Chapter 8 d and f Block Elements – Chemistry free study material by TEACHING CARE online tuition and coaching classes

Chapter 8 d and f Block Elements – Chemistry free study material by TEACHING CARE online tuition and coaching classes

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)1to10

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.


Zn, Cd,   Hg    are called non typical transition element. Some exceptional electronic configuration of


transition element are :

Cr = 3d 5 4s1 ,

Nb = 4d 4 5s1 ,

Pb = 4d10 5s 0 ,

Ag = 4d10 5s1 ,

Cu = 3d10 4s1 ,


Mo = 4d 5 5s1 ,

Ru = 4d 7 5s1 ,

Pt = 5d 0 6s1 ,

Au = 5d10 6s1 . These irregularities can be explain on the basis of half


filled and full filled stability of d-orbital.

Classification : Transition element are classified in following series :


  • First transition series (3d)

= 21

to 30 i.e. Sc to Zn .


  • Second transition series (4d) = 39 to 48 e. Y to Cd .


  • Third transition series

(5d) = 57 La

and 72 – 80 Hf to Hg .


  • Fourth transition series or 6d series = 89 Ac and 104 – 112 Rf .

So there are 39 transition element at present in the periodic table.

1       2                                                                                                                                                                 18


H   Transition Elements (d-block elements) 13 14 15 16 17 He
    3         4         5         6         7         8         9         10        11       12            






















p-block elements




















    * 57




















    * 69 104 105 106 107 108 109 110    
  58-71 Lanthanide series (Lanthanides)
90-103 Actinide series (actinides)
Physico- Chemical Properties of d-Block Elements .
  • Atomic radii : The atomic, radii of 3d-series of elements are compared with those of the neighbouring
    • nd p-block


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

* in pm units

The atomic radii of transition elements show the following characteristics,



  • The atomic radii and atomic volumes of d-block elements in any series decrease with increase in the atomic 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.

  • 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

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.

  • Ionic radii : 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 (m): Sc Ti V Cr Mn Fe Co Ni Cu Zn
Lonic radius,(M2+)/pm: 90 88 84 80 76 74 72 69 74
Pm:(M3+)/pm: 81 76 74 69 66 64 63

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.

  • Ionisation energies : The ionisation energies of the elements of first transition series are given below:



lonisation E.

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

* in kJ mol-1

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

  • The ionisation energies of these elements are high, and in the most cases lie between those of s- and d- block 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.

  • 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



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.

  • The first ionisation energies of 5d-series of elements are much higher than those of the 3d and 4d series

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.

  • Metallic character : All the transition elements are 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.

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 covalant and metallic bonding in transition elements.

  • Melting and boiling points : The melting and boiling points of transition elements except Cd and Hg, are very high as compared to the s-block and p-block 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.

  • Enthalpies of atomization : Transition metals exhibit high enthalpies of

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.

  • 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 electronic configurations and oxidation states 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


Explanation : The outermost electronic configuration of the transition elements is (n – 1)d1-10ns2. Since, the energy levels of (n-1)d and ns-orbitals are quite close to each other, hence both the ns- and (n-1) d-electrons are available for bonding purposes. Therefore, the number of oxidation states show by these elements depends upon the number of d-electrons it has. For example, Sc having a configuration 3d14s2 may show an oxidation state of + 2 (only s-electrons are lost) and + 3 (when d-electron is also lost). The highest oxidation state which an elements of this group might show is given by the total number of ns- and (n -1) d-electrons.

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.

  • The variable oxidation states shown by the transition elements are due to the participation of outer ns and inner (n -1) d-electrons in
  • 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 elements. This means that after scandium, d-orditals become more stable than the s-orbital.

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

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

  • Ionisation energies and the stability of oxidation states :The values of the ionisation energies can be used in estimating the relative stability of various transition metal compounds (or ions). For example, Ni2+ compounds are found to be thermodynamically more stable than Pt2+, whereas Pt4+ compounds are more stable



than Ni4+ compounds. The relative stabilities of Ni2+ relative to Pt2+, and that of Pt4+ relative to Ni4+ can be explained as follows,

The first four ionisation energies of Ni and Pt


Metal (IE1+IE2) kJmol-1, (IE3+IE4) kJ mol-1,  


Etotal, kJ mo

IE1 + IE2 + IE3





Ni 2490 8800 11290
Pt 2660 6700 9360

Thus, the ionisation of Ni to Ni2+ requires lesser energy (2490 kJ mol-1) as compared to the energy required for the production of Pt2+ (2660 kjmol-1). Therefore, Ni2+ compounds are thermodynamically more stable than Pt2+ compounds.

On the other hand, formation of Pt4+ requires lesser energy (9360 kJ mol-1) as compared to that required for the formation of Ni4+(11290 kJ mol-1). Therefore, Pt4+ compounds are more stable than Ni4+ compounds.

