d-Block elements

A transition element may be defined as an element whose atom in the ground state or ion in common oxidation state has incomplete sub-shell, has electron 1 to 9. It is called transition element due to fact that it is lying between most electropositive (s-block) and most electronegative (p-block) elements and represent a transition from them. The general electronic configuration of these element is (n–1)^{1 to 10} ns ^{0 to 2}

The definition of transition metal excludes Zn, Cd and Hg because they have complete d- orbital. Their common oxidation state is Zn⁺⁺, Cd^++, Hg⁺⁺ They also do not show the characteristics of transition element. Element of group 3 (Sc, Y, La and Ac) and group 12 (Zn, Cd, Hg) are called non typical transition element.

First transition or 3d series :

Element	Symbol	At. No.		Electronic configuration
Scandium	Sc	21	3d-orbitals are filled up	[Ar] 3d¹4s²
Titanium	Ti	22		[Ar] 3d²4s²
Vanadium	V	23		[Ar] 3d³4s²
Chromium	Cr^*	24		[Ar] 3d⁵4s¹
Manganese	Mn	25		[Ar] 3d⁵4s²
Iron	Fe	26		[Ar] 3d⁶4s²
Cobalt	Co	27		[Ar] 3d⁷4s²
Nickel	Ni	28		[Ar] 3d⁸4s²
Copper	Cu^*	29		[Ar] 3d¹⁰4s¹
Zinc	Zn	30		[Ar] 3d¹⁰4s²

Second transition or 4d-series :

Element	Symbol	At. No.		Electronic configuration
Yttrium	Y	39	4d-orbitals are filled up	[Kr] 4d¹5s²
Zirconium	Zr	40		[Kr] 4d²5s²
Niobium	Nb^*	41		[Kr] 4d⁴5s¹
Molybdenum	Mo^*	42		[Kr] 4d⁵5s¹
Technetium	Tc	43		[Kr] 4d⁵5s²
Ruthenium	Ru^*	44		[Kr] 4d⁷5s¹
Rhodium	Rh^*	45		[Kr] 4d⁸5s¹
Palladium	Pd^*	46		[Kr] 4d¹⁰5s⁰
Silver	Ag^*	47		[Kr] 4d¹⁰5s¹
Cadmium	Cd	48		[Kr] 4d¹⁰5s²

Third transition or 5d-series :

Element	Symbol	At. No.		Electronic configuration
Lanthanum	La	57	5d-orbitals are filled up	[Xe] 5d¹6s²
Hafnium	Hf	72		[Xe] 4f¹⁴5d²6s²
Tantalum	Ta	73		[Xe] 4f¹⁴5d³6s²
Tungsten	W	74		[Xe] 4f¹⁴5d⁴6s²
Rhenium	Re	75		[Xe] 4f¹⁴5d⁵6s²
Osmium	Os	76		[Xe] 4f¹⁴5d⁶6s²
Iridium	Ir	77		[Xe] 4f¹⁴5d⁷6s²
Platinum	Pt^*	78		[Xe] 4f¹⁴5d¹⁰6s⁰
Gold	Au^*	79		[Xe] 4f¹⁴5d¹⁰6s^!
Mercury	Hg	80		[Xe] 4f¹⁴5d¹⁰6s

Fourth transition or 6d-series :

Element	Symbol	At. No.		Electronic configuration
Actinium	Ac	89	6d-orbitals are filled up	[Rn] 6d¹7s²
Rutherfordium	Rf	104		[Rn] 5f¹⁴6d²7s²
Hahnium	Ha	105		[Rn] 5f¹⁴6d³7s²
Seaborgium	Sg	106		[Rn] 5f¹⁴6d⁴7s²
Bohrium	Bh	107		[Rn] 5f¹⁴6d⁵7s²
Hassium	Hs	108		[Rn] 5f¹⁴6d⁶7s²
Meitnerium	Mt	109		[Rn] 5f¹⁴6d⁷7s²
Ununnilium	Uun	110		[Rn] 5f¹⁴6d⁸7s²
Unununium	Uuu	111		[Rn] 5f¹⁴6d⁹7s²
Unubium	Uub	112		[Rn] 5f¹⁴6d¹⁰7s²

Elements marked with asterisk have anomalous configurations. These are attributed to factors like nuclear-electron and electron-electron forces and stability of half filled and full filled orbital.

All transition elements are d block elements but all d block elements are not transition elements.

Atomic Radii of D Block Element

Atomic radii

The atomic, radii of 3d-series of elements are compared with those of the neighbouring s and p-block elements.

	Ca	Sc	Ti	V	Cr	Mn
227	197	144	132	122	117	117
Fe	Co	Ni	Cu	Zn	Ga	Ge
117	116	115	117	125	135	122^*

* in pm units

The atomic radii of transition elements show the following characteristics,

(i) The atomic radii and atomic volumes of d-block elements in any series decrease with increase in the atomic number. The decrease however, is not regular. The atomic radii tend to reach minimum near at the middle of the series, and increase slightly towards the end of the series.

Explanation: When we go in any transition series from left, to right, the nuclear charge increases gradually by one unit at each elements. The added electrons enter the same penultimate shell, (inner d- shell). These added electrons shield the outermost electrons from the attraction of the nuclear charge. The increased nuclear charge tends to reduce the atomic radii, while the added electrons tend to increase the atomic radii. At the beginning of the series, due to smaller number of electrons in the d-orbitals, the effect of increased nuclear charge predominates, and the atomic radii decrease. Later in the series, when the number of d-electrons increases, the increased shielding effect and the increased repulsion between the electrons tend to increase the atomic radii. Somewhere in the middle of the series, therefore the atomic radii tend to have a minimum value as observed.

(ii) The atomic radii increase while going down in each group. However, in the third transition series from hafnium (Hf) and onwards, the elements have atomic radii nearly equal to those of the second transition elements.

Explanation: The atomic radii increase while going down the group. This is due to the introduction of an additional shell at each new element down the group. Nearly equal radii of second and third transition series elements is due to a special effect called lanthanide contraction.

1. The atomic radii of the elements are almost same of which series

(a) Fe – Co – Ni (b) Na – K – Rb

Ans: c

2. A reduction in atomic size with increase in atomic number is a characteristic of elements of

(a) High atomic masses (b) d-block

Ans: a

Actinides

The elements with atomic numbers 90 to 103 i.e. thorium to lawrencium (which come immediately after actinium, Z = 89) are called actinides or actinones. These elements involve the filling of 5 f-orbitals. Their general electronic configuration is, [ Rn ]5 f^1-14 6d^0-1 7s².

They include three naturally occuring elements thorium, protactinium and uranium and eleven transuranium elements or transuranics which are produced artificially by nuclear reactions. They are synthetic or man made elements. All actinides are radioactive.

Properties of actinides

(i) Oxidation state: The dominant oxidation state of actinides is +3 which shows increasing stability for the heavier elements. Np shows +7 oxidation state but this is oxidising and is reduced to the most stable state +5. Pu also shows states upto +7 and Am upto +6 but the most stable state drops to Pu (+4) and Am (+3). Bk in +4 state is strongly oxidising but is more stable than Cm and Am in 4 state due to f ⁷ configuration. Similarly, No is markedly stable in +2 state due to its f¹⁴ configuration. When the oxidation number increases to + 6, the actinide ions are no longer simple. The high charge density causes the formation of oxygenated ions e.g., UO²⁺₂ , NpO²⁺₂ etc. The exhibition of large number of oxidation states of actinides is due to the fact that there is a very small energy gap between 5f, 6d and 7s subshells and thus all their electrons can take part in bond formation.

(ii) Actinide contraction: There is a regular decrease in ionic radii with increase in atomic number from Th to Lr. This is called actinide contraction analogous to the lanthanide contraction. It is caused due to imperfect shielding of one 5f electron by another in the same shell. This results in increase in the effective nuclear charge which causes contraction in size of the electron cloud.

(iii) Colour of the ions: Ions of actinides are generally coloured which is due to f – f transitions. It depends upon the number of electrons in 5 f orbitals.

(iv) Magnetic properties: Like lanthanides, actinide elements are strongly paramagnetic. The magnetic moments are lesser than the theoretically predicted values. This is due to the fact that 5f electrons of actinides are less effectively shielded which results in quenching of orbital contribution.

(v) Complex formation: Actinides have a greater tendency to form complexes because of higher nuclear charge and smaller size of their atoms. They form complexes even with p-bonding ligands such as alkyl phosphines, thioethers etc, besides EDTA, β-diketones, oxine etc. The degree of complex formation decreases in the order.

M⁴⁺> MO²⁺₂ > M³ > MO⁺₂

Where M is element of actinide series. There is a high concentration of charge on the metal atom inMO²⁺₂ which imparts to it relatively high tendency towards complex formation.

Oxidation states

Most of the transition elements exhibit several oxidation states i.e., they show variable valency in their compounds. Some common oxidation states of the first transition series elements are given below in table,

Outer Ele. Confi. and O. S. for 3d- elements

Elements

Outer electronic configuration

Oxidation states

Sc

3d¹ 4s²

+ 2, + 3

Ti

3d² 4s²

+ 2, + 3, + 4

V

3d³ 4s²

+ 2,+ 3,+ 4,+ 5

Cr

3d⁵ 4s¹

+ 1, + 2, + 3, + 4, + 5, + 6

Mn

3d⁵4s²

+ 2, + 3, + 4, + 5, + 6, + 7

Fe

3d⁶4s²

+ 2, + 3, + 4, + 5, + 6

Co

3d⁷4s²

+ 2, + 3, + 4

Ni

3d⁸4s²

+ 2, + 3, + 4

Cu

3d¹⁰4s¹

+ 1,+ 2

Zn

3d¹⁰4s²

+ 2

The relative stability of the different oxidation states depends upon the factors such as, electronic configuration, nature of bonding, stoichiometry, lattice energies and solvation energies. The highest oxidation states are found in fluorides and oxides because fluorine and oxygen are the most electronegative elements. The highest oxidation state shown by any transition metal is eight. The oxidation state of eight is shown by Ru and Os.

An examination of the common oxidation states reveals the following conclusions.

(i) The variable oxidation states shown by the transition elements are due to the participation of outer ns and inner (n–1)d-electrons in bonding.

(ii) Except scandium, the most common oxidation state shown by the elements of first transition series is +2. This oxidation state arises from the loss of two 4s electrons. This means that after scandium, d-orbitals become more stable than the s-orbital.

(iii) The highest oxidation states are observed in fluorides and oxides. The highest oxidation state shown by any transition elements (by Ru and Os) is 8.

(iv) The transition elements in the + 2 and + 3 oxidation states mostly form ionic bonds. In compounds of the higher oxidation states (compound formed with fluorine or oxygen), the bonds are essentially covalent. For example, in permanganate ion MnO₄^–, all bonds formed between manganese and oxygen are covalent.

(v) Within a group, the maximum oxidation state increases with atomic number. For example, iron shown the common oxidation state of + 2 and + 3, but ruthenium and osmium in the same group form compounds in the + 4, + 6 and + 8 oxidation states.

(vi) Transition metals also form compounds in low oxidation states such as +1 and 0. For example, nickle in, nickel tetracarbonyl, Ni(CO)₄ has zero oxidation state. Similarly Fe in Fe(CO)₅ has zero oxidation state.

The bonding in the compounds of transition metals in low oxidation states is not always very simple.

Example 1: The highest oxidation state of Cr will be

             (a) 2                                         (b) 3

             (c) 4                                         (d) 6

Ans: d

Catalytic Properties of D block Elements

Catalytic properties

Most of the transition metals and their compounds particularly oxides have good catalytic properties. Platinum, iron, vanadium pentoxide, nickel, etc., are important catalysts. Platinum is a general catalyst. Nickel powder is a good catalyst for hydrogenation of unsaturated organic compound such as, hydrogenation of oils some typical industrial catalysts are,

(i) Vanadium pentoxide (V₂O₅) is used in the Contact process for the manufacture of sulphuric acid,

(ii) Finely divided iron is used in the Haber’s process for the synthesis of ammonia.

