8. d and f block elements 01

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Topic – D AND F BLOCK ELEMENTS
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1. Transition Elements:
The d-block elements are also called transitions because their properties are
intermediate between those of the S and P-block elements. They make up three
complete rows of 10 elements corresponding to the filling of 3d, 4d and 5d orbitals and
an incomplete fourth row in the periodic table
2. Electronic Configuration of Transition Elements:
(n−l) d1−10
ns1−2
(n−1) stands for the inner shell and the d orbitals may have one to ten
electrons and the s orbital of the outermost shell (n) may have one or two electrons
Thus, a transition element is defined as an element which has incompletely filled d
orbitals in its ground state or in any one of its oxidation states
Zinc, cadmium and mercury of group 12 have full d10
configuration in their ground state
as well as in their common oxidation states and hence, are not regarded as transition
metals However, being the end members of the three transition series, their chemistry
is studied along with the chemistry of the transition metals
3. General Characteristics of d−Block Elements:
A. Atomic Radii. Some general trends in the variation of atomic radii across the
period are:
(i) On progressive increase in atomic number, due to poor shielding effect of d
electrons, the net electrostatic attraction between the nuclear charge and the
outermost electron increases. Thus the atomic radii of the d-block elements of
a given series generally decrease with increase in the atomic number
(ii) Decrease in size is, however, very small as we move from chromium to copper.
This is explained in terms of screening effect
(iii) In a transition series, the atomic radius reaches a minimum for the group 8
elements and increases again towards the end of the series. This is explained in
terms of the increased force of repulsion among the added electrons which
exceeds the attractive force due to the increased nuclear charge and results in
the expansion of the electron cloud
B. Ionic Radii. The ionic radii follow the same trend as the atomic radii
C. Densities. The decrease in metallic radius coupled with increase in atomic mass
results in a general increase in the density of these elements
D. Metallic Character. All the d-block elements are metals. They are hard,
malleable and ductile, and have high melting and boiling points
E. Melting and Boiling Points. Melting and boiling points of transition metals are
very high due to strong inter atomic bonding which involves the participation of
both ns and (n−l)d electrons. Melting and boiling points increase on moving from
left to right across the transition series and reach a maximum at d5
where after
they start decreasing. This is explained in terms of covalent bonds formed by
unpaired electrons
F. Enthalpies of Atomization. They have high enthalpies of atomization with a
maxima at about the middle of each series indicating that one unpaired electron

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per d orbital is particularly favorable for strong inter atomic interaction. In
general, greater the number of valence electrons, stronger is the resultant
bonding. Since the enthalpy of atomization is an important factor in determining
the standard electrode potential of a metal, metals with very high enthalpy of
atomization (i.e., very high boiling point) tend to be noble in their reactions
Further, the metals of the second and third series have greater enthalpies of
atomization than the corresponding elements of the first series resulting in
greater metal- metal bonding in compounds of the heavy transition metals.
G. Ionization Enthalpies. Ionization enthalpy of each series of transition elements
increases from left to right with increasing nuclear charge which accompanies
the filling of the inner d orbitals. Although the first ionization enthalpy in
general increases, the increase in the second and third ionization enthalpies for
the successive elements is not of the same magnitude.
I.E.1 value of Cr is lower as compared to the general trend because of the absence of
any change in the d configuration and the value for Zn is higher since it involves
ionization from the 4s level
I.E.2 values increase smoothly with atomic number except for chromium and copper due
to the extra stability of their half or fully filled d orbitals
H. Oxidation States: The transition elements show a large number of oxidation
states which are related to their electronic configurations. This can be
illustrated by taking into consideration the oxidation states of the elements of
the first transition series
The variability of oxidation states arises out of incomplete filling of d orbitals.
I. M2+
/M Standard Electrode Potentials. M2+
/M: Standard reduction potential
of transition elements are generally negative, i.e. lower than that of hydrogen.
All those elements with negative reduction potential liberate hydrogen from
dilute acids
------->+ 2+
2M + 2H M + H (g)
Only exceptions are copper and mercury (with positive −E°). Only oxidizing acids (nitric
and hot concentrated sulphuric) react with Cu, the acids being reduced. The high
energy to transform Cu(s) to Cu2+
(aq) is not balanced by its hydration enthalpy
Some of these metals (e.g. Cr) are unreactive due to the formation of the protective
film of oxide on their surface
J. The highest value for Zn is due to the removal of an electron from the stabled
d10
configuration of Zn+
a) The comparatively high value for Mn shows that Mn2+
(d5
) is particularly stable
b) The comparatively low value for Fe shows the extra stability of Fe3+
(d5
)
c) The comparatively low value for V is related to the stability of V2+
(half-filled
t2g level)
K. Stability of Higher Oxidation States: Some noteworthy points are