This is supported by the fact that [PtCl6]2– complex ion is known, while the corresponding ion for nickel is not known. However, other factors which affect the stability of a compound are,

  • Enthalpy of sublimation of the
  • Lattice and the solvation energies of the compound or
  • Electrode potentials (Eo) : 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,

  • These metals should liberate hydrogen form dilute acids 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.

  • 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



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)

DH1 =

Enthalpy of sublimation, DHsub




M(g) ® M +(g) + e


M + (g) ® M + (aq)

DH2 = Ionisation energy, IE


DH3 = Enthalpy of hydration, DHhyd



Adding these equations one gets,


M(s) ® M +(aq) + e,           DH = DH1 + DH2 + DH3 = DHsub + IE + DHhyd


The DH 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 D H is negative. More negative is the value is of

DH , more favourable will be the formation of that cation from the metal. Thus, the oxidation state for which D 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) + n e

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

DH  = (DHsub   + lE + DHhyd ) or  E°  is  more  negative  (for  less positive) will be more stable.

  • Formation of coloured ions : Most of the compound of the transition elements are coloured in the solid state and /or in the solution The compounds of transition metals are coloured 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, shown in fig.

(l = 380 – 760nm) . Typical splitting for octahedral and tetrahedral geometries are


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 colour). This wavelength (or colour) of the absorbed light depends upon the energy difference of the two levels. Rest of the light gets transmitted. The transmitted light has a colour complementary to the absorbed colour. Therefore, the compound or the solution appears to be of the complementary colour. For example, Cu(H2O)26+ ions absorb red radiation, and appear blue-green (blue-green is complementary colour to red).Hydrated Co2+ ions absorb radiation in the blue-green region, and therefore, appear red in sunlight. Relationship between the colour of the absorbed radiation and that of the transmitted light is given in table

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


Colour of the   Colour 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 colours) except one are absorbed, then the colour of the substance will be the colour of the transmitted radiation. For example, if a substance absorbs all colours except green, then it would appear green to the eyes.

The transition metal ions which have completely filled d-orbitals are colourless, as there are no vacant d– orbitals to permit promotion of the electrons. Therefore, Zn2+ (3d10), Cd2 + (4d10) and Hg2+(5d10) ions are colourless. The transition metal ions which have completely empty d-orbitals are also colourless, Thus, Sc3+ and Ti4+.ions are colourless, unless a coloured anion is present in the compound.

Colours 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
  • 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. Paramgnetism 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).


ms =

BM =                         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 (m 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
Ti 3+ 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.

  • Formation of complex ions : Transition metals and their ions show strong tendency for complex 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 )6 ]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,

  • Small size and high nuclear charge of the transition metal
  • The availability to vacant inner d-orbitals of suitable
  • 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,
  • These are non-stoichiometric compounds and cannot be given definite
  • These compounds show essentially the same chemical properties as the parent metals, but differ in physical properties such as density and Steel and cast iron are hard due to the formation of interstitial compound with carbon. Some non-stoichimetric compounds are, Vse0.98 (Vanadium selenide), Fe0.94O, 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.

  • 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 Nickel powder is a good catalyst for hydrogenation of unsaturated organic compound such as, hydrogenation of oils some typical industrial catalysts are,
  • Vanadium pentoxide (V2O5) is used in the Contact process for the manufacture of sulphuric acid,
  • Finely divided iron is used in the Haber’s process for the synthesis of

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

  • The presence of vacant d-orbitals.
  • The tendency to exhibit variable oxidation
  • The tendency to form reaction intermediates with
  • The presence of defects in their crystal
  • Alloy formation : Transition metals form alloys among 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, e., these are less reactive as compared to s-block elements.

Explanation : Low reactivity of transition elements is due to,

  • Their high ionisation
  • Low heats of hydration of their
  • Their high heats of




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,

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




ç yellow÷

  • H2SO4 ®



ç orange ÷

  • K2SO4 + H2O


æ                   ö                                                                 æ                     ö

è                   ø                                                                 è                     ø

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

  • Manufacture from chromite ore : K2Cr2O7 is generally manufactured from chromite ore (FeCr2O4). The process involves the following
  • 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 Yellow mass due to the formation of sodium chromate is obtained.

4 FeCr2O4 + O2 ® 2Fe2O3 + 4Cr2O3

4Cr2O3 + 8 Na2CO3 + 6O2 ® 8 Na2CrO4 + 8CO2(g)


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

8 Na2CrO4

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,

4 FeCr2O4 + 16 NaOH + 7O2 ® 8 Na2CrO4 + 2Fe2O3 + 8H2O

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


2Na2CrO4  + H2SO4  ®

sodium chromate


sodium dichromate

  • Na2SO4 + H2O


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.

  • 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


Na2CrO7 + 2KCl ®


Physical properties



  • 2NaCl


  • Potassium dichromate forms orange-red coloured
  • It melts at 699
  • It is very stable in air (near room temperature) and is generally, used as a primary standard in the volumetric
  • It is soluble in water though the solubility is



Chemical properties

  • Action of heat : Potassium dichromate when heated decomposes to give oxygen.