Explanation : Most transition elements act as good catalyst because of,

(i) The presence of vacant d-orbitals.

(ii) The tendency to exhibit variable oxidation states.

(iii) The tendency to form reaction intermediates with reactants.

(iv) The presence of defects in their crystal lattices.

Alloy formation : Transition metals form alloys among themselves. The alloys of transition metals are hard and high metals are high melting as compared to the host metal. Various steels are alloys of iron with metals such as chromium, vanadium, molybdenum, tungsten, manganese etc.

Explanation : The atomic radii of the transition elements in any series are not much different from each other. As a result, they can very easily replace each other in the lattice and form solid solutions over an appreciable composition range. Such solid solutions are called alloys.

Chemical reactivity : The d-block elements (transition elements) have lesser tendency to react, i.e., these are less reactive as compared to s-block elements.

Explanation : Low reactivity of transition elements is due to,

(i) Their high ionisation energies.

(ii) Low heats of hydration of their ions.

(iii) Their high heats of sublimation.

Example 1: The catalytic activity of the transition metals and their compounds is ascribed to their

(a) Chemical reactivity

(b) Magnetic behaviour

(d) Ability to adopt multiple oxidation states and their complexing ability

Ans: d

Example 2: Which one of the following statement is true for transition elements

(a) They exhibit diamagnetism

(b) They exhibit inert pair effect

(d) They show variable oxidation states

Ans: d

Chemical Properties of Potassium Dichromate

Chemical properties Of Potassium Dichromate

(i) Action of heat: Potassium dichromate when heated strongly. Decomposes to give oxygen.

4K₂Cr₂O₇(s) ————→ 4K₂CrO₄(s) + 2Cr₂O₃(s) + 3O₂

(ii) Action of acids

(a) In cold, with concentrated H₂SO₄, red crystals of chromium trioxide separate out.

K₂Cr₂O₇(aq) + conc.H₂SO₄ → KHSO₄(aq) + 2CrO₃(s) + H₂O

On heating a dichromate-sulphuric acid mixture, oxygen gas is given out.

2K₂Cr₂O₇ + 8H₂SO₄ → 2K₂SO₄ + 2Cr₂(SO₄)₃ + 8H₂O + 3O₂

(b) With HCl, on heating chromic chloride is formed and Cl₂ is liberated.

K₂Cr₂O₇(aq) + 14HCl(aq) → 2CrCl₃(aq) + 2KCl(aq) + 7H₂O + 3Cl₂(g)

(iii) Action of alkalies: With alkalies, it gives chromates. For example, with KOH,

K₂Cr₂O₄ + 2KOH → 2K₂CrO₄ + H₂O

^{orange yellow}

On acidifying, the color again changes to orange-red owing to the formation of dichromate.

2K₂CrO₄ + H₂SO₄ → K₂Cr₂O₇ + K₂SO₄ + H₂O

Actually, in dichromate solution, the ions are in equilibrium with ions.

Cr₂O₇^2– + H₂O 2CrO₄^2– + 2H⁺

(iv) Oxidizing nature : In neutral or in acidic solution, potassium dichromate acts as an excellent oxidizing agent, and Cr₂O₇^2–gets reduced to Cr³⁺. The standard electrode potential for the reaction,

Cr₂O₇^2– + 14H⁺ + 6e^– → 2Cr⁺³ + 7H₂O is + 1.31V.

This indicates that dichromate ion is a fairly strong oxidizing agent, especially in strongly acidic solutions. That is why potassium dichromate is widely used as an oxidizing agent, for quantitative estimation of the reducing agents such as, Fe²⁺. It oxidizes,

(a) Ferrous salts to ferric salts

K₂CrO₇ + 4H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3[O]

2FeSO₄ + H₂SO₄ + [O] → Fe₂[SO₄]₃ + H₂O × 3

K₂Cr₂O₇ + 6FeSO₄ + 7H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 3Fe₂(SO₄)₃ + 7H₂O

Ionic equation

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

(b) Sulphites to sulphates and arsenites to arsenates.

K₂Cr₂O₇ + 4H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3[O]

Na₂SO₃ + [O] → Na₂SO₄] × 3

K₂Cr₂O₇ + 4H₂SO₄ + 3Na₂SO₃ → K₂SO₄ + Cr₂(SO₄)₃ + 3Na₂SO₄ + 4H₂O

Ionic equation

Cr₂O₇^2– + 8H⁺ + 3SO₃^2– → 2Cr³⁺ + 3SO₄^2– + 4H₂O

Similarly, arsenites are oxidised to arsenates.

Cr₂O₇^2– + 8H⁺ + 3AsO₃^3– → 2Cr³⁺ + 3AsO₄^3– + 4H₂O

K₂Cr₂O₇ + 4H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3[O]

2HX + O → H₂O + X₂] × 3

K₂Cr₂O₇ + 4H₂SO₄ + 6HX → K₂SO₄ + Cr₂(SO₄)₃ + 7H₂O + 3X₂

where, X may be Cl, Br, I.

Cr₂O₇^2– + 8H⁺ + 6HX → 2Cr³⁺ + 3X₂ + 7H₂O

(d) Iodides to iodine

K₂Cr₂O₇ + H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3[O]

2KI + H₂O + [O] → 2KOH + I₂] × 3

K₂Cr₂O₇ + 7H₂SO₄ + 6KI → 4K₂SO₄ + Cr₂(SO₄)₃ + 3I₂ + 7H₂O

Ionic equation :

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

Thus, when KI is added to an acidified solution of K₂Cr₂O₇ iodine gets liberated.

(e) It oxidizes H₂S to S.

K₂Cr₂O₇ + 4H₂SO₄ → K₂SO₄ + Cr₂(SO₄)₃ + 4H₂O + 3[O]

H₂S + [O] → H₂O + S] × 3

K₂Cr₂O₇ + 4H₂SO₄ + 3H₂S → K₂SO₄ + Cr₂(SO₄)₃ + 7H₂O + 3S

Ionic equation

Cr₂O₇^2– + 8H⁺ + H₂S → 2Cr³⁺ + 3S + 7H₂O

(v) Formation of insoluble chromates : With soluble salts of lead, barium etc., potassium dichromate gives insoluble chromates. Lead chromate is an important yellow pigment.

2Pb(NO₃)₂ + K₂Cr₂O₇ + H₂O → 2PbCrO₄ + 2KNO₃ + 2HNO₃

(vi) Chromyl chloride test : When potassium dichromate is heated with conc. H₂SO₄ in the presence of a soluble chloride salt, the orange-red vapors of chromyl chloride (CrO₂Cl₂) are formed.

_heat

K₂Cr₂O₇ + 4NaCl + 6H₂SO₄ ————→ 2KHSO₄ + 4NaHSO₃ + 2CrO₂Cl₂

^{chromyl chloride}

sup> (orange-red vapours)

Chromyl chloride vapors when passed through water give yellow-colored solution containing chromic acid.

CrO₂Cl₂ + 2H₂O → 2HCl + H₂CrO₄

^{Chromic acid}

^{(yellow solution)}

Chromyl chloride test can be used for the detection of chloride ion is any mixture.

Uses : Potassium dichromate is used as,

(i) An oxidizing agent

(ii) In chrome tanning

(iii) The raw material for preparing large number of chromium compounds

(iv) Primary standard in the volumetric analysis.

Structures of Chromate and Dichromate Ions

Chromates and dichromate are the salts of chromic acid (H₂CrO₄). In solution, these ions exist in equilibrium with each other. Chromate ion has four oxygen atoms arranged tetrahedrally around Cr atom. (see Fig). Dichromate ion involves a Cr–O–Cr bond as shown in Fig.

Chemical Properties of Potassium Permanganate

Chemical properties of Potassium Permanganate

Some important chemical reactions of KMnO₄ are given below,

Action of heat: KMnO₄ is stable at room temperature, but decomposes to give oxygen at higher temperatures.

_heat

2KMnO₄(s) ————→ K₂MnO₄(s) + MnO₂ + O₂(g)

Oxidizing actions : KMnO₄ is a powerful agent in neutral, acidic and alkaline media. The nature of reaction is different in each medium. The oxidising character of KMnO₄ (to be more specific, of MnO₄^–) is indicated by high positive reduction potentials for the following reactions.

Acidic medium:

MnO₄^– + 8H⁺ + 5e^– ? Mn²⁺ + 4H₂O E^o = 1.51 V

Alkaline medium

MnO₄^– + 2H₂O + 3e^– ? MnO₂ + 4OH^– E^o = 1.23 V

In strongly alkaline solutions and with excess of MnO₄^–, the reaction is

MnO₄^– + e^– ? MnO₄^2– E^o = 0.56 V

There are a large number of oxidation-reduction reactions involved in the chemistry of manganese compounds. Some typical reactions are,

In the presence of excess of reducing agent in acidic solutions permanganate ion gets reduced to manganous ion, e.g.,

5Fe²⁺ + MnO₄^– + 8H⁺ → 5Fe³⁺ ? Mn²⁺ + 4H₂O

An excess of reducing agent in alkaline solution reduces permanganate ion only to manganese dioxide e.g.,

3NO₂^– + MnO₄^– + 2OH^– → 3NO₃^– + MnO₂ + 4H₂O

In faintly acidic and neutral solutions, manganous ion is oxidized to manganese oxidized to manganese dioxide by permanganate.

2MnO₄^– + 3Mn⁺² + 2H₂O → 5MnO₂ + 4H⁺

In strongly basic solutions, permangante oxidizes manganese dioxide to manganate ion.

MnO₂ + 2MnO₄^– + 4OH^– → 3MnO₄^2– + 2H₂O

In acidic medium, KMnO₄ oxidizes,

Ferrous salts to ferric salts

2KMnO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

2FeSO₄ + H₂SO₄ + [O] → Fe2(SO₄)₃ + H₂O] × 5

2KMnO₄ + 8H₂SO₄ + 10FeSO₄ → K₂SO₄ + 2MnSO₄ + 5Fe₂(SO₄)₃ + 8H₂O

Ionic equation

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

The reaction forms the basis of volumetric estimation of Fe²⁺ in any solution by KMnO₄.

Oxalic acid to carbon dioxide

2KmNO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

(COOH)₂ + [O] → 2CO₂ + H₂O] × 3

2KMnO₄ + 3H₂SO₄ + 5(COOH)₂ → K₂SO₄ + 2MnSO₄ + 10CO₂ + 8H₂O

Ionic equation

2MnO₄^– + 6H⁺ + 5(COOH)₂ → 2Mn²⁺ + 10CO₂ + 8H₂O

Sulphites to sulphates

2KMnO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

Na₂SO₃ + [O] → Na₂SO₄] × 5

2KMnO₄ + 3H₂SO₄ + 5Na₂SO₃ → K₂SO₄ + 2MnSO₄ + 5Na₂SO₄ + 3H₂O

Ionic equation

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

Iodides to iodine in acidic medium

2KMnO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

2KI + H₂O + [O] → I₂ + 2KOH × 5

2KOH + H₂SO₄ → K₂SO₄ + 2H₂O] × 5

2KMnO₄ + 8H₂SO₄ + 10 KI → 6K₂SO₄ + 2MnSO₄ + 5I₂ + 8H₂O

Ionic equation

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

Hydrogen peroxide to oxygen

2KMnO₄ + 3H₂SO₄ → K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

H₂O₂ + [O] → H₂O + O₂ × 5

2KMnO₄ + 3H₂SO₄ + 5H₂O₂ → K₂SO₄ + 2MnSO₄ + 8H₂O + 5O₂

Manganous sulphate (MnSO₄) to manganese dioxide (MnO₂)

2KMnO₄ + H₂O → 2KOH + 2MnO₂ + 3[O]

MnSO₄ + H₂O + [O] → MnO₂ + H₂SO₄ × 3

2KOH + H₂SO₄ → K₂SO₄ + 2H₂O

2KMnO₄ + 3MnSO₄ + 2H₂O → 5MnO₂ + K₂SO₄ + 2H₂SO₄

Ionic equation

2MnO₄^– + 3Mn²⁺ + 2H₂O → 5MnO₂ + 4H⁺

Ammonia to nitrogen

2KMnO₄ + H₂O → 2MnO₂ + 2KOH + 3[O]

2NH₃ + 3[O] → N₂(g) + 3H₂O

2KMnO₄ + 2NH₃ → 2MnO₂ + 2KOH + 2H₂O + N₂(g)

Uses : KMnO₄ is used,

(i) As an oxidizing agent. (ii) As a disinfectant against disease-causing germs. (iii) For sterilizing wells of drinking water. (iv) In volumetric estimation of ferrous salts, oxalic acid etc. (v) Dilute alkaline KMnO₄ solution known as Baeyer’s reagent.

