metals like Zn and Hg.

2. Liquation: A low melting metal like tin (Sn) can be made to flow on a sloping surface leaving higher melting impurities.

3. Electrolytic Refining: Impure metal is made anode. A strip of pure metal is made the cathode. The electrolyte (bath) contains soluble salt of the same metal.
Anode: M → Mn⁺ + e^–
Cathode: Mn⁺ + ne^– → M

Copper is refined by the electrolytic method. Here acidified solution of copper sulphate acts as an electrolyte.
Anode: Cu(s) → Cu²⁺ + 2e
Cathode: Cu²⁺ + 2e → Cu(s)

4. Zone Refining: This method is based upon the principle that the impurities are more soluble in the melt than in the solid state of the metal. The molten zone moves along with the mobile heater fixed at one end of the impure metal. As the heater moves forward the pure metal crystallises out of the metal and the impurities pass on into the adjacent molten zone. At one end, impurities get concentrated. This end is cut off.

Zone-Refining Process

Vapour Phase Refining: Mond’s Process for Nickel Refining

The carbonyl complex is subjected to a higher temperature so that it is decomposed to give pure metal.

van Arkel Method for Refining Zirconium or Titanium: This method is very useful for removing all the oxygen and nitrogen present in the form of impurity in certain metals like Zr and Ti. The crude metal is heated in an evacuated vessel with iodine. The metal iodide, being more covalent, volatilises.
Zr + 2I₂ → ZrI₄.

Chromatographic Methods: This method is based on the principle that different components of a mixture are differently adsorbed on an adsorbent.

Column Chromatography: It is very useful for the purification of the elements which are available in minute quantities and the impurities are not very different in chemical properties from the element to be purified.

There are several chromatographic techniques such as

1. Paper chromatography
2. Gas chromatography

Uses of Aluminium, Copper, Zinc and Iron: Alloys containing Al are light and are very useful. A1 wires conduct electricity. It is used as a reductant.

Copper is used to making wires in the electrical industry. Several alloys of copper with Zn, Sn and Ni are largely used.
Zinc is used for galvanising iron. It is used as a reducing agent. Similarly different forms of iron: Cast iron, Wrought iron, Steel find wide applications.

A Summary of the Occurrence and Extraction of Some Metals is Presented in the following Table:

By going through these CBSE Class 12 Chemistry Notes Chapter 5 Surface Chemistry, students can recall all the concepts quickly.

Surface Chemistry Notes Class 12 Chemistry Chapter 5

The branch of chemistry which deals with the nature of surface and species present on it is called surface chemistry. Adsorption on solid or on solution surfaces and colloidal properties are important surface effects.

Adsorption: The phenomenon of higher concentration of molecular species (gases or liquids) on the surface of solids than in the bulk is called adsorption.

The solid on the surface of which adsorption occurs is called adsorbent. The substances that get adsorbed on the solid surface is called adsorbate. The adsorbent may be a solid or a liquid and the adsorbate may be a gas or a liquid.

The phenomenon of absorption differs from adsorption as:

Types Of Adsorption:
Depending upon the nature of forces between the molecules of the adsorbate and the adsorbent, the adsorption maybe classified as: physical adsorption and chemical adsorption.
1. Physical adsorption: When the particles of the adsorbate are held to the surface of the adsorbent by the weak forces such as van der Waals forces, the adsorption is called physical adsorption or physisorption. The attractive forces are weak and therefore, these can be easily overcome either by increasing the temperature or by decreasing the pressure. In other words, physical adsorption can be easily reversed or decreased.

2. Chemical adsorption: When the molecules of the adsorbate are held to the surface of the adsorbent by the chemical forces, the adsorption is called chemical adsorption or chemisorption, hi this case, a chemical reaction occurs between the adsorbed molecules and the molecules or atoms of adsorbent.

Table: Comparison between Physisorption and Chemisorption:

Freundlich Adsorption Isotherm:
The adsorption of a gas on the surface of the solid depends upon the pressure of the gas. The extent of adsorption is generally expressed as x/m where m is the mass of the adsorbent and x is the mass of adsorbate when equilibrium has been attained. On the basis of experimental studies, Freundlich gave the following relationship between the amount of gas adsorbed (x) per unit mass of the adsorbent (m) and the pressure (p).
\(\frac{x}{m}\) = kp^1/n
where n is a constant (whole number) which depends upon the nature of adsorba te and adsorbent.

A graph between the amount (x/m) adsorbed by an adsorbent and the equilibrium pressure (or concentration for solutions) of the adsorbate at constant temperature is called Adsorption Isotherm.

At low pressure \(\frac{x}{m}\) ∝ p¹ ………..(i)

At high pressure \(\frac{x}{m}\) ∝ p⁰ ………..(ii)

In the intermediate range of pressure, combining (i) and (ii)
\(\frac{x}{m}\) ∝ p^0-1
∝ p1/n where n is an integer (n > 1)
or
\(\frac{x}{m}\) = kp^1/n ………..(iii)
where k is a constant depending upon the nature of the adsorbate and adsorbent.