4 K2Cr2O7 (s) ¾¾D ® 4 K2CrO4 (s) + 2Cr2O3(s) + 3O2

  • Action of acids
  • In cold, with concentrated H2SO4, red crystals of chromium trioxide separate

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

  • With HCl, on heating chromic chloride is formed and Cl2 is

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

  • Action of alkalies : With alkalies, it gives For example, with KOH,

K2Cr2O4 + 2KOH ® 2K2CrO4 + H2O

orange                                               yellow

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

2K2CrO4 + H2SO4 ® K2Cr2O7 + K2SO4 + H2O

Actually, in dichromate solution, the Cr O2 ions are in equilibrium with CrO2 ions.

2   7                                                                      4


Cr O2 + H O  2CrO2 + 2H +

2   7           2                     4

  • Oxidising nature : In neutral or in acidic solution, potassium dichromate acts as an excellent oxidising agent, and Cr O2 gets reduced to Cr3+. The standard electrode potential for the reaction,

2 7

Cr O2 + 14H+ + 6e ® 2Cr+3 + 7H O is +1.31 V. This indicates that dichromate ion is a fairly strong

2   7                                                           2

oxidising agent, especially in strongly acidic solutions. That is why potassium dichromate is widely used as an oxidising agent, for quantitative estimation of the reducing agents such as, Fe2+. It oxidises,

  • Ferrous salts to ferric salts

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

2FeSO4 + H2SO4 + [O] ® Fe2[SO4 ]3 + H2O ´ 3

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

Ionic equation: Cr O2 + 14 H+ + 6Fe2+ ® 2Cr3+ + 6Fe3+ + 7H O

2   7                                                                                 2

  • Sulphites to sulphates and arsenites to

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

Na2 SO3 + [O] ® Na2 SO4 ] ´ 3

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

Ionic equation: Cr O2 + 8H + + 3SO2 ® 2Cr3+ + 3SO2 + 4 H O

2   7                             3                                  4              2



Similarly, arsenites are oxidised to arsenates.

Cr O2 + 8H+ + 3 AsO3 ® 2Cr3+ + 3 AsO3 + 4 H O

2   7                                3                                    4              2

  • Hydrogen halides to

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

2HX + O ® H2O + X2] ´ 3

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

where, X may be Cl, Br, I.

Ionic equation :   Cr O2 + 8H + + 6HX ® 2Cr 3+ + 3X   + 7H O

2   7                                                           2           2

  • Iodides to iodine

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

2KI + H 2 O + [O] ® 2KOH + I 2 ] ´ 3

2KOH + H2SO4 ® K2SO4 + 2H2O] ´ 3

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



Ionic equation :

Cr2 O 2 + 14 H + + 6I ® 2Cr 3+ + 7H

2O + 3I 2



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

  • It oxidises H2S to S.

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

H2S + [O] ® H2O + S] ´ 3

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

Ionic equation : Cr O2 + 8H + + 3H S ® 2Cr3+ + 3S + 7H O

2   7                           2                                          2

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

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

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


K2Cr2O7  + 4 NaCl  + 6H2SO4  ¾¾he¾at ® 2KHSO4  + 4 NaHSO4  +


chromyl chloride

(orange red vapours )


Chromyl chloride vapours when passed through water give yellow-coloured solution containing chromic acid.


CrO2Cl2 + 2H2O ® 2HCl +


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,



  • An oxidising agent
  • In chrome tanning
  • The raw meterial for preparing large number of chromium compounds
  • Primary standard in the volumetric

Structures of Chromate and Dichromate Ions

Chromates and dichromates 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.

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.


MnO2 + 2KOH ¾¾fus¾e¾d ®



+ 4 H2O] ´ 2


2K2MnO3 + O2 ® 2K2MnO4 + 2 H2O




  • 4 KOH + O2 ¾¾fuse¾d ®



édark green massù

  • 2H2O


ëê                                                úû

Oxidation of potassium manganate to potassium permanganate : The potassium manganate so obtained is oxidised 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 oxidised to potassium permanganate and the hydrated manganese dioxide precipitates out. The reactions taking place are,

When CO2 is passed


3K2MnO4 +


2H2O ®


potassium permanganate

  • MnO2 ¯ +4 KOH



2CO2 + 4 KOH ® 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 oxidised electrolytically. The electrode reactions are,

At anode: 2MnO2 ® 2MnO + 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 KMnO4 has a purple colour.

Chemical properties : 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.

2KMnO4 (s) ¾¾he¾at ® K2MnO4 (s) + MnO2 + O2 (g )


Oxidising 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 MnO ) is indicated by high

positive reduction potentials for the following reactions.