Structure of Permanganate Ion (MnO₄^–) : Mn in MnO₄^– is in +7 oxidation state. Mn⁷⁺ exhibits sp³hybridisation in this ion. The structure of MnO₄^– is, shown in fig.

Chromium containing Compounds

Chromium Containing Compounds

Potassium dichromate, (K₂Cr₂O₇)

Potassium dichromate is one of the most important compound of chromium, and also among dichromates. In this compound Cr is in the hexavalent (+6) state.

Preparation : It can be prepared by any of the following methods,

(i) From potassium chromate : Potassium dichromate can be obtained by adding a calculated amount of sulphuric acid to a saturated solution of potassium chromate.

2K₂CrO₄ + H₂SO₄ → K₂Cr₂O₇ + K₂SO₄ + H₂O

^{potassium chromate potassium dichromate}

^{(yellow) (orange)}

K₂Cr₂O₇ Crystals can be obtained by concentrating the solution and crystallisation.

(ii) Manufacture from chromite ore : K₂Cr₂O₇ is generally manufactured from chromite ore (FeCr₂O₄). The process involves the following steps.

(a) Preparation of sodium chromate : Finely powdered chromite ore is mixed with soda ash and quicklime. The mixture is then roasted in a reverberatory furnace in the presence of air. Yellow mass due to the formation of sodium chromate is obtained.

4FeCr₂O₄ + O₂ → 2FeO₃ + 4Cr₂O₃

4FeCr₂O₄ + 8Na₂CO₃ + 6O₂ → 8Na₂CrO₄ + 8CO₂(g)

4FeCr₂O₄ + 8Na₂CO₃ + 7O₂ → 2Fe₂O₃ + 8CO₂(g) + 8Na₂CrO₄

^{sodium chromate}

The yellow mass is extracted with water, and filtered. The filtrate contains sodium chromate.

The reaction may also be carried out by using NaOH instead of Na₂CO₃. The reaction in that case is,

4FeCr₂O₄ + 16NaOH + 7O₂ → 8Na₂CrO₄ + 2FeO₃ + 8H₂O

(b) Conversion of chromate into dichromate : Sodium chromate solution obtained in step (a) is treated with concentrated sulphuric acid when it is converted into sodium dichromate.

2Na₂CrO₄ + H₂SO₄ → Na₂Cr₂O₇ + Na₂SO₄ + H₂O

^{sodium chromate sodium dichromate}

On concentration, the less soluble sodium sulphate, Na₂SO₄.10H₂O crystallizes out. This is filtered hot and allowed to cool when sodium dichromate, Na₂Cr₂O₇.2H₂O, separates out on standing.

(c) Concentration of sodium dichromate to potassium dichromate : Hot concentrated solution of sodium dichromate is treated with a calculated amount of potassium chloride. When potassium dichromate being less soluble crystallizes out on cooling.

Na₂Cr₂O₇ + 2KCl → KrCr₂O₇ + 2NaCl

^{sod. dichromate pot. dichromate}

Physical properties

(i) Potassium dichromate forms orange-red coloured crystals.

(ii) It melts at 699 K.

(iii) It is very stable in air (near room temperature) and is generally, used as a primary standard in the volumetric analysis.

(iv) It is soluble in water though the solubility is limited.

Compounds of Copper

(1) Halides of copper : Copper (II) chloride, CuCl₂ is prepared by passing chlorine over heated copper. Concentrated aqueous solution of CuCl₂ is dark brown but changes first to green and then to blue on dilution.

On heating, it disproportionates to copper (I) chloride and chlorine

_Heat

2CuCl₂ ————→ 2CuCl + Cl₂

It is used as a catalyst in the Daecon’s process for the manufacture of chlorine.

Copper (I) chloride, CuCl is a white solid insoluble in water. It is obtained by boiling a solution of CuCl₂ with excess of copper turnings and conc. HCl.

_HCl

CuCl₂ + Cu ————→ 2CuCl

It dissolves in conc. HCl due to the formation of complex H[CuCl₂]

CuCl + HCl ————→ H [CuCl₂]

It is used as a catalyst alongwith NH₄Cl in the preparation of synthetic rubber.

(2) Cuprous oxide Cu₂O: It is a reddish brown powder insoluble in water but soluble in ammonia solution, where it forms diammine copper (I) ion. Cu⁺ + 2NH₃ —→ [Cu(NH₃)₂]⁺. It is used to impart red colour to glass in glass industry.

(3) Cupric oxide CuO: It is dark black, hygroscopic powder which is reduced to Cu by hydrogen, CO etc. It is used to impart light blue colour to glass. It is prepared by heating copper nitrate.

2Cu(NO₃)₂ ————→ 2CuO + 4NO₂ + O₂

(4) Copper sulphate CuSO₄.5H₂O (Blue vitriol) : It is prepared by action of dil H₂SO₄ on copper scrap in presence of air.

2Cu + 2H₂SO₄ + O₂ —→ CuSO₄ + 2H₂O

^(air)

(i) On heating this blue salt becomes white due to loss of water of crystallization.

_503K

CuSO₄.5H₂O ————→ CuSO₄ + 5H₂O

^{Blue White}

At about 1000 K, CuSO₄ decomposes to give CuO and SO₃.

_1000K

CuSO₄ ————→ CuO + SO₃

(ii) It gives a deep blue solution of tetrammine copper (II) sulphate with

Cu₂SO₄ + 4NH₄OH, ————→ [Cu(NH₃)₄]SO₄ + 4H₂O

^{Blue colour}

(iii) With KCN it first gives yellow precipitate of CuCN which decomposes of give Cu₂(CN)₂.Cu₂(CN)₂dissolves in excess of KCN to give K₃[Cu(CN)₄]

(iv) With KI it gives white ppt. of Cu₂I₂

4KI + 2CuSO₄ → 2K₂SO₄ + Cu₂I₂ + I₂

^{white ppt.}

(v) With K₄[Fe(CN)₆], CuSO₄ gives a reddish brown ppt. of Cu₂[Fe(CN)₆]

2CuSO₄ + K₄[Fe(CN)₆] → Cu₂[Fe(CN)₆] + 2K₂SO₄

^{Reddish brown ppt.}

Uses: For electroplating and electrorefining of copper. As a mordant in dyeing. For making Bordeaux mixture (11 parts lime as milk of lime + 16 parts copper sulphate in 1,000 parts of water). It is an excellent fungicide. For making green pigments containing copper carbonate and other compounds of copper. Like Verdigris which is Cu(CH₃COO)₂Cu(OH)₂ i.e. basic copper acetate and is used as a green pigment in paints. As a fungicide in starch paste for book binding work.

(5) Cupric sulphide CuS It is prepared as follows : Cu(NO₃)₂ + H₂S → CuS + 2HNO₃.

^{Black ppt.}

(6) Basic copper carbonates : Because of lower solubility of the hydroxide, the normal carbonate does not exist. Two basic copper carbonates occur in nature viz malachite CuCO₃.Cu(OH)₂ which has a fine green colour, and azurite, 2CuCO₃, Cu(OH)₂ which is deep blue in colour.

Malachite is prepared by heating a mixture of CuSO₄ solution and limestone in a sealed tube at 423 – 443 K

_{423-443 K}

2CuSO₄ + 2CaCO₃ + H₂O —————→ CuCO₃Cu(OH)₂ + 2CaSO₄ + CO₂

^Malachite

At lower temperature azurite is formed

_{< 423}

3CuSO₄ + 3CaCO₃ + H₂O —————→ 2CuCO₃Cu(OH)₂ + 3CaSO₄ + CO₂

^Azurite

On heating, both decompose to give black cupric oxide, water and CO₂.

They are used as green and blue painter’s pigments under the name ‘malachite green’ and azurite blue’.

Compounds of Iron

Compounds of iron

(1) Oxides of Iron : Iron forms three oxides FeO, Fe₂O₃ (Haematite), Fe₃O₄ (magnetite also called magnetic oxide or load stone).

(i) Ferrous oxide, FeO: It is a black powder, basic in nature and reacts with dilute acids to give ferrous salts.

FeO + H₂SO₄ → FeSO₄ + H₂O; It is used in glass industry to impart green colour to glass.

(ii) Ferric oxide Fe₂O₃: It is a reddish brown powder, not affected by air or water; amphoteric in nature and reacts both with acids and alkalis giving salts. It can be reduced to iron by heating with C or CO.

Fe₂O₃ + 3C → 2Fe + 3CO; Fe₂O₃ + 3CO → 2Fe + 3CO₂

It is used as red pigment to impart red colour to external walls and as a polishing powder by jewellers.

(iii) Ferroso ferricoxide Fe₃O₄(FeO.Fe₂O₃): It is more stable than FeO and Fe₂O₃ magnetic in nature and dissolves in acids giving a mixture of iron (II) and iron (III) salts.

Fe₃O₄ + 4H₂SO₄ (dil.) → FeSO₄ + Fe₂(SO₄)₃ + 4H₂O

(2) Ferrous sulphide FeS: It is prepared by heating iron filing with sulfur. With dilute H₂SO₄ it gives H₂S.FeS + H₂SO₄ (dil) → FeSO₄ + H₂S ↑

(3) Ferric chloride FeCl₃: (i) preparation: It is prepared by Fe(OH)₃ treating with HCl

Fe(OH)₃ + 3HCl → FeCl₃ + 3H₂O

The solution on evaporation give yellow crystals of FeCl₃.6H₂O

(ii) Properties: (a) Anhydrous FeCl₃ forms reddish-black deliquescent crystals.

(b) FeCl₃ is hygroscopic and dissolves in H₂O giving brown acidic solution due to formation of HCl

FeCl₃ + 3H₂O → Fe(OH)₃ + 3HCl

^(Brown)

2Fe³⁺ + Cu → 2Fe²⁺ + Cu²⁺ (aq)

(d) In vapour state FeCl₃ exists as a dimer, Fe₂Cl₆

(e) FeCl₃ is used as stypic to stop bleeding from a cut.

(4) Ferrous sulphate, FeSO₄, 7H₂O (Green vitriol): It is prepared as follow,

Fe + H₂SO₄ → FeSO₄ + H₂

(i) One pressure to moist air crystals become brownish due to oxidation by air.

4FeSO₄ + 2H₂O + O₂ → 4Fe(OH)SO₄

(ii) On heating, crystals become anhydrous and on strong heating it decomposes to Fe₂O₃, SO₂ andSO₃.

heat strong

FeSO₄.7H₂O ———→ FeSO₄ + 7H₂O 2FeSO₄ ————→ Fe₂O₃ + SO₂ + SO₃

heating

(iii) It can reduce acidic solution of KMnO₄ and K₂Cr₂O₇

(iv) It is generally used in double salt with ammonium sulphate.