This relationship is called Freundlich Adsorption Isotherm
Taking logs on both sides of (iii)
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log p

A graph between log \(\frac{x}{m}\) against log p should, therefore, be a straight line with slope equal to \(\frac{1}{n}\) and ordinate intercept equal to log K.

Adsorption From Solution Phase:
Solids can adsorb solutes from solutions so when a solution of acetic acid in water is shaken with charcoal, a part of the acid is adsorbed by charcoal and the concentration of the acid decreases in the solution.

The following conclusions have been made regarding adsorption from solution phase:

1. The extent of adsorption decreases with increase in temperature.
2. The extent of adsorption increases with the increase in the surface area of the absorbent.
3. The extent of adsorption depends upon the concentration of the solute in the solution.
4. The extent of adsorption depends upon the nature of the absorbent and the adsorbate.

Freundlich equation as applied to solutions is modified as \(\frac{x}{m}\) = k C^1/n  i.e., where C is the equilibrium concentration,
or
log \(\frac{x}{m}\) = log k + \(\frac{1}{n}\) log C

Applications of Adsorption:
1. Production of high vacuum: The remaining traces of air can be adsorbed by charcoal from a vessel evacuated by a vacuum pump to give a very high vacuum.

2. Gas masks: Gas mask (a device which consists of activated charcoal or mixture of adsorbents) is usually used for breathing in coal mines to adsorb poisonous gases.

3. Control of humidity: Silica and aluminium gels are used as adsorbents for removing moisture and controlling humidity.

4. Removal of colouring matter from solutions: Animal charcoal removes colours of solutions by adsorbing coloured impurities.

5. Heterogeneous catalysis: Adsorption of reactants on the solid surface of the catalysts increases the rate of reaction. There are many gaseous reactions of industrial importance involving solid catalysts. Manufacture of ammonia using iron as a catalyst, manufacture of H₂SO₄ by contact process and use of finely divided nickel in the hydrogenation of oils are excellent examples of heterogeneous catalysis.

6. Separation of inert gases: Due to the difference in degree of adsorption of gases by charcoal, a mixture of noble gases can be separated by adsorption on coconut charcoal at different temperatures.

7. In curing diseases: A number of drugs are used to kill germs by getting adsorbed on them.

8. Froth floatation process: A low grade sulphide ore is concentrated by separating it from silica and other earthy matter by this method using pine oil and frothing agent (see Unit 6).

9. Adsorption indicators: Surfaces of certain precipitates such as silver halides have the property of adsorbing some dyes like eosfn, fluorescein, etc. and thereby producing a characteristic colour at the end point.

10. Chromatographic analysis: Chromatographic analysis based on the phenomenon of adsorption finds a number of applications in analytical and industrial fields.

Catalysis: Potassium chlorate when heated strongly decomposes slowly giving dioxygen. The decomposition between 653 – 873 K.
2 KClO₃ → 2 KCl + 3O₂

However with a little of MnO a decomposition occurs only at 473 – 633 K and also at a much accelerated rate. MnO2 is a catalyst for this reaction.

A substance which accelerates the rate of a chemical reaction, itself remaining chemically and quantitatively unchanged after the reaction is called a catalyst.

Promoter is a substance that enhances the activity of a catalyst, while Poison is a substance that decreases the activity of a catalyst. In the reaction for the manufacture of NH₃ by Haber’s process.

Fe is catalyst and Molybdenum (MO) is a promoter.

Homogeneous And Heterogeneous Catalysis:
When the reactants and the catalyst are in the same phase (liquid or gas), the process is said to be homogeneous catalysis.

Example:

When the reactants and the catalyst are in different phases, the process is called heterogeneous catalysis.

Important Features of solid catalysts are:
(a) Activity: It depends upon the strength of chemisorption to a large extent. .
(b) Selectivity: Selectivity of a catalyst is its ability to direct a reaction to yield a particular product. For example, starting with H₂ and CO and using different catalysts, we get different products.

Thus a catalyst is highly selective in nature, i.e., a given substance can act as a catalyst only in a particular reaction and not for all the reactions. A catalyst for a particular reaction may fail to catalyse any other reaction.

Shape-Selective Catalysis by Zeolites: The catalytical reaction that depends upon the pore structure of the catalyst and the size of the reactant and product molecules is called shape selective catalysis. Zeolites are good shape-selective catalysts because of their honey comb¬like structure.

An important Zeolite catalyst used in the petroleum industry is ZSM-5. It converts alcohols directly into gasoline (petrol) by dehydrating them to give a mixture of hydrocarbons.

Enzyme Catalysis: Enzymes are complex nitrogeneous organic compounds which are produced by living plants and animals. They are actually protein molecules of high molecular mass, and form colloidal solutions in water. The enzymes are also referred to as Bio-Chemical Catalysts as they also occur in the bodies of animals and plants and such a phenomenon is known as Bio-Chemical Catalysis.