Acidic medium: MnO + 8H+ + 5eMn2+ + 4 H O

Eo = 1.51 V



Alkaline medium:



MnO + 2H O + 3eMnO



+ 4OH  Eo = 1.23 V


4            2                               2


In strongly alkaline solutions and with excess of MnO , the reaction is



MnO4 –  + e


Eo = 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+ + MnO + 8H+ ® 5Fe3+ + Mn2+ + 4 H O

4                                                           2

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

3NO + MnO + 2OH ® 3NO + MnO + H O

2               4                                 3               2         2



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

2MnO + 3Mn+2 + 2H  O ® 5MnO  + 4 H+

4                             2                      2

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

MnO2 + 2MnO + 4OH ® 3MnO2 + 2H O

4                                    4              2


In acidic medium, KMnO4 oxidises,

Ferrous salts to ferric salts

2KMnO4 + 3H 2 SO4  ® K2 SO4  + 2MnSO4  + 3H 2O + 5[O]

2FeSO4 + H2SO4 + [O] ® Fe2 (SO4 )3 + H2O] ´ 5

2KMnO4 + 8H 2 SO4  + 10FeSO4  ® K 2 SO4  + 2MnSO4  + 5Fe 2 (SO4 )3  + 8H 2 O



Ionic equation:

2MnO + 16H + + 10Fe 2+  ® 2Mn2+ + 10Fe 3+ + 8H 2 O



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

Oxalic acid to carbon dioxide

2KMnO4 + 3H 2 SO4  ® K 2 SO4  + 2MnSO4  + 3H 2 O + 5[O]

(COOH )2 + [O] ® 2CO2 + H 2 O] ´ 5

2KMnO4 + 3H 2 SO4  + 5(COOH )2  ® K 2 SO4  + 2MnSO4  + 10CO2  + 8H 2 O



Ionic equation : 2MnO + 6H + + 5(COOH )2

Sulphites to sulphates

® 2Mn2+  + 10CO2

+ 8H 2 O


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

Na2SO3 + [O] ® Na2SO4 ] ´ 5

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

Ionic equation : 2MnO + 6H+ + 5SO2 ® 2Mn2+ + 5SO2 + 3H O

4                            3                                    4              2

Iodides to iodine in acidic medium

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

2KI + H 2 O + [O] ® I 2  + 2KOH ´ 5

2KOH + H2SO4 ® K2SO4 + 2H2O ] ´ 5

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



Ionic equation : 2MnO + 16H+ + 10I ® 2Mn2+ + 5I2

Hydrogen peroxide to oxygen

+ 8H2O


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 : 2MnO + 3Mn2+ + 2H O ® 5MnO + 4 H+

4                              2                      2

Ammonia to nitrogen

2KMnO4 + H 2 O ® 2MnO2 + 2KOH + 3[O]

2NH 3 + 3[O] ® N 2 (g ) + 3H 2 O

2KMnO4 + 2NH 3 ® 2MnO2 + 2KOH + 2H 2 O + N 2 (g )

Uses : KMnO4 is used,

  • As an oxidising (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.

Structure of Permanganate Ion (MnO4) : Mn in MnO4 is in +7 oxidation state.

Mn7+ exhibits sp3 hybridisation in this ion. The structure of MnO4 is, shown in fig.



  • Ores of iron : Haematite Copper pyrities (CuFeS2 )

Fe 2O3 , Magnetite

(Fe 3 O4 ),


(Fe 2 O3 .3H 2 O), Iron pyrites

(FeS2 ),


  • Extraction : Cast iron is extracted from its oxides by reduction with carbon and carbon monoxide in a blast furnace to give pig

Roasting : Ferrous oxide convert into ferric oxide.


Fe2 O3 . 3H 2 O ® Fe2 O3  + 3H 2 O ;

2FeCO3 ® 2FeO + 2CO2 ;

4 FeO + O2  ® 2Fe 2 O3


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

Note : ® The CO is the essential reducing agent.



Fe 2 O3  + 3CO


2Fe + 3CO2 . It is a reversible and oxothermic reaction. Hence according to



chatelier principle more iron will be produced in the furnace at lower temp.

Fe2O3 + CO ® 2FeO + CO2

(it is not reversible)



FeO + C ¾¾107¾3¾K ®

endothermic reaction

Fe + CO



Note : ®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;

  • Cast iron or Pig-iron : It is most impure form of iron and contains highest proportion of carbon (2.5–


  • Wrought iron or Malleable iron : It is the purest form of iron and contains minimum amount of carbon


  • 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 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,

  • To develop certain special properties like hardness, strength, ductility without changing the chemical composition.
  • To remove some undesirable properties or gases like entrapped gases, internal stresses and The various methods of heat treatment are,
  • Annealing : It is a process of heating steel to redness followed by slow
  • 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
  • Tempering : It is a process of heating the hardened or quenched steel to a temperature much below redness (473–623K) followed by slow
  • 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


  • Nitriding : It is a process of heating steels at about imparts a hard coating of iron nitride on the surface of

700 o C

in an atmosphere of ammonia. This process


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.