(NH₄)₂SO₄ + FeSO₄ + 6H₂O → FeSO₄.(NH₄)₂SO₄.6H₂O

^{Mohr's salt}

Mohr’s salt is resistant to atmospheric oxidation.

(v) It is used in the ring test for nitrate ions where it gives brown coloured ring of compound FeSO₄.NO.

FeSO₄ + NO → FeSO₄ . NO

(vi) FeSO₄ is used in manufacture of blue black ink.

(vii) FeSO₄ + H₂O₂ is known as a name of Fenton’s reagent.

(5) Mohr's salt FeSO₄.(NH₄)₂SO₄.6H₂O: It is a double salt and is prepared by crystallising a solution containing equivalent amounts of FeSO₄.7H₂O and (NH₄)₂SO₄. It may be noted that Mohr’s salt contains only Fe²⁺ ions without any trace of Fe³⁺ ions. In contrast FeSO₄.7H₂O always contains some Fe³⁺ ions due to aerial oxidation of Fe²⁺ ions. Mohr salt is, therefore, used as a primary standard in volumetric analysis since a standard solution of Fe²⁺ ions can be obtained directly by weighing a known amount of the Mohr salt.

It acts as a reducing agent and as such reduces acidified KMnO₄ and K₂Cr₂O₇ solutions.

MnO₄^– + 5Fe²⁺ + 8H⁺ → 5Fe³⁺ + Mn²⁺ + 4H₂O

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

^{(From mohrs salt)}

Compounds Of Mercury

Compounds of Mercury

(1) Mercuric oxide, HgO : It is obtained as a red solid by heating mercury in air or oxygen

for a long time at 673 K

_679K

2Hg + O₂ → 2 HgO ( red )

or by heating mercuric nitrate alone or in the presence of Hg

_Heat

2Hg ( NO₃ )₂ → 2HgO + 4NO₂ + O₂

^red

When NaOH is added to a solution of HgCI₂, yellow precipitate of HgO are obtained.

Hg₂ CI₂ + 2NaOH → HgO ↓+ H₂O + 2NaCl

^(yello)

Red and yellow forms of HgO differ only in their particle size. On heating to 673 K, yellow form changes to red form.

_673K

HgO → HgO

^Yellow ^red

It is used in oil paints or as a mild antiseptic in ointments.

(2) Mercuric chloride, HgCl₂ : It is obtained by treating Hg with CI₂ or by heating a mixture of NaCland HgSO₄ in presence of small amount of MnO₂ (which oxidizes any Hg(I) salts formed during the reaction).

Heat

HgSO₄ + 2NaCI → HgCI₂ + Na₂SO₄

MnO₂

It is a white crystalline solid and is commonly known as corrosive sublimate. It is a covalent compound since it dissolves in organic solvents like ethanol and ether.

It is extremely poisonous and causes death. Its best antidote is white of an egg.

When treated with stannous chloride, it is first reduced to white ppt. of mercurous chloride and then to mercury (black).

2HgCI₂ + SnCI₂ → Hg₂CI₂+ SnCI₄

^{white ppt.}

Hg₂CI₂ + SnCI₂ → 2Hg + SnCI₄

^grey

With ammonia it gives a white ppt. known as infusible white ppt.

HgCI₂ + 2NH₃ → Hg ( NH₂ ) CI + NH₄ CI

A dilute solution of HgCI₂ is used as an antiseptic.

(3) Mercuric iodide, HgI₂ : It is obtained when a required amount of KI solution is added to a solution of . HgCI₂

HgCI₂ + 2KI → Hgl₂ + 2KCl

^(red)

Below 400 K, Hgl₂ is red but above 400 K, it turns yellow

Hgl₂ readily dissolves in excess of KI solution to form the ( Hgl₄ )^2- complex ion.

Hgl₂ + 2Kl → K₂Hgl₄

_{Red ppt. soluble colourless solution}

An alkaline solution of K₂ [ Hgl₄] is called Nessler’s reagent and is used to test NH⁺₄ ions.

It gives a brown ppt. of NH₂ - Hg - O Hg - I (Iodide of Millon’s base) with NH⁺₄ ions.

2K₂ [ HgI₄ ] + NH₃ + 3KOH → NH₂. HgO. Hgl + 7KI + 2H₂ O

It is used in ointments for treating skin infections.

(4) Mercurous chloride, Hg₂Cl₂ : It is obtained as under :

(a) Hg₂( NO₃) + 2NaCl → Hg₂Cl₂+ 2NaNo₃

^{white ppt.}

Heat in an iron retort

(b) HgCl₂ + Hg → Hg₂ Cl₂(condenses on cooling)

It is purified by sublimation.

Mercurous chloride is also called calomel. It is a white powder insoluble in H₂O. On heating, it decomposes to give HgCl₂ and Hg.

_Heat

_{Hg₂ Cl₂ → Hgcl₂ + Hg}

It dissolves in chlorine water forming mercuric chloride.

Hg₂Cl₂ + Cl₂ → 2HgCl₂

With ammonia, it turns black due to the formation of a mixture of finely divided black Hg and mercuric amino chloride.

Hg₂Cl₂ + 2NH₃ → Hg + NH₂ HgCl + NH₄Cl

(black)

It is used to prepare standard calomel electrode and as a purgative in medicine.

(5) Mercuric sulphide, HgS : The solubility product of HgS is lower than that of ZnS and hence it gets precipitated as black solid when H₂S is passed through an acidic solution of any mercury (II) salt.

HgCl₂ + H₂S → HgS + 2HCl

It is insoluble in water and HCl but dissolves in aqua regia (1 part conc.HNO₃ + 3 parts conc. HCl)

3HCl + HNO₃ → NOCl + 2H₂O + 2 [ Cl ]

^{Aqua regia Nitrosyl chloride Nacent chlorine}

HgS + 2| Cl | → HgCl₂ + S ↓

(Soluble)

On sublimation, its colour changes to red and hence it is used as a red pigment.

(6) Mercuric sulphate, HgSO₄ : It is obtained when HgS is treated with conc.H₂SO₄.

Hg + 2H₂ SO₄ → HgSo₄ + SO₂ + 2H₂O

It is a white solid which decomposes on heating to give mercurous sulphate.

3HgSo₄ → Hg₂ SO₄ + Hg + 2SO₂ + 2O₂

^675K

It is used as a catalyst in the hydration of alkynes to give aldehydes or ketones. It is also used as a cosmetic under the name Vermillon and in ayurvedic medicine as makardhwaj.

(7) Amalgams : Mercury forms alloys commonly known as amalgams, with all metals except iron and platinum. Hence it is transported in iron containers.

Compounds of Silver

Compounds of Silver:

(1) Silver oxide (Ag₂O): It is unstable and decomposes into Ag and O₂ on slow heating.

2Ag₂O → 4Ag + O₂

(2) Silver halides (AgF, AgCl, AgBr and Agl) : Only AgF is soluble in H₂O. AgCl is insoluble in H₂O but dissolves in NH₄OH, Na₂S₂O₃ and KCN solutions. AgBr is partly soluble whereas Agl is completely insoluble in NH₄OH. Except AgF, all the remaining three silver halides are photosensitive.

AgCl + 2NH₄OH → [Ag(NH₃)₂]Cl + 2H₂O

^{Diamine silver (I) chloride}

AgCl + 2KCN → K [Ag(CN)₂] + KCl

^{Pot.Dicyano argentate (I)}

AgCl + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaCl

^{Sod. Dithiosulphato argentate (I)}

(3) Silver nitrate (AgNO₃) : Silver nitrate (AgNO₃) is called lunar caustic silver nitrate on heating above its m.p. (485 K) decomposes to silver nitrite but on heating to red heat gives silver.

Above 485 K

2AgNO₃ ——————→ 2AgNO₂ + O₂

2AgNO₃ → 2AG + 2NO₂ + O₂

When treated with alkali, AgNO₃ forms silver oxide which in case of NH₄OH dissolves to form complex ion.

2AgNO₃ + 2NaOH → Ag₂O + 2NaNO₃ + H₂O

2AgNO₃ + 2NH₄OH → Ag₂O + 2NH₄NO₃ + H₂O

Ag₂O + 4NH₄OH → 2[Ag(NH₃)₂]OH + 3H₂O^{Diamine silver hydroxide}

AgNO₃ reacts with iodine in two ways

6AgNO₃ (excess) + 3l₂ + 3H₂O → AglO₃ + 5Agl + 6HNO₃

5AgNO₃ + 3l₂ (excess) + 3H₂O → HIO₃ + 5Agl + 5HNO₃

In contact with organic matter (skin, cloth, paper etc.) AgNO₃ is reduced to metallic silver (black)

2AgNO₃ + H₂O → 2Ag + 2HNO₃ + [O] → oxidises organic matter

AgNO₃ gives different coloured ppt. with different anions like Cl^–, Br^–, I^–, S^2–, S₂O₃^2–, CrO₄^3–, PO₄^3–etc.

AgNO₃ is used in the preparation of ink and hair dyes.

Photography : The photographic plate is coated with a colloidal gelatinised solution of AgBr. During exposure, AgBr is reduced to metallic silver.

2AgBr → 2Ag + Br₂

The exposed film is developed. The developer used is an alkaline solution of hydroquinone or quinol which reduces some of the exposed AgBr to black silver.

C₆H₄(OH)₂ + 2AgBr → 2Ag + C₆H₄O₂ + 2HBr

^{Quinol Quinone}

The film is finally fixed by dipping in a solution of sodium thiosulphate or hypo which removes unchanged AgBr as complex ion.

AgBr + 2Na₂S₂O₃ → Na₃[Ag(S₂O₃)₂] + NaBr

After taking a print of the photograph it is finally toned by dipping in a dilute solution of gold chloride to impart a beautiful golden colour or it is dipped in potassium chloro platinate K₂PtCl₆ solution to get a shining grey tinge.

AuCl₃ + 3Ag → 3AgCl + Au

Copper Ores

Copper Ores and its Extraction

(1) Ores: Copper pyrites (chalcopyrite) CeFeS₂ Cuprite (ruby copper) Cu₂O Copper glance (Cu₂S), Malachite [Cu(OH)₂.CuCO₃] Azurite [Cu(OH)₂.2CuCO₃]

(2) Extraction : Most of the copper (about 75%) is extracted from its sulphide ore, copper pyrites.

Concentration of ore : Froth floatation process.

Roasting: Main reaction:

2CuFeS₂ + O₂ → Cu₂S + 2FeS + SO₂.

Side reaction:

2Cu₂S + 3O₂ → 2Cu₂O + 2SO₂

2FeS + 3O₂ → 2FeO + 2SO₂.

Smelting:

FeO + SiO₂ → FeSiO₃ (slag)

Cu₂O + FeS → FeO + Cu₂S

The mixture of copper and iron sulphides melt together to form 'matte' (Cu₂S + FeS) and the slag floats on its surface.

Conversion of matte into Blister copper (Bessemerisation): Silica is added to matte and a hot blast of air is passed FeO + SiO₂ → FeSiO₃ (slag). Slag is removed. By this time most of iron sulphide is removed.

Cu₂S + 2Cu₂O → 6Cu + SO₂

Blister copper: Which contain about 98% pure copper and 2% impurities (Ag, Au, Ni, Zn etc.)

Properties of copper: It has reddish brown colour. It is highly malleable and ductile. It has high electrical conductivity and high thermal conductivity. Copper is second most useful metal (first being iron). It undergoes displacement reactions with lesser reactive metals e.g. with Ag. It can displace Ag from AgNO₃. The finally divided Ag so obtained is black in colour.

Copper shows oxidation states of +1 and +2. Whereas copper (I) salts are colourless, copper (II) salts are blue in colour. Cu (I) salts are less stable and hence are easily oxidised to Cu (II) salts (2Cu⁺ → Cu²⁺ + Cu) . This reaction is called disproportionation.