The following are examples of enzyme-catalysed reactions:
1. Inversion of cane-sugar

2. Conversion of glucose into ethyl alcohol.

3. Decomposition of urea into ammonia and CO₂.

Characteristics Of Enzyme Catalysis:
Enzyme catalysis is unique in its efficiency and high degree of specificity. The following characteristics are exhibited by enzyme catalysts:
1. Most highly efficient: One molecule of an enzyme may transform one million molecules of the reactant per minute.

2. Highly specific nature: Each enzyme is specific for a given reaction, i.e., one catalyst cannot catalyse more than one reaction. For example, the enzyme urease catalyses the hydrolysis of urea only. It does not catalyse hydrolysis of any other amide.

3. Highly active under optimum pH: The rate of an enzyme- catalysed reaction is maximum at a particular pH called optimum pH, which is between pH value 5-7.

4. Highly active under optimum temperature: The rate of an enzyme reaction becomes,maximum at a definite temperature, called the optimum temperature. On either side of the optimum temperature, the enzyme activity decreases. The optimum temperature range for enzymatic activity is 298-310 K. Human body temperature being 310 K is suited to enzyme-catalysed reaction.

5. Increasing activity in presence of activators and co¬enzymes: The enzymatic activity is increased in the presence of certain substances, known as co-enzymes. It has been observed that when a small non-protein (vitamin) is present along with an enzyme, the catalytic activity is enhanced considerably.

Activators are generally metal ions such as Na⁺ Mn²⁺, CO²⁺, Cu²⁺, etc. These metal ions, when weakly bonded to enzyme molecules, increase their catalytic activity. Amylase in presence of sodium chloride i.e., Na⁺ ions are catalytically very active.

6. Influence of inhibitors and poisons: Like ordinary catalysts enzymes are also inhibited or poisoned by the presence of certain substances. The inhibitors or poisons interact with the active functional groups on the enzyme surface and often reduce or completely destroy the catalytic activity of the enzymes. The use of many drugs is related to their action as enzyme inhibitors in the body.

Mechanism of Enzyme Catalysis: There are a number of cavities present on the surface of collodial particles of enzymes. These cavities are of characteristic shape and possess active groups such as – NH₂, – COOH, – SH, – OH etc. These are actually the active centres on the surface of enzyme particles. The molecules of the reactant (substrate), which have complementary shape, fit into these cavities just like a key fits into a lock.

An activated complex is formed which then decomposes to yield the products in two steps as outlined below:
Step I: E + S → ES*
Step II: ES* → E + P

Some Industrial Catalytical Processes:

Colloids: A colloid is a heterogeneous system in which one substance is dispersed (dispersed phase) as very fine particles in another substance called dispersion medium.

The essential difference between a solution and a colloid is that of particle size. Their size is in between that of true solution and suspension

Classification of Colloids: Colloids are classified on the basis of the following criteria:
(a) Physical state of dispersed phase and dispersion medium.
(b) Nature of interaction between dispersed phase and dispersion medium.
(c) Type of particles of the dispersed phase.
(a) Physical state of dispersed phase and dispersion medium: Depending upon whether the dispersed phase and the dispersion medium are solids, liquids or gases, eight types of colloidal systems are possible.

Types of Collodial Systems:

(b) Depending upon the nature of interactions between the dispersed phase and dispersion medium, the colloids can be classified as Lyophilic Colloids and Lyophobic Colloids.

1. Lyophilic collids: The colloidal solutions in which the particles of the dispersed phase have a great affinity (or love) for the dispersion medium are called lyophilic colloids. Such solutions are reversible in nature. In case water acts as the dispersion medium, the lyophilic colloid is called hydrophilic colloid. The common examples of lyophilic colloids are glue, gelatin, starch, proteins, rubber etc.

2. Lyophobic colloids: The colloidal solutions in which the particles of the dispersed phase have no affinity or love, rather have hatred for the dispersion medium, are called lyophobic collids. The solutions of metals like Ag and Au, hydroxides like Al (OH)₃ and Fe (OH)₃ and metal sulphides like As₂S₃ are examples of lyophobic colloids.
Such sols are formed with difficulty. They are irreversible in nature.

Multimolecular Macromolecular And Associated Colloids:
Depending upon the molecular size, the colloids can be classified as:
1. Multimolecular colloids: In this type, the particles consist of an aggregate of atoms or small molecules with molecular size less than 1 nm. For example, sols of gold atoms and sulphur (S₈) molecules. In these colloids, the particles are held together by Van der Waals forces.

2. Macromolecular colloids: In this type, the particles of the dispersed phase are sufficiently big in size (macro) to be of colloidal dimensions. In this case, a large number of small molecules are joined together through their primary valencies to form giant molecules.

These molecules are called macro molecules and each macromolecule may consist of hundreds or thousands of simple molecules. The solution of such moleucles are called macromolecular soluions. For example, colloidal solution of starch, cellulose, etc.

3. Associated colloids: These are the substances which behave as normal electrolytes at low concentration but behave, as colloidal particles at higher concentration. These associated particles are also called miscelles. For example, in aqueous solution, soap (sodium stearate) ionises as:

In concentrated solutions, these ions get associated to form an aggregate of colloidal size.