  • Low carbon or soft steels contain carbon upto 25%.
  • Medium carbon steels or mild steels contain 0.25–0.5%
  • High carbon or hard steels contains 1 – 1.5 percent carbon.


  • Alloy steels or special steels are alloys of steel with

Ni, Cr, Co,W, Mn, V

etc., For example – stainless steel


is an alloy of

Fe, Cr

and Ni and it is used for making automobile parts and utensils. Tool steel is an alloy of


Fe,W, V etc.

Uses of steel: In general, steels are used for making machinery parts, girders, tools, knives, razors, household utensils, etc. The specific use of steel depend upon the nature of metal added to iron.

Compounds of iron




  • Oxides of Iron : Iron forms three oxides magnetic oxide or load stone).

FeO, Fe2 O3


Fe3 O4 (magnetite also called


  • Ferrous oxide, FeO : It is a black powder, basic in nature and reacts with dilute acids to give ferrous

FeO + HSO4 ® FeSO4 + H 2 O ; It is used in glass industry to impart green colour to glass.


  • Ferric oxide

Fe 2 O3 :

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.

Fe2 O3 + 3C ® 2Fe + 3CO ;  Fe2 O3  + 3CO ® 2Fe + 3CO2

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


  • Ferrosoferricoxide

Fe 3 O4 (FeO. Fe2 O3 ) :

It is more stable than FeO and

Fe 2 O3 ,

magnetic in nature


and dissolves in acids giving a mixture of iron (II) and iron (III) salts.

Fe3 O4 + 4 H 2 SO4 (dil) ® FeSO4  + Fe2 (SO4 )3  + 4 H 2 O


(2)  Ferrous sulphide


It is prepared by heating iron filing with sulphur. With dilute

H 2 SO4 ,

it gives


H 2 S.        FeS + H 2 SO4 (dil) ® FeSO4 + H 2 S ­


(3)  Ferric chloride

FeCl3 : It is prepared by treating

Fe(OH)3 with HCl


Fe(OH)3 + 3HCl ® FeCl3  + 3H 2 O

The solution on evaporation give yellow crystals of FeCl 3 . 6H 2 O


Properties : (i) Anhydrous

FeCl 3

forms reddish-black deliquescent crystals.



FeCl 3

is hygroscopic and dissolves in

H 2 O giving brown acidic solution due to formation of HCl


FeCl 3 + 3H 2 O ® Fe(OH)3 + 3HCl



  • Due to oxidising nature

Fe 3+


FeCl3 is used in etching metals such as copper


2Fe 3+ + Cu ® 2Fe 2+ + Cu2+ (aq)


  • In vapour state




FeCl 3




exists as a dimer, Fe 2 Cl6





(4)  Ferrous sulphate,

FeSO4 , 7H 2O

(Green vitriol) : It is prepared as follow ,


Fe + H 2 SO4 ® FeSO4  + H 2

  • One pressure to moist air crystals become brownish due to oxidation by

4 FeSO4 + 2H 2 O + O2 ® 4 Fe(OH)SO4

  • On heating, crystals become anhydrous and on strong heating it decomposes to




Fe 2 O3 , SO2 and SO3 .


FeSO4 .7H 2 O ¾¾he¾at ® FeSO4  + 7H 2 O ;

2FeSO4  ¾¾Stro¾ng ® Fe 2 O3  + SO2  + SO3



  • It can reduce acidic solution of



K 2 Cr2 O7


  • It is generally used in double salt with ammonium sulphate.




(NH4 )2 SO4  + FeSO4  + 6H 2 O ® FeSO4 .(NH 4 )2 SO4 .6H 2 O

Mohr’s salt

Mohr’s salt is resistant to atmospheric oxidation.

  • It is used in the ring test for nitrate ions where it gives brown coloured ring of compound

FeSO4 + NO ® FeSO4 .NO



FeSO4 . NO.


(5)  Mohr’s salt

FeSO4  .(NH 4 )2  SO4 . 6H 2 O :

It is also known as ferrous ammonium sulphate and is a light


green coloured double salt.




Ores :      Copper pyrites (chalcopyrite)

CuFeS2 ,

Cuprite (ruby copper)

Cu2 O,

Copper glance

(Cu2 S) ,


Malachite [Cu(OH)2 . CuCO3 ], Azurite [Cu(OH)2 . 2CuCO3 ]

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

Concentration of ore : Froth floatation process.


Roasting : Main reaction :

2CuFeS2  + O2  ® Cu2 S + 2FeS + SO2 .


Side reaction : 2Cu2 S + 3O2 ® 2Cu2 O + 2SO2 ; 2FeS + 3O2  ® 2FeO + 2SO2 .


Smelting :

FeO + SiO2 ® FeSiO3 (slag) ;

Cu2O + FeS ® FeO + Cu2 S


Note : ® The mixture of copper and iron sulphides melt together to form ‘matte’ 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.