(1) In presence of atmospheric CO₂ and moisture, copper gets covered with a green layer of basic copper carbonate (green layer) which protects the rest of the metal from further acton.

Cu + O₂ + CO₂ + H₂O → Cu(OH)₂ CuCO₃

_{^{green layer}}

(2) In presence of oxygen or air, copper when heated to redness (below 1370K) first forms red cuprous oxide which changes to black cupric oxide on further heating. If the temperature is too high, cupric oxide changes back to cuprous oxide

_{Below 1370K O2}

4Cu + O₂ ————————→ 2Cu₂O ————————→ 4CuO

^{(Red) Above 1370K (Black)}

_{High temp}

CuO + Cu ————————→ Cu₂O

(3) Action of acids. Non oxidising dil. acids such as HCl, H₂SO₄ have no action on copper. However, copper dissolves in these acids in presence of air.

Cu + 2HCl + 1/2 O₂(air) → CuCl₂ + H₂O

With dil. HNO3, Cu liberates NO (nitric oxide)

3Cu + 8HNO₃ →, copper gives NO₂

Cu + 4HNO₃ → 3Cu(NO₃)₂ + 2NO₂ + 2H₂O

With hot conc. H₂SO₄, copper gives SO₂

Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O

Electrode potential of D block Elements

Electrode potentials(E^o)

Standard electrode potentials of some half–cells involving 3d-series of transition elements and their ions in aqueous solution are given in table,

Standard electrode potentials for 3d-elements

Elements	Ion	Electrode reaction	E°/ volt
Sc	Sc³⁺	Sc³⁺+ 3e^– Sc	– 2.10
Ti	Ti²⁺	Ti²⁺+ 2e^– Ti	– 1.60
V	V²⁺	V²⁺+ 2e^– V	– 1.20
Cr	Cr³⁺	Cr³⁺ + 3e^– Cr	– 0.71
Mn	Mn²⁺	Mn²⁺+ 2e^– Mn	– 1.18
Fe	Fe²⁺	Fe²⁺ + 2e^– Fe	– 0.44
Co	Co²⁺	Co²⁺ + 2e^– Co	– 0.28
Ni	Ni²⁺	Ni²⁺ + 2e^– Ni	– 0.24
Cu	Cu²⁺	Cu²⁺ + 2e^– Cu	+ 0.34
Zn	Zn²⁺	Zn²⁺ + 2e^– Zn	– 0.76

The negative values of E° for the first series of transition elements (except for Cu²⁺/ Cu ) indicate that,

(i) These metals should liberate hydrogen from dilute acids i.e., the reactions,

M + 2H⁺ → M²⁺ + H₂ (g); 2M + 6H⁺ → 2M³⁺ + 3H₂(g)

are favourable in the forward direction. In actual practice however, most of these metals react with dilute acids very slowly. Some of these metals get coated with a thin protective layer of oxide. Such an oxide layer prevents the metal to react further.

(ii) These metals should act as good reducing agents. There is no regular trend in the E° values. This is due to irregular variation in the ionisation and sublimation energies across the series.

Relative stabilities of transition metal ions in different oxidation states in aqueous medium can be predicted from the electrode potential data. To illustrate this, let us consider the following,

M(s) → M(g) ; ?H₁ = Enthalpy of sublimation, ?H_sub

M(g) → M⁺(g) + e^– ; ?H₂ = Ionisation energy, IE

M⁺(g) → M⁺(aq) ; ?H₃ = Enthalpy of hydration, ?H_hyd

Adding these equations one gets,

M(s) → M+(aq) + e–

?H = ?H₁ + ?H₂ + ?H₃ = ?H_sub + IE + ?H_hyd

The represents the enthalpy change required to bring the solid metal M to the monovalent ion in aqueous medium, M⁺(aq).

The reaction, M(s) → M⁺(aq) + e^–, will be favourable only if ?H is negative. More negative is the value is of ?H, more favourable will be the formation of that cation from the metal. Thus, the oxidation state for which ?H value is more negative will be stable in the solution.

Electrode potential for a Mⁿ⁺/M half-cell is a measure of the tendency for the reaction, Mⁿ⁺(aq) + ne^–→ M(s)

Thus, this reduction reaction will take place if the electrode potential for Mⁿ⁺/M half- cell is positive. The reverse reaction, M(s) → Mⁿ⁺(aq) + ne^–

Involving the formation of Mⁿ⁺(aq) will occur if the electrode potential is negative, i.e., the tendency for the formation of Mⁿ⁺(aq) from the metal M will be more if the corresponding E° value is more negative. In other words, the oxidation state for which E° value is more negative (or less positive) will be more stable in the solution.

When an elements exists in more than one oxidation states, the standard electrode potential (E°) values can be used in the predicting the relative stabilities of different oxidation states in aqueous solutions. The following rule is found useful.

The oxidation state of a cation for which ?H = (?H_sub + lE + ?H_hyd) or E° is more negative (for less positive) will be more stable.

1. Which of the following will have standard oxidation potential less than

(a) Zn (b) Cu

Ans: b\

Enthalpy of Atomization

Enthalpy of atomization

Transition metals exhibit high enthalpies of atomization.

Explanation : This is because the atoms in these elements are closely packed and held together by strong metallic bonds. The metallic bond is formed as a result of the interaction of electrons in the outermost shell. Greater the number of valence electrons, stronger is the metallic bond.

Solved example 1: Which one of the following characteristics of the transition metals is associated with their catalytic activity

(a) Variable oxidation states

(b) High enthalpy of atomization

(d) Colour of hydrated ions

Ans: a

Formation of colored ions

Most of the compound of the transition elements are colored in the solid state and /or in the solution phase. The compounds of transition metals are colored due to the presence of unpaired electrons in their d-orbitals.

Explanation : In an isolated atom or ion of a transition elements, all the five d-orbitals are of the same energy (they are said to be regenerate). Under the influence of the combining anion (s), or electron- rich molecules, the five d-orbitals split into two (or sometimes more than two) levels of different energies. The difference between the two energy levels depends upon the nature of the combining ions, but corresponds to the energy associated with the radiations in the visible region, (λ = 380 – 760nm). Typical splitting for octahedral and tetrahedral geometries are shown in fig.

Promotion of d-electrons to a higher level by

The splitting of d-orbital energy levels in (a) an octahedral, (b) a tetrahedral, geometry. This spllitting is termed as the crystal field splitting.

The transition metals in elements form or in the ionic form have one or more unpaired electrons. When visible light falls on the sample, the electrons from the lower energy level get promoted to a higher energy level due to the absorption of light of a characteristic wavelength (or color). This wavelength (or color) of the absorbed light depends upon the energy difference of the two levels. Rest of the light gets transmitted. The transmitted light has a color complementary to the absorbed color. Therefore, the compound or the solution appears to be of the complementary color. For example, Cu(H₂O)₆²⁺ ions absorb red radiation, and appear blue-green (blue-green is complementary color to red). Hydrated Co²⁺ ions absorb radiation in the blue-green region, and therefore, appear red in sunlight. Relationship between the color of the absorbed radiation and that of the transmitted light is given in table.

Relationship between the colors of the absorbed and transmitted light: the complementary colors.

Color of the		Color of the
Absorbed light	Transmitted light	Absorbed light	Transmitted light
IR	White	Blue-green	Red
Red	Blue-green	Blue	Orange
Orange	Blue	Indigo	Yellow
Yellow	Indigo	Violet	Yellow-green
Yellow-green	Violet	UV	White
Green	Purple

However, if radiations of all the wavelengths (or colors) except one are absorbed, then the color of the substance will be the color of the transmitted radiation. For example, if a substance absorbs all colors except green, then it would appear green to the eyes.

The transition metal ions which have completely filled d-orbitals are colorless, as there are no vacant d-orbitals to permit promotion of the electrons. Therefore, Zn²⁺ (3d¹⁰), Cd² + (4d¹⁰) and Hg²⁺(5d¹⁰) Sc ³⁺, Ti⁴⁺, Cu⁺ions and Zn, Cd, Hg are colorless and diamagnetic. The transition metal ions which have completely empty d-orbitals are also colorless, Thus, Sc³⁺ and Ti^4+.ions are colorless, unless a colored anion is present in the compound.

Colors and the outer- electronic configurations of the some important ions of the first transition series elements are given bellow,

Ion	Outer configuration	Number of unpaired electrons	Colour of the ion
Sc³⁺	3d⁰	0	Colourless
Ti³⁺	3d¹	1	Purple
Ti⁴⁺	3d⁰	0	Colourless
V³⁺	3d²	2	Green
Cr³⁺	3d³	3	Violet
Mn²⁺	3d⁵	5	Light pink
Mn³⁺	3d⁴	4	Violet
Fe²⁺	3d⁶	4	Green
Fe³⁺	3d⁵	5	Yellow
Co³⁺	3d⁷	3	Pink
Ni²⁺	3d⁸	2	Green
Cu²⁺	3d⁹	1	Blue
Cu⁺	3d¹⁰	0	Colourless
Zn²⁺	3d¹⁰	0	Colourless

Formation of complex Ions

Formation of complex ions

Transition metals and their ions show strong tendency for complex formation. The cations of transition elements (d-block elements) form complex ions with certain molecules containing one or more lone-pairs of electrons, viz., CO, NO, NH₃ etc., or with anions such as, F^–, Cl^–, CN^– etc. A few typical complex ions are,

[Fe(CN)]^4–, [Cu(NH₃)₄]²⁺, [Y(H₂O)₆]²⁺,

[Ni(CO)₄], [Co(NH₃)₆]³⁺ [FeF₆]^3–

Explanation : This complex formation tendency is due to,

(i) Small size and high nuclear charge of the transition metal cations.

(ii) The availability to vacant inner d-orbitals of suitable energy.

Formation of interstitial compounds: Transition elements form a few interstitial compounds with elements having small atomic radii, such as hydrogen, boron, carbon and nitrogen. The small atoms of these elements get entrapped in between the void spaces (called interstices) of the metal lattice. Some characteristics of the interstitial compound are,

(i) These are non-stoichiometric compounds and cannot be given definite formulae.

(ii) These compounds show essentially the same chemical properties as the parent metals, but differ in physical properties such as density and hardness. Steel and cast iron are hard due to the formation of interstitial compound with carbon. Some non-stoichimetric compounds are, VSe_0.98 (Vanadium selenide), Fe_0.94O and titanium nitride.

Explanation : Interstital compounds are hared and dense. This is because, the smaller atoms of lighter elements occupy the interstices in the lattice, leading to a more closely packed structure. Due to greater electronic interactions, the strength of the metallic bonds also increases.

Gold and its compounds

Gold and its Compounds

(1) Occurrence of gold : Gold is mainly found in native state either as vein gold, placer gold or alluvial gold. It is also present to a small extent in the combined state as sulphide, telluride and arsenosulphide. It is considered to be the king of metal.

Some important ores of gold are:

(i) Calaverite, AuTe₂ (ii) Sylvanite, AuAgTe₂ and

(iii) Bismuth aurite, BiAu₂

(2) Extraction of gold : (i) Mac-Arthur-Forest Cyanide process : The powdered gold ore, after concentration by Froth-floatation process, is roasted to remove easily oxidisable impurities of tellurium, arsenic and sulphur. The roasted ore is then treated with a dilute solution of KCN in presence of atmospheric oxygen when gold dissolves due to the formation of an aurocyanide complex.

4Au + 8KCN + 2H₂O + O₂ → 4K[Au(CN)₂] + 4KOH

^solution

The metal is then extracted by adding zinc dust.

2K[Au(CN)₂] + Zn → K₂[Zn(CN)₄] + 2Au↓

^ppt

(ii) Plattner’s chlorine process : The roasted ore is moistened with water and placed in wooden vats with false perforated bottoms. It is saturated with current of chlorine, gold chloride thus formed is leached with water and the solution is treated with a reducing agent such as FeSO₄ or H₂S to precipitate gold.