General Methods Of Preparation Of Sols:
Lyophilic sols are readily formed by simply mixing the dispersed phase and the dispersion medium under ordinary conditions.

Lyophobic sols can generally be prepared by two methods.
1. Condensation methods,
2. Dispersion methods.

→ Condensation methods: In these methods, the smaller particles are condensed suitably to be of colloidal size. This can be done by chemical reactions or by exchange of solvent.

→ Dispersion methods: In these methods, the large particles of a substance (suspension) are broken into smaller particles. This can be done by mechanical dispersion, by electrical dispersion or Bredig’s arc method and by peptisation.

Purification Of Colloidal Solutions:
The colloidal solutions prepared usually contain impurities especially electrolytes which can destabilize the sols. These impurities must be eliminated to make the colloidal solution stable.

The following methods are commonly used for the purification of colloidal solutions.
1. Dialysis: The method is based upon the fact that colloidal particles cannot pass through a parchment or cellophane membrane while the ions of the electrolyte can pass through it. The colloidal solution is taken in a bag made of cellophane or parchment.

The bag is suspended in fresh water. The impurities slowly diffuse out of the bag leaving behind pure colloidal solution. For example, dialysis can be used for removing HCl from the ferric hydroxide sol.

2. Electrodialysis: The ordinary process of dialysis is slow:

An apparatus for electrodialysis

To speed up the process of purification, the dialysis is carried out by applying electric field. This process is called electrodialysis.

3. Ultra-filtration: It is the process of removing the impurities from the colloidal solution by passing it through graded filter paper called ultrafilter papers. These filter papers are made from ordinary filter papers by impregnating them with colloidal solutions.

As a result, the size of the pores gets reduced. These filter papers allow the ions and molecules of the impurities to pass but retain colloidal particles. Ordinary filter papers cannot be used for this purpose since the colloidal particles also easily pass through the pores of these papers.

Properties Of Colloidal Solutions: The important properties of colloidal solutions are:
1. Heterogeneous nature: The colloidal solutions are heterogenous in nature consisting of dispersed phase and dispersion medium.

2. Visibility: The colloidals are not visible to naked eye and these can be seen with ultra microscopes.

3. Brownian movement: The colloidal particles have continuous zigzag motion called Brownian movement.

Brownian Movement

4. Tyndall effect: When the light is passed through the colloidal solution, the path of the light becomes visible when viewed from a direction at right angle to the incident beam. This phenomenon was studied by Tyndall and is known as Tyndall effect. The phenomenon of scattering oflight’by colloidal particles as a result of which the path of the beam becomes visible is called Tyndall effect.

5. Electrical properties: The particles of the colloidal solutions possess electrical charge, positive or negative. The presence of charge is responsible for the stability of these solutions. It may be noted that only the sol particles carry some charge while the dispersion medium has no charge. For example, the collodial solutions of gold, arsenious sulphide (AS₂S₃) are negatively charged while those of Fe (OH)₃ and Al (OH)₃ have positive charge. In the case of silver chloride sol, the particles may either be positively or negatively charged.

The presence of the charge on the sol particles and its nature whether positive of negative can be determined with the help of a phenomenon known as electrophoresis. In this experiment, the colloidal particles move towards positive or negative electrodes depending upon their charge under the influence of electrical field.

The phenomenon of movement of colloidal particles under an applied electric field is called electrophoresis.

If the particles accumulate near the negative electrode, the charge on the particles is positive. On the other hand, if the sol particles accumulate near the positive electrode the charge on the particles is negative.

A set up for electrophoresis

Origin of charge: The charge on the colloidal particles may be due to selective adsorption of ions. The particles contributing the dispersed phase adsorb only those ions preferentially which are common with their own lattice ions. For example, if silver nitrate solution is added to an aqueous solution of potassium iodide, the silver iodide will adsorb negative ions (I^–) from the dispersion medium to form a negatively charged sol.

However, if silver iodide is formed by adding potassium iodide to silver nitrate solution, the sol will be positively charged due to the adsorption of Ag⁺ ions present in the dispersion medium.

Coagulation Of Colloidal Solution: The phenomenon of precipitation of a colloidal solution by the addition of excess of an electrolyte is called coagulation or floculation.

→ Factors governing coagulation
1. Nature of the electrolytes: The coagulation capacity of different electrolytes depends upon the valency cf the active ion or called floculating ion. It is the ion carrying charge opposite to the charge on the colloidal particles. According to Hardy Schulz law, greater the valency of the active ion or floculating ion greater will be its coagulating power. Thus, to coagulate negative sol of As₂S₃, the coagulating power of different cations has been found to decrease in the order as:
Al³⁺ > Mg²⁺ > Na⁺

Similarly, to coagulate a positive sol such as Fe (OH)₃, the coagulating power of different anions has been found to decrease in the order:
[Fe(CN)₆]^4- > PO₄ ^3- > SO₄^2- > Cl^–

The minimum concentration of an electrolyte which is required to cause the coagulation or flocculation of a sol is known as flocculation value. It is usually expressed as milli moles per litre.