Cu2 S + 2Cu2O ® 6Cu + SO2

Note : ® 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. In presence of


and moisture Cu is covered with a green layer of


CuCO3 . Cu(OH)2 . 2Cu + H 2 O + CO2 + O2 ® CuCO3 .Cu(OH)2 . It undergoes displacement reactions with lesser


reactive metals e.g. with Ag. It can displace Ag from

Compounds of copper

AgNO3 . The finally divided Ag so obtained is black in colour.


Cuprous oxide

Cu2 O : It is a reddish brown powder insoluble in water but soluble in ammonia solution,


where it forms diammine copper (I) ion. glass industry.

Cu+ + 2NH 3 ® [Cu(NH 3 )2 ]+ . It is used to impart red colour to glass in


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  ¾¾D ® 2CuO + 4 NO2  + O2




Copper sulphate

CuSO4 . 5H 2O

(Blue vitriol) : It is prepared by action of dil


on copper scrap in


presence of air.

2Cu + 2H 2 SO4 + O2  ® CuSO4 + 2H 2 O



  • On heating this blue salt becomes white due to loss of water of

CuSO4 . 5H 2 O ¾¾503¾K ® CuSO4 + 5H 2 O

Blue                    White

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



¾¾100¾0¾K ® CuO + SO3


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

Cu2 SO4 + 4 NH 4 OH, ® [Cu(NH 3 )4 ]SO4 + 4 H 2 O

Blue colour

NH 4 OH.


  • With KCN it first gives yellow precipitate of CuCN which decomposes of give

Cu2 (CN)2 .

Cu2 (CN)2


dissolves in excess of KCN to give

K 3 [Cu(CN)4 ]


2CuSO4 + 4 KCN ® Cu2 (CN)2  + 2K 2 SO4  + (CN)2

  • With KI it gives white of Cu2 I 2

4 KI + 2CuSO4 ® 2K 2 SO4 +   Cu2 I 2   + I 2

White ppt.

  • With K 4 [Fe(CN)6 ], CuSO4 gives a reddish brown of Cu2 [Fe(CN)6 ]


2CuSO4 + K 4 [Fe(CN)6 ] ®

Cu2 [Fe(CN)6 ] + 2K 2 SO4

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. As a fungicide in starch paste for book binding work.


Cupric sulphide

CuS : It is prepared as follows : Cu(NO3 )2 + H 2 S ®

CuS +

Black ppt.

2HNO3 .


Cupric chloride

CuCl2   :

It is a dark brown solid soluble in water and its aqueous solution first changes to


green and then to blue on dilution.


Cuprous chloride

Cu2 Cl2

: It is a white solid insoluble in water and dissolves in conc.  HCl due to


formation of

H[CuCl2 ] complex.





Ores : Argentite (silver glance)

Ag 2 S, Horn silver (AgCl), Ruby silver (Pyrargyrite) 3 Ag 2 S. Sb2 S3 .


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.

4 Ag + 8 NaCN + 2H 2 O + O2 ⇌ 4 Na[Ag(CN)2 ] + 4 NaOH




or    4 Ag + 8CN + 2H 2 O + O2 ⇌ 4[Ag(CN)2 ] + 4OH

Ag 2 S + 4 NaCN ⇌ 2Na[Ag(CN)2 ] + Na 2 S


AgCl + 2NaCN Na[Ag(CN)2 ] + NaCl.

The reaction with the sulphide is reversible and accumulation of



must be prevented. A free excess of


air is continuously passed through the solution which oxidizes

2Na 2 + 2O2  + H 2 O ® Na 2 S2 O3  + 2NaOH

Na 2 S2 O3 + 2NaOH + 2O2  ® 2Na 2 SO4  + H 2 O


into sulphate and thiosulphate.


2Na[Ag(CN)2 ] + 4 NaOH + Zn ® Na 2 ZnO2 + 4 NaCN + 2H 2 O + 2Ag


Compounds of silver :

AgNO3 , Ag 2 S, AgCl, AgBr, AgI,






Ores : Bismuthaurite ( BiAu2 ), Syvanite (AgAuTe 2 ) , Calverite (AuTe 2 ) .

Extraction : By cyanide or Mac-Arther forest cyanide process,

4 Au + 8 NaCN + 2H 2 O + O2 ® 4[NaAu(CN)2 + 4 NaOH]

Sodium aurocyanide


2Na[Au(CN)2 + Zn ® Na2 [Zn(CN)4 ] + 2Au

Refining : Anode : Crude gold; Cathode : Pure gold. Electrolytic Solution : Gold chloride in hydrochloric acid. Plattner chlorine extraction process,

AuCl 3 + 3FeSO4 ® FeCl 3 + Fe 2 (SO4 )3 + Au AuCl 2 + 3H 2 S ® 6HCl + 3S + 2Au

Quartation process : Refining of gold carried out by this method. It involves separation of gold and Ag by

H 2 SO4 .


Gold is soft and hence for making ornaments it is generally hardened by adding Ag or gold is expressed in terms of Carats. Pure gold is taken as 24 carats.