AuCl₃ + 3FeSO₄ → Au↓ + FeCl₃ + Fe₂(SO₄)₃

2AuCl₃ + 3H2S → 6HCl + 3S + 2Au↓

The impure gold thus obtained contains impurities of Ag and Cu. The removal of Ag and Cu from gold is called parting. This is done by heating impure gold with conc. H₂SO₄ (or HNO₃ when Ag and Cu dissolve leaving behind Au.

Cu + 2H₂SO₄ → CuSO₄ + SO₂ + 2H₂O

2Ag + 2H₂SO₄ → Ag₂SO₄ + SO₂ + 2H₂O

Properties of Gold: Gold is a yellow, soft and heavy metal. Gold and Ag are called noble metals since they are not attacked by atmospheric oxygen. However, Ag gets tarnished when exposed to air containing traces of H₂S. Gold is malleable, ductile and a good conductor of heat and electricity.

Pure gold is soft. It is alloyed with Ag or Cu for making jewellery. Purity of gold is expressed in terms of carats. Pure gold is 24 carats. Gold ’14 carats’ means that it is an alloy of gold which contains 14 parts by weight of pure gold and 10 parts of copper per 24 parts by weight of the alloy. Thus the percentage of gold in ’14 carats” of gold is = 100/24 × 14 = 58.3%.

Most of the jewellery is made from 22 carat gold (91.66% pure gold). Gold is quite inert. It does not react with oxygen, water and acids but dissolves in aqua regia

3HCl + HNO₃ → NOCl + 2H₂O + 2Cl] × 3

Au + 3Cl → AuCl₃] × 2

2Au + 9HCl + 3HNO₃ → 2AuCl₃ + 6H₂O + 3NOCl

^{Auric chloride Nitrosyl chloride}

Oxidation states of gold: The principal oxidation states of gold are + 1 and + 3 though + 1 state is more stable than + 3.

Compounds of gold

(1) Auric chloride, AuCl₃ : It is prepared by passing dry Cl₂ over finely divided gold powder at 573 K

_573K

2Au + 3Cl₂ ————→ 2AuCl₃

It is a red coloured crystalline solid soluble in water and decomposes on heating to give gold (I) chloride and Cl₂

_Heat

AuCl₃ ————→ AuCl + Cl₂

It dissolves in conc. HCl forming chloroauric acid

AuCl₃ + HCl → H[AuCl₄]

Chloroauric acid is used in photography for toning silver prints and as an antidote for snake poisoning.

(2) Aurous sulphide, Au₂S : It is prepared when H₂S is passed through an acidified solution of potassium aurocyanide, K[Au(CN)₂]

2K[Au(CN)₂] + H₂S → Au₂S + 2KCN + 2HCN

It is a dark brown solid, not attached by dilute mineral acids and hence is probably the most stable gold compound.

Hydroboration of Alkenes

HYDROBORATION OF ALKENES: PREPARATION OF ALKANES

(a) On treatment with acetic acid

B₂H₆ CH₃COOH

R – CH = CH₂ —————→ (R – CH₂ – CH₂)₃ B ——————→ R – CH₂ – CH₃

^{Alkene Trialkyl borane Alkane}

(b) Coupling of alkyl boranes by means of silver nitrate

2B₂H₆ AgNO₃25^oC

6[R – CH = CH₂] —————→ [2R – CH₂ – CH₂ – ]₃ B ——————→ 3[RCH₂CH₂ – CH₂CH₂R]

Example 1: A gas decolourised KMnO₄ solution but gives no precipitate with ammoniacal cuprous chloride is or Which of the following gases does not give a precipitate with ammoniacal solution of silver nitrate but decolourizes KMnO₄ (neutral or slightly alkaline)

(a) Ethane (b) Methane

Ans (c)

Example 2: A hydrocarbon containing 2 carbon atoms gives Sabatier and Senderen's reaction but does not give precipitate with ammoniacal silver nitrate solution. The hydrocarbon in the question is

(a) Ethane (b) Acetylene

Ans (c)

Ionic Radius of D Block Elements

Ionic Radius

For ions having identical charges, the ionic radii decrease slowly with the increase in the atomic number across a given series of the transition elements.

Elements (m):

Ionic radius,(M²⁺)/pm

	Pm:(M³⁺)/pm:
Sc	–	81
Ti	90	76
V	88	74
Cr	84	69
Mn	80	66
Fe	76	64
Co	74	63
Ni	72	–
Cu	69	–
Zn	74	–

Explanation :The gradual decrease in the values of ionic radius across the series of transition elements is due to the increase in the effective nuclear charge.

Solved example 1: Which of the following trivalent ion has the largest atomic radii in the lanthanide series

(a) La (b) Ce

Ans: a

Solved example 2:The atomic radii of the elements are almost same of which series

(a) Fe – Co – Ni (b) Na – K – Rb

Ans: a

Ionization energy of D block Elements

Ionization Energy

The ionization energies of the elements of first transition series are given below:

Elements	I₁	I₂	I₃
Sc	632	1245	2450
Ti	659	1320	2721
V	650	1376	2873
Cr	652	1635	2994
Mn	716	1513	3258
Fe	762	1563	2963
Co	758	1647	3237
Ni	736	1756	3400
Cu	744	1961	3560
Zn	906	1736	3838

* in kJ mol^–1

The following generalizations can be obtained from the ionisation energy values given above.

(i) The ionisation energies of these elements are high and in the most cases lie between those of s- and p-block elements. This indicates that the transition elements are less electropositive than s-block elements.

Explanation : Transition metals have smaller atomic radii and higher nuclear charge as compared to the alkali metals. Both these factors tend to increase the ionisation energy, as observed.

(ii) The ionisation energy in any transition series increases in the nuclear with atomic number; the increase however is not smooth and as sharp as seen in the case of s and p-block elements.

Explanation : The ionisation energy increases due to the increase in the nuclear charge with atomic number at the beginning of the series. Gradually, the shielding effect of the added electrons also increases. This shielding effect tends to decrease the attraction due to the nuclear charge. These two opposing factors lead to a rather gradual increase in the ionisation energies in any transition series.

(iii) The first ionisation energies of 5d-series of elements are much higher than those of the 3d and 4d series elements.

Explanation : In the 5d-series of transitions elements, after lanthanum (La), the added electrons go to the next inner 4f orbitals. The 4f electrons have poor shielding effect. As a result, the outermost electrons experience greater nuclear attraction. This leads to higher ionisation energies for the 5d- series of transition elements.

Solved example 1: Ionisation potential values of d-block elements as compared to ionization potential value of f-block elements are

(a) Higher (b) Equal

Ans: a

Iron and its properties

(1) Ores of iron : Haematite Fe₂O₃, Magnetite (Fe₃O₄) Limonite (Fe₂O₃.3H₂O), Iron pyrites (FeS₂) Copper pyrities (CuFeS₂) etc.

(2) Extraction : Cast iron isextracted from its oxides by reduction with carbon and carbon monoxide in a blast furnace to give pig iron.

Roasting : Ferrous oxide convert into ferric oxide.

Fe₂O₃.3H₂O → Fe₂O₃ + 3H₂O; 2FeCO₃ → 2FeO + 2CO₂

4FeO + O₂ → 2Fe₂O₃

Smelting : Reduction of roasted ore of ferric oxide carried out in a blast furnace.

(i) The reduction of ferric oxide is done by carbon and carbon monoxide (between 1473k to 1873k) 2C + O₂ → 2CO

_673k

(ii) Fe₂O₃ + 3CO ? 2Fe + 3CO₂. It is a reversible and exothermic reaction. Hence according toLe-chatelier principle more iron will be produced in the furnace at lower temp. Fe₂O₃ + CO → 2FeO + CO₂

_1073K

(iii) FeO + C ————————→ Fe + Co

^{endothermic reaction}

The gases leaving at the top of the furnace contain up to 28% CO and are burnt in cowper's stove to pre-heat the air for blast

Varieties of iron : The three commercial varieties of iron differ in their carbon contents. These are;

(1) Cast iron or Pig-iron : It is most impure form of iron and contains highest proportion of carbon (2.5–4%).

(2) Wrought iron or Malleable iron : It is the purest form of iron and contains minimum amount of carbon (0.12–0.25%).

(3) Steel : It is the most important form of iron and finds extensive applications. Its carbons content (Impurity) is mid-way between cast iron and wrought iron. It contains 0.2–1.5% carbon. Steels containing 0.2–0.5% of carbon are known as mild steels, while those containing 0.5–1.5% carbon are known as hard steels.

Steel is generally manufactured from cast iron by three processes, viz, (i) Bessemer Process which involves the use of a large pear-shaped furnace (vessel) called Bessemer converter, (ii) L.D. process and (iii) open hearth process, Spiegeleisen (an alloy of Fe, Mn and C) is added during manufacture of steel.

Heat treatment of steels : Heat treatment of steel may be defined as the process of carefully heating the steel to high temperature followed by cooling to the room temperature under controlled conditions. Heat treatment of steel is done for the following two purposes,

(i) To develop certain special properties like hardness, strength, ductility etc. without changing the chemical composition.

(ii) To remove some undesirable properties or gases like entrapped gases, internal stresses and strains. The various methods of heat treatment are,

(a) Annealing : It is a process of heating steel to redness followed by slow cooling.

(b) Quenching or hardening : It is a process of heating steel to redness followed by sudden cooling by plunging the red hot steel into water or oil.

(c) Tempering : It is a process of heating the hardened or quenched steel to a temperature much below redness (473–623K) followed by slow cooling.

(d) Case-hardening : It is a process of giving a thin coating of hardened steel to wrought iron or to a strong and flexible mild steel by heating it in contact with charcoal followed by quenching in oil.

(e) Nitriding : It is a process of heating steels at about 700^oC in an atmosphere of ammonia. This process imparts a hard coating of iron nitride on the surface of steel.

Properties of steel : The properties of steel depend upon its carbon contents. With the increase in carbon content, the hardness of steel increases while its ductility decreases.

(i) Low carbon or soft steels contain carbon upto 0.25%.

(ii) Medium carbon steels or mild steels contain 0.25–0.5% carbon.

(iii) High carbon or hard steels contains 0.1 – 1.5 percent carbon.

(iv) Alloy steels or special steels are alloys of steel with Ni, Cr, Co, W, Mn, V etc., For example

(a) Stainless steel (Fe = 73%, Cr = 18%, Ni = 8% + C) is resistant to corrosion and is used for making ornamental pieces, cutlery etc.

(b) Invar (Fe = 64%, Ni = 36%) has small coefficient of expansion and is used for making metre scales, pendulum rods and watches.

(c) Manganese steel (Fe = 86%, Mn 13% + carbon) is very hard and resistant to wear and hence is used for making rock drills, safes etc.

(d) Tungsten steel (Fe = 94%, W = 5% + carbon) is quite hard and is used for making high speed cutting tools.

(e) Permalloy (Fe = 21%, Ni = 78% + carbon) is strongly magnetised by electric current but loses magnetism when current is cut off. It is used for making electromagnets, ocean cables etc.

Properties of iron

(1) Dry or moist air has no action on pure iron but impure iron when exposed to moist air is covered with a layer of rust Fe₂O₃ + Fe(OH)₃. However, finely divided pure iron burns in air or oxygen forming Fe₃O₄ (magnetic oxide of iron).

3Fe + 2O₂ → Fe₃O₄

(2) Iron decomposes steam at red heat

_{Red heat}

3Fe + 4H₂O ——————→ Fe₃O₄ + 4H₂

^Steam

(3) Action of acids: Iron reacts with dil. HCl and dil. H₂SO₄ liberating hydrogen. with hot conc.H₂SO₄, it gives SO₂, with dil.HNO₃, it gives NH₄NO₃ and moderately conc. HNO₃ reacts with iron forming NO₂.