Protection of Colloids: The process of protecting the lyophobic colloidal solution from precipitation by the electrolytes due to the previous addition of some lyophilic colloid is called protection. The colloid which is added to achieve such a protection is called protecting colloid.

Gold Number: The different protecting colloids differ in their portecting powers. Zsigmondy introduced a term called gold number to describe the protective power of different colloids. This is defined as the minimum number of milligrams of the protective colloid required to just prevent the coagulation of a 10 ml of a given gold sol when 1 ml of a 10% solution of sodium chloride is added to it.

The coagulation of gold sol is indicated by change incolour from red to blue. The gold number of a few protective colloids are as follows:

It may be noted that smaller the value of gold number, greater will be protecting power of the protective colloid. Therefore, reciprocal of gold number is a measure of the protective power of a colloid. Thus, out of the list given above, gelatin is the best protective colloid.

Emulsions: Emulsions are the colloidal solutions of two immiscible liquids in which the liquid acts as the dispersed phase as well as the dispersion medium. Normally they are obtained by mixing an oil with water. Since the two do not mix well, the emulsion is generally unstable and is stabilised by adding a suitable outside reagent called emulsifier or emulsifying agent. The substances that are commonly employed for the purposes are gum, soap, glass powder, etc.

Types of Emulsions: These are of two types:
1. Oil-in-water emulsions: In this case, oils acts as the dispersed phase (small amount) and water as the dispersion medium (excess) e.g., milk is an emulsion of soluble fats in water and here casein acts as an emulsifier. Vanishing cream is another example of this class. Such emulsions are called aqueous emulsions.

2. Water-in-oil emulsions: In this case water acts as the dispersed phase while the oil behaves as the dispersion medium e.g., butter, cod liver oil, cold cream etc. Such types of emulsions are called oily emulsions.

Types of Emulsions

Demulsification: It is the process of decomposing an emulsion back into its constituent liquids. The demulsification can be done by centrifugation, filtration, boiling, freezing and some chemical methods.

Following are the interesting noteworthy examples of colloids which we come across in daily life.

1. Blue colour of the sky: Dust particles along with water suspended scatter blue light due to which sky looks blue to, us.
2. Fog, mist and rain: Clouds are aerosols. It is possible to cause artificial rain by throwing electrified sand.
3. Food articles like milk, butter, fruit juices are all colloids.
4. Blood-: Blood is a colloidal solution of an albuminoid substahce. Alum and FeCl₃ solution stop oozing blood due to coagulation. ,
5. Soils: Fertile soils are colloidal in nature in which humus . acts as a protective colloid.
6. Formation of delta: River water is a colloidal solution of clay.

Sea water contains several electrolytes. When river water meets sea water, the electrolytes present in sea water, coagulate the colloidal solution of clay resulting in its deposition with the formation of delta.

Applications of colloids: Colloids are widely used in the industry.
Some examples are:
1. Cottrell Smoke Precipitator: Smoke is a colloidal solution of solid particles such as carbon, arsenic compounds, dust etc. in air. The smoke is led through a chamber containing platgs having a charge opposite to that carried by smoke particles.” The particles on coming in contact with these plates lose their charge and get precipitated. The particles thus settle down on the floor of the chamber.

Cottrell Smoke Precipitator

2. Purification of drinking water: Alum is added to water to coagulate the suspended impurities present in water from natural resources to make it fit for drinking purposes.

3. Medicines: Most of the medicines are colloidal in nature. . Milk of magnesia, an emulsion, is used for stomach disorder. Argyrol is a silver sol used as an eye lotion. Colloidal antimony is used in curing Kalazar. Colloidal gold is used for intramuscular injection. Colloidal medicines are more effective because of large surface area and hence are easily assimilated.

4. Tanning: Animal hides due to colloidal nature bear positive charge when it is soaked in tanning or chromium salts which Contain negatively charged colloidal particles, mutual coagulation takes place leading to hardening of leather. The process is called tanning.

5. Cleansing action of soaps and detergents: Soaps and detergents act as emulsifiers and remove greasy impurities and dust of clothes which is washed away.

6. Photographic plates & films are prepared by coating an emulsion of light sensitive AgBr in gelatin over glass plates or celluloid films.

7. Rubber industry: Latex is a colloidal solution of rubber particles which are negatively charged. Rubber is ob tained by coagulation of latex.

8. Industrial products: Paints, inks, synthetic plastics and rubber, graphite lubricants, cement etc. are all colloidal solutions.

By going through these CBSE Class 12 Chemistry Notes Chapter 4 Chemical Kinetics, students can recall all the concepts quickly.

Chemical Kinetics Notes Class 12 Chemistry Chapter 4

Thermodynamics tells us about the feasibility of a reaction. A reaction is feasible if ΔG < 0 at constant temperature and pressure. Chemical equilibrium tells us the extent to which a reaction can proceed. But none of them tell us the speed of a reaction, i.e., time taken by a reaction to reach equilibrium.