20 carats means, it contain 20 parts by wt. of gold in 24 parts by wt. of given alloy.

\ percentage of gold in 20 carat gold sample = 20 ´ 100 = 250 = 83.33%


The weight of



Properties : Gold is not affected by conc.


H 2 SO4 , conc.


HNO3 or by strong alkalis. However it dissolves in


aqua regia to form H[AuCl 4 ] ;

Compounds of gold

2Au + 3HNO3  + 11HCl ® 2H[AuCl 4 ] + 6H 2 O + 3NOCl .


AuCl3 : It is a reddish solid soluble in water. It reacts with HCl to give H[Au(Cl)4 ] which is used in toning


process in photography.

HCl + AuCl 3  ® H[Au(Cl)4 ]



Au2S : It is a dark brown solid insoluble in water prepared as follows :

2K[Au(CN)2 ] + H 2 S ® Au2 S + 2KCN + 2HCN


Ores : Cinnabar (HgS)

Extraction : Roasting : The concentrated ore roasted at 770 K to 780 K in the pressure of air.

2HgS + 3SO2  ® 2HgO + 2SO2 ;     2HgO ® 2Hg + O2

Refining : By filtering impure Hg through thick canvass or chamois leather. It is then dropped into 5% HNO3 .

Compounds of Mercury


Mercuric chloride

HgCl 2 (Corrosive sublimate) : It is a colourless solid, sparingly soluble in water. It forms


red ppt. of

HgI 2


KI :

HgCl2 + 2KI ® HgI 2 + 2KCl . With


it gives white ppt. of

Hg(NH 2 )Cl.


HgCl 2 + 2NH 4 OH ® Hg(NH 2 )Cl+ NH 4 Cl + 2H 2 O .

white ppt.


Mercurous chloride

Hg 2Cl2 (Calomel) : It is a white solid insoluble in water. With


it forms a


black mixture composed of black metallic mercury and white mercuric amino chloride,

Hg 2 Cl2  + 2NH 4 OH ® 1Hgu+uHug2(NuHu2 u)3Cl + NH 4 Cl + 2H 2 O

Black mixture

It is used as purgative in medicine and it sublimes on heating.

Hg(NH 2 )Cl.


Mercuric iodide

HgI 2 : It is a yellow solid below 400K but changes to red solid above 400K.


HgI 2 400 K


HgI 2



It dissolves in excess of KI forming

KHgI 4 ;

HgI 2 + 2KI ® K 2 HgI 4


Alkaline solution of K 2 HgI 4

is called Nessler’s reagent.





Ores : Zincite (red zinc ore)

ZnCO3 .



(ZnOFe2 O3 ) , Zinc blende


Calamine ( Zinc spar)


Extraction : Concentration : Froth floatation


Roasting :

ZnS + 3O2  ¾¾120¾0¾K ® 2ZnO + 2SO2 ;

ZnS + 2O2 ¾¾D ® ZnSO4


2ZnSO4  ¾¾D ® 2ZnO + 2SO2  + O2 .

Reduction of ZnO : The oxide ore is mixed with crushed coke and heated to about 1670K in fire clay retorts (Belgian process). The crude metal obtained called Zinc spelter.

Refining : By distillation and by electrolytic method

Anode : Spelter; Cathode : Pure zinc wire; Electrolyte : Zinc sulphate.



Note : ® Zinc is a volatile metal (easily vaporisable)

  • At ordinary temperature zinc metal is brittle but on heating at 120 – 150 o C it is malleable and

Compounds of zinc

Zinc oxide ZnO : Zincite (ZnO) is also called Philospher’s wool. It is white powder, become yellow on heating and again white on cooling. It is amphoteric in nature. It is used as a white pigment under the name Zinc white or Chinese white.


Zinc Sulphate (white vitriol),

ZnSO4 .7H 2 O : It is a colourless transparent crystal highly soluble in water.


It is used as an eye-lotion and for preparing double salts. On heating it looses its molecules of water as,

ZnSO4 .7H 2 O ¾¾375¾K ® ZnSO4 .H 2 O ¾¾725¾K ® ZnSO4  ¾¾107¾5¾K ® ZnO + SO2  + O2



Ores of Tin : Cassiterite or tin stone (SnO2 )


  • The concentrated and roasted ore is reduced with carbon to get impure tin which is purified by liquation process.
  • Tin forms two series of salts, e., Sn (II) and Sn (IV). Whereas Sn (II) salts are ionic Sn (IV) salts are covalent.


  • Tin dissolves in hot NaOH forming

Na 2 SnO3 and evolving

H 2 gas.


  • Tin reacts with

HNO3 forming metastannic acid (H 2 SnO3 ) .


  • Tin is not attacked by organic acids and hence is used for tinning of utensils to resist Tin foils are used for wrapping cigarettes, confectionary items and for making tooth-paste tubes.


  • SnO2

is an amphoteric oxide.