Cold conc. HNO₃ makes iron passive due to the deposit of a thin layer of iron oxide (Fe₃O₄) on the surface.

Hot conc. HNO₃ reacts with iron liberating NO.

Fe + 4HNO₃ (hot conc.) → Fe(NO₃)₃ + NO + 2H₂O

(4) Iron does not react with alkalies.

(5) It displaces less electropositive metals (e.g., Cu, Ag etc.) from their salts

CuSO₄ + Fe → FeSO₄ + Cu

(6) Finely divided iron combines with CO forming penta carbonyl

Fe + 5CO → Fe(CO)₅

(7) Iron does not form amalgam with Hg.

(8) Iron is the most abundant and most widely used transition metal.

Lanthanides

F Block Elements & Lanthanides

Lanthanides and actinides are collectively called f-block elements because last electron in them enters into f-orbitals of the antepenultimate (i.e., inner to penultimate) shell partly but incompletely filled in their elementary or ionic states. The name inner transition, elements is also given to them because they constitute transition series with in transition series (d-block elements) and the last electron enters into antepenultimate shell (n-2)f. In addition to incomplete d-subshell, their f-subshell is also incomplete. Thus, these elements have three incomplete outer shells i.e., (n–2), (n–1) and n shells and the general electronic configuration of f-block elements is (n–2) f^1-14 ( n- 1 ) d^0-10 ns².

(1) Lanthanides : The elements with atomic numbers 58 to 71 i.e. cerium to lutetium (which come immediately after lanthanum Z = 57) are called lanthanides or lanthanones or rare earths. These elements involve the filling of 4 f-orbitals. Their general electronic configuration is [ Xe ] 4f^1-14 4d^0-10 6s²,. Promethium (Pm), atomic number 61 is the only synthetic (man made) radioactive lanthanide.

Properties of lanthanides

(i) These are highly dense metals and possess high melting points.

(ii) They form alloys easily with other metals especially iron. e.g. misch metal consists of a rare earth element (94–95%), iron (upto 5%) and traces of S, C, Ca and Al, pyrophoric alloys contain Ce (40–5%),La + neodymium (44%), Fe (4–5%), Al (0–5%) and the rest is Ca, Si and C. It is used in the preparation of ignition devices e.g., trace bullets and shells and flints for lighters and cigarette.

(iii) Oxidation state : Most stable oxidation state of lanthanides is +3. Oxidation states + 2 and + 4 also exist but they revert to +3 e.g. Sm²+, Eu²⁺, Yb²⁺ lose electron to become +3 and hence are good reducing agents, where as Ce⁴⁺, Pr⁴⁺, Tb⁴⁺ in aqueous solution gain electron to become + 3 and hence are good oxidizing agents. There is a large gap in energy of 4 f and 5 d subshells and thus the number of oxidation states is limited.

(iv) Colour : Most of the trivalent lanthanide ions are coloured both in the solid state and in aqueous solution. This is due to the partly filled f-orbitals which permit f–f transition. The elements with xfelectrons have a similar colour to those of (14 – x) electrons.

(v) Magnetic properties : All lanthanide ions with the exception of Lu³⁺, Yb³⁺ and Ce ⁴⁺ are paramagnetic because they contain unpaired electrons in the 4 f orbitals. These elements differ from the transition elements in that their magnetic moments do not obey the simple “spin only” formula μ_eff= √n(n+2) B.M. where n is equal to the number of unpaired electrons. In transition elements, the orbital contribution of the electron towards magnetic moment is usually quenched by interaction with electric fields of the environment but in case of lanthanides the 4f-orbitals lie too deep in the atom for such quenching to occur. Therefore, magnetic moments of lanthanides are calculated by taking into consideration spin as well as orbital contributions and a more complex formula

μ_eff= √4S ( S+1 ) + L (L+1 ) B.M.

Which involves the orbital quantum number L and spin quantum number S.

(vi) Complex formation : Although the lanthanide ions have a high charge (+3) yet the size of their ions is very large yielding small charge to size ratio i.e., low charge density. As a consequence, they have poor tendency to form complexes. They form complexes mainly with strong chelating agents such as EDTA, β -diketones, oxine etc. No complexes with π - bonding ligands are known.

(vii) Lanthanide contraction : The regular decrease in the size of lanthanide ions from to is known as lanthanide contraction. It is due to greater effect of the increased nuclear charge than that of the screening effect.

Consequences of lanthanide contraction

(a) It results in slight variation in their chemical properties which helps in their separation by ion exchange

(b) Each element beyond lanthanum has same atomic radius as that of the element lying above it in the group (e.g. Zr 145 pm, Hf 144 pm); Nb 134 pm, Ta 134 pm ; Mo 129 pm, W 130 pm).

(c) The covalent character of hydroxides of lanthanides increases as the size decreases from La³⁺ to Lu³⁺ . However basic strength decreases. La (OH)₃Thus is most basic whereas Lu (OH)₃ is least basic. Similarly, the basicity of oxides also decreases in the order from La³⁺ to Lu³⁺.

(d) Tendency to form stable complexes from La³⁺ to Lu³⁺ increases as the size decreases in that order.

(e) There is a slight increase in electronegativity of the trivalent ions from La to Lu.

(f) Since the radius of Yb³⁺ ion (86 pm) is comparable to the heavier lanthanides Tb, Dy, Ho and Er, therefore they occur together in natural minerals.

Magnetic properties of D block Elements

Magnetic properties

Most of the transition elements and their compounds show paramagnetism. The paramagnetism first increases in any transition element series, and then decreases. The maximum paramagnetism is seen around the middle of the series. The paramagnetism is described in Bohr Magneton (BM) units. The paramagnetic moments of some common ions of first transition series are given below in Table

Explanation :A substance which is attracted by magnetic filed is called paramagnetic substance. The substances which are repelled by magnetic filed are, called diamagnetic substances. Paramagnetism is due to the presence of unpaired electrons in atoms, ions or molecules.

The magnetic moment of any transition element or its compound/ion is given by (assuming no contribution from the orbital magnetic moment).

μ_s = √4S(S+1) BM = √n(n+2) BM

where, S is the total spin (n×s) n is the number of unpaired electrons and s is equal to ½ (representing the spin of an unpaired electron).

From the equation given above, the magnetic moment (μ_s) increases with an increase in the number of unpaired electrons.

Magnetic moments of some ions of the 3d-series elements

Ion	Outer configuration	No. of unpaired electrons	Magnetic moment (BM)
			Calculated	observed
Sc³⁺	3d⁰	0	0	0
Ti³⁺	3d¹	1	1.73	1.75
Ti²⁺	3d²	2	2.84	2.86
V²⁺	3d³	3	3.87	3.86
Cr²⁺	3d⁴	4	4.90	4.80
Mn²⁺	3d⁵	5	5.92	5.95
Fe²⁺	3d⁶	4	4.90	5.0-5.5
Co²⁺	3d⁷	3	3.87	4.4-5.2
Ni²⁺	3d⁸	2	2.84	2.9-3.4
Cu²⁺	3d⁹	1	1.73	1.4-2.2
Zn²⁺	3d¹⁰	0	0	0

In d-obitals belonging to a particular energy level, there can be at the maximum five unpaired electrons in d⁵ cases. Therefore, paramagnetism in any transition series first increases, reaches a maximum value for d⁵ cases and then decreases thereafter.

Example 1: Which ion has maximum magnetic moment

(a) V⁺³ (b) Mn⁺³

Ans: c

Manganese Containing Compounds

Manganese containing compounds

Potassium Permanganate, (KMnO₄)

Potassium permanganate is a salt of an unstable acid HMnO₄ (permanganic acid). The Mn is an +7 state in this compound.

Preparation : Potassium permanganate is obtained from pyrolusite as follows.

Conversion of pyrolusite to potassium manganate : When manganese dioxide is fused with potassium hydroxide in the presence of air or an oxidising agent such as potassium nitrate or chlorate, potassium manganate is formed, possibly via potassium manganite.

_fused

MnO₂ + 2KOH ––––––→ K₂MnO₃ + 4H₂O] × 2

^{potassium manganite}

2K₂MnO₃ + O₂ ––––––→ 2K₂MnO₄ + 2H₂O

_fused

2MnO₂ + 4KOH + O₂ ––––––→ 2K₂MnO₄ + 2H₂O

^pyrolusite ^{potassium manganate}

^{[dark-green mass]}

Oxidation of potassium manganate to potassium permanganate: The potassium manganate so obtained is oxidized to potassium permanganate by either of the following methods.

By chemical method : The fused dark-green mass is extracted with a small quantity of water. The filtrate is warmed and treated with a current of ozone, chlorine or carbon dioxide. Potassium manganate gets oxidized to potassium permanganate and the hydrated manganese dioxide precipitates out. The reactions taking place are,

When CO₂ is passed

3K₂MnO₄ + 2H₂O → 2KMnO₄ + MnO₂ ↓ + 4KOH

^{potassium manganate potassium permanganate}

2CO₂ + 4KOH → 2K₂CO₃ + 2H₂O

When chlorine or ozone is passed

2K₂MnO₄ + Cl₂ → 2KMnO₄ + 2KCl

2K₂MnO₄ + O₃ + H₂O → 2KMnO₄ + 2KOH + O₂(g)

The purple solution so obtained is concentrated and dark purple, needle-like crystals having metallic lustre are obtained.

Electrolytic method : Presently, potassium manganate (K₂MnO₄) is oxidized electrolytically. The electrode reactions are,

At anode: 2MnO₄^2– → 2MnO₄^– + 2e^–

At cathode: 2H⁺ + 2e^– → H₂(g)

The purple solution containing KMnO₄ is evaporated under controlled condition to get crystalline sample of potassium permanganate.

Physical properties

KMnO₄ crystallizes as dark purple crystals with greenish luster (m.p. 523 K).

It is soluble in water to an extent of 6.5g per 100g at room temperature. The aqueous solution of KMnO₄has a purple colour.

Melting and Boiling Points

MELTING AND BOILING POINTS OF D BLOCK ELEMENTS

The melting and boiling points of transition elements except Cd and Hg, are very high as compared to the s-block and p-block elements. The melting and boiling points first increase, pass through maxima and then steadily decrease across any transition series. The maximum occurs around middle of the series.

Explanation : Atoms of the transition elements are closely packed and held together by strong metallic bonds which have appreciable covalent character. This leads to high melting and boiling points of the transition elements.

The strength of the metallic bonds depends upon the number of unpaired electrons in the outermost shell of the atom. Thus, greater is the number of unpaired electrons stronger is the metallic bonding. In any transition element series, the number of unpaired electrons first increases from 1 to 5 and then decreases back to the zero .The maximum five unpaired electrons occur at Cr (3d series). As a result, the melting and boiling points first increase and then decrease showing maxima around the middle of the series.

The low melting points of Zn, Cd, and Hg may be due to the absence of unpaired d-electrons in their atoms.

Solved example 1: Mercury is the only metal which is liquid at 0^oC. This is due to its

(a) Very high ionisation energy and weak metallic bond

(b) Low ionisation potential

(d) High vapour pressure

Ans: a

Mercury Ores

Mercury Ores and Extraction

(1) Occurrence and extraction of mercury: Cinnabar (HgS) is the only important ore of Hg. It is concentrated by froth floatation method and mercury is extracted from this ore by heating it in air at 773-873 K (auto reduction).

_273-873K

HgS + O₂ → Hg + SO₂

The mercury vapours thus obtained are condensed to give liquid metal. Hg thus obtained contains impurities of Zn, Sn and Pb. These are removed by treating the impure metal with dil HNO₃ , mercurous nitrate, Hg₂ ( NO₃ )₂ thus formed react with metals present as impurities forming their nitrates which pass into solution leaving behind pure mercury. However, it is best purified by distillation under reduced pressure.