The branch of chemistry which deals with the rates of reactions and their mechanisms is called Chemical Kinetics.
Rate Of A Chemical Reaction

From a kinetics point of view, reactions can be classified into 3 categories:
1. Very fast reactions: Which take place instantaneously e.g., ionic reactions like
AgNO₃ (aq) + NaCl (aq) → ↓AgCl (s) + NaNO₃ (aq)
BaCl₂ (aq) + H₂SO₄ (aq) → BaS04 (s) + 2 HCl (aq)
NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)

2. Very slow reactions: Which may take days or months for

Decomposition of NH₄NO₂

Decomposition of Hydrogen peroxide

Rate Of A Chemical Reaction: The rate of the reaction means the speed of the reaction.
It is defined as the change in the concentration of any one of the reactants or products per unit time.
Rate of the reaction = \(\frac{\text { Decrease in the concentration of a reactant }}{\text { Time taken }}\)
= \(\frac{\text { Increase in the concentration of a product }}{\text { Time taken }}\)

Consider a hypothetical reaction, R → P assuming the volume of the system is constant. One mole of reactant R produces one mole of the product P. If [R]₁ and [P]₁ are the concentrations of R and P at time t₁ and [R]₂ and [P]₂ are their concentrations at time t₂ then
ΔT = t₂ – t₁
Δ[R] = [R]₂ – [R]₁
Δ[P] = [P]₂ – [P]₁
The square brackets are used to express molar concentrations.
Rate of disappearance of R = – \(\frac{\Delta[\mathrm{R}]}{\Delta \mathrm{t}}\) …….(i)
Rate of appearance of P = + \(\frac{\Delta[\mathrm{P}]}{\Delta \mathrm{t}}\) ………(ii)

Equations (i) and (ii) given above represent the average rate of a. reaction, r_ave. r_ave depends upon (a) change in the cone, of reactants or products and (ii) time taken for that change to occur.

Units of Rate of a Reaction: Units of rate are concentration time-1, i.e., mol L^-1 S^-1. However for gaseous reactions, when the concentrations are expressed in terms of partial pressure, then the units of the rate will be atm s^-1.

For example,
PCl₅ → PCl₃ + Cl₂

Any general reaction
A + B → C + D

For this reaction, the rate of decomposition (disappearance) of

Rate of formation of O₂ = + \(\frac{\mathrm{d}\left[\mathrm{NO}_{2}\right]}{\mathrm{d} \mathrm{t}}\)

Rate law: The dependence of the concentration of reactants on the rate of reaction is given by the law. It gives the experimental dependence on concentration of reactants. For example, for the reaction
aA + bB → Products

From the kinetic study of the reaction, the dependence of the concentration of reactants on the rate of the reaction has been found to be

Rate = k[A]^m[B]ⁿ
where m and n are constant numbers or the powers of the concentration terms of the reactants A and B respectively on which the rate of reaction depends. It may be noted that the values of m and n are determined experimentally and may or may not be equal to a and b coefficients in the reaction. The above expression is the rate law. It may be defined as:

The mathematical expression denotes the observed or actual rate of a reaction in terms of the molar concentration of the reacting species which influences the rate of the reaction.

Order And Molecularity Of A Reaction:
Order of a Reaction: The dependence of reaction rate on concentration may be expressed in terms of the order of a reaction. The order of a reaction is defined as

the sum of the powers to which the concentration terms in the rate law are raised to express the observed rate of a reaction.
Thus, if the rate of a reaction,
xA + yB + zC → Products is given by the rate law,

Rate = –\(\frac{\mathrm{d} \mathrm{x}}{\mathrm{dt}}\) = k[A]x [B]y [C]z then, the order of the reaction, n is
n = x + y + z
where x, y and z are the order with respect to individual reactants and overall order of the reaction is sum of these exponents, i.e., x + y + z.

For example,
1. decomposition of ammonium nitrite is a first-order reaction.
NH₄NO₂ → N₂ + 2H₂O
Rate = k [NH₄NO₂] Order = 1

2. decomposition of hydrogen iodide is a second order reaction
2HI → H₂ + I₂
Rate = k[HI]₂ Order = 2

3. the reaction between nitric oxide and oxygen is a third-order reaction.
2NO (g) + O₂ (g) → 2NO₂ (g)
Rate = k[NO]²[O₂] Order = 3

The order of a reaction may be fractional or even zero.
For example,
CH₃CHO → CH₄ + CO
Rate = k[CH₃CHO]^3/2 Order = 3/2

Rate = k [NH₃]° Order = 0

Molecularity of Reaction: The molecularity of a reaction is defined as the number of reacting molecules that collide simultaneously to bring about a chemical reaction.

For example, the decomposition of H₂O₂ involves only one molecule, so it is a unimolecular reaction.
H₂O₂ → H₂O + \(\frac{1}{2}\)O₂

Dissociation of hydrogen iodide is a biomolecular reaction.
2HI (g) → H₂ (g) + I₂ (g)

The reaction involving three or more than three molecules is uncommon because the occurrence of such reactions would require simultaneous collisions of three or more than three molecules. The chances of the occurrence of such collisions are very small.