  • Stannous chloride

(SnCl2 ) acts as a good reducing agent. It reduces

HgCl 2

to first

Hg 2 Cl2

and then to


Hg. It also reduces FeCl 3 to FeCl 2 .

  • Stannic chloride (SnCl4 ) is a liquid and fumes in air due to It acts as a Lewis acid and dissolves


in concentration HCl forming

H 2 SnCl6 .


  • SnCl4 .5H 2 O is called butter of tin.
  • SnS dissolves in yellow ammonium




Ores of lead : Galena (PbS), Anglesite (PbSO4) and Cerussits (PbCO3 ).

  • Lead is extracted from The ore is concentrated by froth-floatation process and roasted when a part


of the ore is converted into PbO and

PbSO4 . The unchanged galena then brings about the reduction of PbO or



to Pb.


  • Lead dissolves in hot NaOH forming sodium plumbite (Na 2PbO2 ) and evolving H 2



  • Leads forms two series of salts, i.e., Pb (II) and Pb (IV) but Pb (II) compounds are more stable than Pb (IV) Lead (II) compounds are essentially ionic while lead (IV) compounds are covalent.
  • Lead is used in making bullet shots, lead accumulators, tetraethyl lead (antiknocking agent) and a number of pigments such as red lead (Pb3O4 ) , white lead or basic lead carbonate [2 PbCO3 .Pb(OH)2 ] and lead chromate

(PbCrO4 ) .


  • Litharge is PbO . It is obtained by heating

Pb(NO3 )2

or PbCO3 . It is an amphoteric oxide and is reduced


back to Pb by

H 2 , C and CO.




  • Red lead

(Pb3 O4 )

or Sindhur is a mixed oxide

(2PbO.PbO2 ) . It acts as an oxidising agent and as such


oxidises HCl to Cl2.

  • Lead dioxide (PbO2 )

is obtained either by treating


with conc.


or by treating lead acetate


with bleaching powder. It acts as an oxidising agent and oxidises HCl to Cl2.


  • The ionic character of lead dihalides decreases the order:

PbF2 > PbCl 2 > PbBr2 > PbI 2 .


  • PbF4



are stable while



PbI 4

are however, unknown. The non-existence of


PbBr4 and

PbI 4 is due to strong oxidising character of

Pb4 +

ions and reducing character of


and I





  • PbF4

is ionic while


is a volatile liquid.


  • Lead is readily corroded by water containing dissolved air forming


which has appreciable


solubility in water. This action of water on lead called is Plumbosolvency. 2Pb + 2H 2O + O2 ® 2Pb(OH)2



Whereas organic acids,

NH +

salts and nitrates increase while salts like carbonates, phosphate and sulphates



decrease Plumbosolvency. Hard water, however, has no solvent action on lead.


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 0-10ns 2 .



  • 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]4 f 1-14 5d 0-10 6s 2 . Promethium (Pm), atomic number 61 is the only synthetic (man made) radioactive

Properties of lanthanides

  • These are highly dense metals and possess high melting
  • They form alloys easily with other metals especially 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.
  • Oxidation state : Most stable oxidation state of lanthanides is + Oxidation states + 2 and + 4 also

exist but they revert to + 3 e.g. Sm2+ , Eu 2+ , Yb 2+ 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.

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


their magnetic moments do not obey the simple “spin only” formula

m eff  =

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 4 f– 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


m eff =



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

  • 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, b -diketones, oxine etc. No complexes with p -bonding ligands are


  • Lanthanide contraction : The regular decrease in the size of lanthanide ions from

La 3+

to Lu 3+ 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

  • It results in slight variation in their chemical properties which helps in their separation by ion exchange
  • 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).





  • The covalent character of hydroxides of lanthanides increases as the size decreases from

La 3+

to Lu3+ .


However basic strength decreases. Thus


is most basic whereas


is least basic. Similarly, the


basicity of oxides also decreases in the order from

  • Tendency to form stable complexes from

La 3+

La 3+

to Lu3+ . to Lu 3+


increases as the size decreases in that order.


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


  • Since the radius of

Yb 3+

ion (86 pm) is comparable to the heavier lanthanides Tb, Dy, Ho and Er,


therefore they occur together in natural minerals.

  • Actinides : The elements with atomic numbers 90 to 103 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 7s 2 .

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

  • Oxidation state : The dominant oxidation state of actinides is +3 which shows increasing stability for the heavier 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., UO 2+ , NpO 2+ etc.

2               2

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.

  • 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 5 f electron by another in the same This results in increase in the effective nuclear charge which causes contraction in size of the electron cloud.
  • Colour of the ions : Ions of actinides are generally coloured which is due to f f It depends upon the number of electrons in 5 f orbitals.
  • 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 5 f electrons of actinides are less effectively shielded which results in quenching of orbital
  • 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, b-diketones, oxine The degree of complex formation decreases in the order.


M 4 +

  • MO2+
  • M 3+
  • MO+






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


MO 2+