_warm

6Hg + 8HNO₃ (dil. ) → 3 Hg ( NO₃ )₂ + 4H₂O + 2NO

Zn + Hg₂ ( NO₃ )₂ → Zn ( NO₃ )₂ + 2Hg

Similar reaction is given by Pb and Sn.

Properties of mercury : Mercury is less reactive than Zn. It is a liquid at room temperature and has low thermal and electrical conductivity. Mercury forms dimeric mercury (I) ions, Hg⁺²₂ in which the two atoms are bonded by a covalent bond. It is slowly oxidised to HgO at about its boiling point. Hg does not react with dil. HCl or dil. H₂SO₄ but reacts with hot concentrated H₂SO₄ to form HgSO₄, it reacts with both warm dil. and conc. HNO₃ evolving NO and NO₂ respectively.

Hg + 2H₂SO₄ ( hot, conc. ) → HgSO₄ + SO₂ + 2H₂O

Hg + 4HNO₃ ( conc. ) → Hg ( NO₃)₂ +2NO₂ + 2H₂O

Hg does not react with steam or water hence can’t form any hydroxide.

Metallic character of D block Elements

METALLIC CHARACTER OF D BLOCK ELEMENTS

All the transition elements are metals. These are hard, and good conductor of heat and electricity. All these metals are malleable, ductile and form alloys with other metals. These elements occur in three types e.g., face- centered cubic (fcc), hexagonal close-packed (hcp) and body-centered cubic (bcc), structures.

The transition elements shows both covalent as well as metallic bonding amongst their atoms.

Explanation : The ionisation energies of the transition elements are not very high. The outermost shell in their atoms have many vacant, partially filled orbitals. These characteristics make these elements metallic in character. The hardness of these metals, suggests the presence of covalent bonding in these metals. The presence of unfilled d-orbitals favour covalent bonding. Metallic bonding in these metals is indicated by the conducting nature of these metals. Therefore, it appears that there exists covalent and metallic bonding in transition elements.

Solved Example 1: In human body if necessary, the plate, screw or wire used for surgery are made up of

(a) Ni (b) Au

Ans: d

Solved Example 2: A hard and resistant metal (alloy) generally used in tip of nib of fountain pen is

(a) Os.lr (b) Pt.Cr

Ans : b

Silver Ores and Extraction

Silver ores and its extraction

(1) Ores : Argentite (silver glance) Ag₂S Horn silver (AgCl) Ruby silver (Pyrargyrite) 3Ag₂S.Sb₂S₃.

(2) Extraction : Cyanide process or Mac Arthus-Forrest cyanide process : This method depends on the fact that silver, its sulphide or chloride, forms soluble complex with alkali cyanides in the silver. This implies that silver compounds will dissolve in solution of alkali cyanides in the presence of blast of air.

4Ag + 8NaCN + 2H₂O + O₂ ? 4Na[Ag(CN)₂] + 4NaOH

air

or 4Ag + 8CN^– + 2H₂O + O₂ ? 4[Ag(CN)₂]^– + 4OH^–

Ag₂S + 4NaCN ? 2Na[Ag(CN)₂] + Na₂S

AgCl + 2NaCN ? Na[Ag(CN)₂] + NaCl

The reaction with the sulphide is reversible and accumulation of Na₂S must be prevented. A free excess of air is continuously passed through the solution which oxidizes Na₂S into sulphate and thiosulphate.

2Na₂ + 2O₂ + H₂O → Na₂S₂O₃ + 2NaOH

Na₂S₂O₃ + 2NaOH + 2O₂ → 2Na₂SO₄ + H₂O

2Na[Ag(CN)₂] + 4NaOH + Zn → Na₂ZnO₂ + 4NaCN + 2H₂O + 2Ag

(3) Extraction of Ag from argentiferrous lead (PbS + Ag₂S)– Parke’s Process : It is based upon the following facts (i) Molten Zn and Pb are immiscible, zinc forms the upper layer (ii) Ag is more soluble in molten Zn (iii) Zn-Ag alloy solidifies earlier than molten Pb (IV) Zn being volatile can be separated from Ag by distillation. Ag is purified by cupellation.

Properties of Silver: Silver is a white lustrous metal, best conductor of heat and electricity. Being soft, it is alloyed. The silver alloy used for making jewellery contain 80% Ag and 20% Cu. The composition of a silver alloy is expressed as its purity i.e. the amount of Ag present in 1000 parts of the alloy Ag does not react with dilute HCl or dil. H₂SO₄ and aqua regia but reacts with dil. HNO₃ and conc. HNO₃ forming NO and NO₂ respectively. Chlorine also reacts with Ag to form AgCl.

2Ag + Cl₂ → 2AgCl

Hot conc. H₂SO₄ reacts with Ag forming SO₂ like Cu

Zinc and its Compounds

(1) Occurrence of zinc: Zinc does not occur in the native form since it is a reactive metal. The chief ores of zinc are (i) Zinc blende (ZnS) (ii) Calamine or zinc spar (ZnCO₃) and (iii) Zincite (ZnO)

(2) Extraction of zinc : Zinc blende, after concentration by Froth floatation process, is roasted in air to convert it into ZnO. In case of calamine, ore is calcined to get ZnO. The oxide thus obtained is mixed with crushed coke and heated at 1673 K in fire clay retorts (Belgian Process) when ZnO gets reduced to metallic zinc. Being volatile at this temperature, the metal distils over and is condensed leaving behind Cd, Pb and Fe as impurities. The crude metal is called spelter. The metal may be refined either by electrolysis or by fractional distillation.

Properties of Zn : Zinc is more reactive than mercury. It is a good conductor of heat and electricity. Zinc readily combines with oxygen to form ZnO. Pure zinc does not react with non-oxidising acids (HCl or H₂SO₄) but the impure metal reacts forming Zn²⁺ ions and evolving H₂ gas.

Zn + 2HCl → ZnCl₂ + H₂↑

Hot and conc. H₂SO₄ attacks zinc liberating SO₂ gas

Zn + 2H₂SO₄ → ZnSO₄ + SO₂ + 2H₂O

Zinc also reacts with both dilute (hot and cold) HNO₃ and conc. HNO₃ liberating nitrous oxide (N₂O), ammonium nitrate (NH₄NO₃) and nitrogen dioxide (NO₂) respectively.

4Zn + 10HNO₃ (warm, dilute) → 4Zn(NO₃)₂ + N₂O + 5H₂O

4Zn + 10HNO₃ (cold very dilute) → 4Zn(NO₃)₂ + 2NO₂ + 2H₂O

Zn + 4HNO₃ (hot and conc.) → Zn(NO₃)₂ + 2N₂O + 2H₂O

Zinc dissolves in hot concentrated NaOH forming the soluble sod. Zincate

Zn + 2NaOH + 2H₂O → Na₂ [Zn(OH)₄] + H₂

Or Zn + 2NaOH → Na₂ZnO₂ + H₂

(3) Special varieties of zinc. (i) Zinc dust : It is prepared by melting zinc and then atomising it with a blast of air.

(ii) Granulated zinc : It is prepared by pouring molten zinc into cold water.

Both these varieties of zinc are used as reducing agents in laboratory.

Compounds of zinc

(1) Zinc oxide (Zinc white or Chinese white), ZnO : It is obtained by burning zinc in air or by heating zinc carbonate or zinc nitrate.

2Zn + O₂ → 2ZnO

ZnCO₃ → ZnO + CO₂

2Zn(NO₃)₂ → 2ZnO + 4NO₂ + O₂

It is a white powder but becomes yellow and heating and again white on cooling.

It is insoluble in water and is vey light and hence commonly known as philosopher's wool.

It is amphoteric in nature.

ZnO + 2HCl → ZnCl + H₂O

^(Black)

ZnO + 2NaOH → Na₂ZnO₂ + H₂O

^{(Acidic) Sod. sincate}

Or ZnO + 2NaOH + H₂O → Na₂[Zn(OH)₄]

^{Sod. tetrahydrozincate(II)}

It is reduced both by carbon adn H₂ and is used as a white paint

ZnO + C → Zn + CO; ZnO + H₂ → Zn + H₂O

(2) Zinc chloride, ZnCl₂ : It is obtained when Zn metal, ZnO or ZnCO₃ is treated with dil. HCl. It crystallizes as ZnCl₂.2H₂O and becomes anhydrous on heating. ZnCl₂ is highly deliquescent and is highly soluble in H₂O and also readily dissolves in organic solvents like acetone, alcohol, ether etc. its aqueous solution is acidic due to hydrolysis.

ZnCl₂ + H₂O → Zn(OH)Cl + HCl

Anhydrous ZnCl₂ is used as a Lewis acid catalyst in organic reactions. Mixed with moist zinc oxide, it is used for filling teeth and its solution is used for preserving timber. Anhydrous ZnCl₂ used as a Lucas reagent with conc. HCl.

(3) Zinc sulphide, ZnS : It is a white solid. It is soluble in dil. HCl and thus does not get precipitated by H₂S in the acidic medium.

ZnS + 2HCl → ZnCl₂ + H₂S

It is a constituent of lithopone (ZnS + BaSO₄)

(4) Zinc sulphate, ZnSO₄.7H₂O : It is commonly known as white vitriol and is obtained by the action of dil. H₂SO₄ on zinc metal, ZnO or ZnCO₃. On heating, it first loses six molecules of water of crystallization at 373 K. At 723 K, it becomes anhydrous and on further heating, it decomposes.

_{373K 723K}

ZnSO₄.7H₂O —————→ ZnSO₄.H₂O ————→

_1073K

2ZNSO₄ —————→ 2ZnO + 2SO₂ + O₂

It is used to prepare lithopone (BaSO4 + ZnS), a white paint and also in galvanishing iron.

ZnSO₄ + BaS → ZnS + BaSO₄

Elements	Outer electronic configuration	Oxidation states
Sc	3d¹ 4s²	+ 2, + 3
Ti	3d² 4s²	+ 2, + 3, + 4
V	3d³ 4s²	+ 2,+ 3,+ 4,+ 5
Cr	3d⁵ 4s¹	+ 1, + 2, + 3, + 4, + 5, + 6
Mn	3d⁵4s²	+ 2, + 3, + 4, + 5, + 6, + 7
Fe	3d⁶4s²	+ 2, + 3, + 4, + 5, + 6
Co	3d⁷4s²	+ 2, + 3, + 4
Ni	3d⁸4s²	+ 2, + 3, + 4
Cu	3d¹⁰4s¹	+ 1,+ 2
Zn	3d¹⁰4s²	+ 2

Thursday 22 January 2015

The D and F Block Elements

d-Block elements

Atomic Radii of D Block Element

Atomic radii

Actinides

Actinides

Properties of actinides

Oxidation states

Catalytic Properties of D block Elements

Catalytic properties

Chemical Properties of Potassium Dichromate

Chemical properties Of Potassium Dichromate

Chemical Properties of Potassium Permanganate

Chemical properties of Potassium Permanganate

Iodides to iodine in acidic medium

Hydrogen peroxide to oxygen

Chromium containing Compounds

Chromium Containing Compounds

Compounds of Copper

Compounds of Iron

Compounds of iron

Compounds Of Mercury

Compounds of Mercury

Compounds of Silver

Compounds of Silver:

Copper Ores and its Extraction

Electrode potentials(Eo)

Enthalpy of atomization

Formation of colored ions

Formation of complex ions

Gold and its Compounds

HYDROBORATION OF ALKENES: PREPARATION OF ALKANES

Ionic Radius of D Block Elements

Ionic Radius

Ionization Energy

Iron and its properties

F Block Elements & Lanthanides

Magnetic properties

Manganese containing compounds

MELTING AND BOILING POINTS OF D BLOCK ELEMENTS

Mercury Ores

Mercury Ores and Extraction

METALLIC CHARACTER OF D BLOCK ELEMENTS

Silver ores and its extraction

Zinc and its Compounds

Zinc and its Compounds

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Electrode potentials(E^o)