However, some of these reactions are found to be quite fast. This means that even though the balanced equation involves a large number of molecules, yet the reaction does not proceed by the simultaneous collision of all these reacting particles.

Such types of reactions take place through a sequence of two or more consecutive steps and are called complex reactions. The detailed description of various steps by which reactants change into the products is called the mechanism of the reaction. The steps which contribute to the overall reaction are called elementary processes.

Mechanism and Rate Law: In the case of multi-step reactions, some of the steps will be very fast while others will be slow. If one step takes place much more slowly than all other steps, it controls the overall reaction rate. This means that all the steps have to wait for the occurrence of this slowest step. But once this slowest step has occurred, the other steps will take place to form the products. In other words, the rate of the reaction is determined by the slowest step in the sequence.

Let us consider reaction between NO₂ and F₂ to form NO₂F;
2NO₂(g) + F₂(g) → 2NO₂F(g)

The experimental observations reveal that the rate of the reaction is proportional to the product of the concentrations of nitrogen peroxide and fluorine. This indicates that the rate-determining step in the mechanism of this reaction must be the reaction between NO₂ and F₂ only. Keeping this in mind, a mechanism of this reaction may be suggested as:

The rate of the overall reaction is determined by the first step which is the slower of the two steps. Accordingly, the experimentally observed rate of the reaction is given by the expression :
fix
Rate = – \(\frac{\mathrm{d} x}{\mathrm{dt}}\) = k[NO₂][F₂]

The above is the rate law for the reaction. It is, therefore, evident that complex reactions involving more than three molecules in the stoichiometric equation must take place in more than one step.
e.g.KClO₃ + 6 FeSO₄ + 3 H₂SO₄ → KCl + 3Fe₂ (SO₄)₃ + 3H₂O
should be, apparently, of the tenth order is actually a second-order reaction. This shows that the reaction occurs in several steps. The slowest step is called the rate-determining step.

Consider

The rate equation for this reaction is found to be
Rate = – \(\frac{\mathrm{d}\left[\mathrm{H}_{2} \mathrm{O}_{2}\right]}{\mathrm{dt}}\) = k [H202][I-]

This reaction is of first order w.r.t. both H2Oz and I-. Evidences suggest that this reaction takes place in 2 steps :

1. H₂O₂ + I → H₂O + IO^– – Slow Step
2. H₂O₂ + IO^– → H₂O + I^– + O₂ – Fast Step

∴ The first step, being slow, is the rate determining step.

It can be concluded from above

1. The order of a reaction is an experimental quantity. It can be zero and even a fraction, but molecularity cannot be zero or a non-integer.
2. Order is applicable to elementary as well as complex reactions whereas molecularity is applicable only for elementary reactions. For complex reactions, molecularity has no meaning/significance.
3. For a complex reaction, the order is given by the slowest step and generally, the molecularity of the slowest step is the same as the order of the overall reaction.

Units of Reaction Rate Constants: The rate is the change in concentration with time. Therefore, the rate of reaction is expressed by concentration units divided by time. If the concentrations are expressed in mol litre^-1 and time in seconds, then units of rate constants for different orders are:

• Units for first-order reaction = sec^-1
• Units for second-order reaction = litre mol^-1 sec^-1
• Units for third-order reaction = litre² mol^-2 sec^-1

In the case of gaseous reactions, if concentrations are expressed in units of atm, then

• Units of rate constant for the first-order reaction = sec^-1
• Units of rate constant for second-order reaction = atm-1 sec^-1
• Units of rate constant for third-order reaction = atm^-2 sec^-1

Rate Constant For First Order Reaction:
For a first-order reaction,
A → Products
Let the initial cone, of [A] be an ML^-1.

Let x change into products so that equilibrium concentration after time t is


If the initial concentration is ‘a’ moles per liter, x moles of A change in time t and k is the rate constant, then the integrated rate equation is

where [A]₀ is the initial concentration and [A] is the concentration at time t. The value of k can be calculated by substituting the values of a, t, and x.

A plot between In (A) and t gives a straight line with a slope equal to – k.

Half-Life Period Of A Reaction: The half-life period of a reaction is defined as the time during which the concentration of a reactant is reduced to one-half of its initial concentration. It is generally denoted as t_1/2. The half-life period of a first-order reaction may be calculated as given below:

For the first-order reaction,
t = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{[\mathrm{A}]_{0}}{[\mathrm{~A}]}\)

Now half-life period corresponds to the time during which the initial concentration, [A]₀ = a, is reduced to half i.e. [A] = a/2

The half-life period, t_1/2becomes

Thus, half life period of a first-order reaction is independent of the initial concentration of the reactant.

Similarly, the relation for the time required to reduce the concentration of the reactant to any fraction of the initial concentration can be calculated. For example,
∴ t_3/4 = \(\frac{2.303}{\mathrm{k}}\) log \(\frac{\mathrm{a}}{\mathrm{c} / 4}\) = \(\frac{2.303}{\mathrm{k}}\) log 4

Zero order Reactions are those reactions in which the rate of the reaction is proportional to zero power of the concentration of the reactants.
Rate = – \(\frac{\mathrm{d}[\mathrm{R}]}{\mathrm{dt}}\) = k[R]_o = k
or
R = \(\frac{[\mathrm{R}]_{0}-[\mathrm{R}]}{\mathrm{t}}\)

Collision Theory of Reaction Rate: The number of collisions that take place per second per unit volume of the reaction mixture is known as collision frequency (Z). All the collisions are not effective and do not give products. The collisions which actually produce the products and therefore, result in chemical reactions are called effective collisions.

There are two conditions for effective collisions.
1. Energy barrier: For the reacting species to make effective collisions, they should have sufficient energy to break the chemical bond in the reacting molecules. The minimum amount of energy that the colliding molecules must possess is known as threshold energy. Thus, only those collisions of reactants will give products that possess energies greater than threshold energy.

2. Orientation barrier: The colliding molecules should also have proper orientation so that the old bonds may break and new bonds are formed.

Thus, the collisions in which the colliding molecules do not possess the threshold energy of proper orientation do not form products. Therefore, only a small fraction of collision is effective.

Orientation of colliding NO₂ molecules is proper in (a) but ‘not’ in (b) for the reaction to take place.

Dependance Of Reaction Rate On Temperature: Temperature has a great influence on reaction rates. In general, the rate of a reaction becomes almost double for every 10° rise in temperature. The increase in reaction rate is not due to an increase in collision frequency but it is due to an increase in the fraction of effective collisions. It has been found that the fraction of effective collisions becomes almost double for a 10° rise in temperature.

Arrhenius Equation: The quantitative relationship between the rate constant and temperature was proposed by Arrhenius known as the Arrhenius equation.
k = A e^-E_a/RT

where k is a constant called frequency factor, E_a is the activation energy. If k₁ and k₂ are rate constants at two different temperatures T₁ and T₂ respectively then, integrated form of Arrhenius equation is
log\(\frac{\mathrm{k}_{2}}{\mathrm{k}_{1}}=\frac{\mathrm{E}_{\mathrm{a}}}{2.303 \mathrm{R}}\left[\frac{\mathrm{T}_{2}-\mathrm{T}_{1}}{\mathrm{~T}_{1} \mathrm{~T}_{2}}\right]\)

Activation Energy: The excess of energy (over and above the average energy of the reactants) required by the reactants to undergo chemical reactions is called activation energy. It is equal to the difference between threshold energy needed for the reaction and the average energy of reactant molecules.

Activation Energy = Threshold energy – Average energy of all reacting molecules.

When the molecules possess the energy equal to E , the atomic configuration of species formed at this stage is different from the reactants as well as the products. This stage is called the activated state or the transition state and specific configuration of the state is called Activated Complex. For example in the reaction between H₂ (g) and I₂ (g), activated complex has configuration in which H-H arid 1-1 bonds are breacking and H-I bonds are forming as shown below.

The change of reactants to products, i.e., progress of a reaction is shown below.

E_a = Activation energy for forward reaction
E_a‘ = Activation energy for backward reaction

Effect Of Catalyst On Reaction Rate:
The substances which increase the rate of a reaction and can be recovered chemically unchanged in mass and composition after the reaction are called catalysts. The phenomenon of increasing the rate of reactions by the use of catalyst is known as catalysis.

The function of a catalyst is that it provides a new path for the reaction, in which the reactants are converted into products quickly. It is believed that the catalyst forms a new activated complex of lower potential energy. This means that the activation energy becomes lower for the catalysed reaction than that for uncatalysed reaction.

Consequently, the fraction of the total number of collisions possessing lower activation energy is increased and hence, the rate of reaction also increases. This is shown below. The solid lines shows the path for uncatalysed reaction and dotted line shows the path adopted by catalysed reaction.

E_a = Activation energy without catalyst
E’_a (Cat) = Activation energy with catalyst

Effect Of Radiation On Reaction Rate:
Rate of reaction is increased by use of certain radiations. Such type of reactions which are initiated by the absorption of radiation are called photochemical reactions. For example, reaction of hydrogen and chlorine takes pla,ce very slowly in the absence of light. However, in the presence of light, the reaction occurs rapidly.

Reversible Reactions: At any stage, the net rate of the reaction is determined by the difference in the rates of forward and backward reactions. At equilibrium the overall rate of reaction in either direction becomes zero. Therefore,
Net rate of reaction = Rate of forward reaction – Rate of backward reaction At equilibrium, net rate is zero, so that

Rate of forward reaction = Rate of backward reaction
For the simple reaction A + B ⇌ C + D

If the rate constant for the forward reaction is Rf and that for the backward reaction is kf then.
Rate of forward reaction a [A] [B] = k [A] [B]
Similarly, rate of backward reaction a [C] [D] = kb [C] [D]
∴ At equilibrium rate of forward reaction = rate of backward reaction

where K = Equilibrium constant
Thus, K = \(\frac{[C] \times[D]}{[A] \times[B]}\)