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a) Fluorine stabilizes the highest oxidation state due to either higher lattice
energy as in the case of CoF3, or higher bond enthalpy terms for the higher
covalent compounds, e.g., VF5 and CrF6
b) Fluorides are unstable in the low oxidation states, e.g. CuX. On the other hand,
all CuII
halides are known except the iodide. In this case, Cu2+
oxidizes I-
to I2
2Cu2+
+4I−
→Cu2I2 (s) +I2
c) Many copper (I) compounds are unstable in aqueous solution and undergo
disproportionation
2Cu+
→Cu2+
+Cu
The stability of Cu2+
(aq) rather than Cu+
(aq) is due to the much more negative
∆hyd Ho
of Cu2+
(aq) than Cu*, which more than compensates for the second
ionization enthalpy of Cu
d) The ability of oxygen to stabilize the highest oxidation state is demonstrated in
the oxides
e) The highest oxidation number in the oxides coincides with the group number and
is attained in Sc2O3 to Mn2O7
f) Oxygen is superior to fluorine in stabilizing the higher oxidation states due to
its ability to form multiple bonds to metals. For example, the highest Mn
fluoride is MnF4 whereas the highest oxide is Mn2O7. In Mn2O7, each Mn is
tetrahedrally surrounded by O's including a Mn-O-Mn bridge
L. Chemical reactivity and E° values: Some noteworthy points are,
1) Many transition metals are sufficiently electropositive to dissolve in mineral
acids, although a few are unaffected ('noble') by simple acids
2) The metals of the first series with the exception of copper are relatively more
reactive and are oxidized by 1M H+
3) The E° values for M2+
/M indicate a decreasing tendency to form divalent cations
across the series. This general trend towards less negative E° values is related
to the increase in the sum of the first and second ionization enthalpies
4) E° values for Mn and Zn are more negative than expected from the general
trend due to the stabilities of half-filled d subshell (d5
) in Mn2+
and completely
filled d subshell (dl0
) in zinc
5) Eo
values for Ni is more negative than expected from the general trend due to
its highest negative enthalpy of hydration
6) Mn+3
and Co+3
ions are the strongest oxidizing agents in aqueous solutions due
their high positive E° values for the redox couple M3+
/M 2+
7) The ions Ti2+
, V2+
and Cr2+
are strong reducing agents (due their negative E°
values for the redox couple M3+
/M2+
) and will liberate hydrogen from a dilute
acid, e.g.
2 Cr2+
(aq) + 2 H+
(aq) → 2 Cr3+
(aq) + H2(g)
M. Tendency to form Complexes: The cations of d-block elements have a tendency
to form complex ions with certain molecules (e.g., CO, NO and NH3) or ions (e.g.,
F-
, Cl-
, CN-
, etc.)

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Two factors which account for this tendency are:
(i) Transition metal cations because of their small size, and high effective nuclear
charge, have a high charge density. These can, therefore, accept lone pairs of
electrons from other molecules or ions
(ii) These have vacant inner d-orbitals with appropriate energy to accept lone pairs
of electrons and form coordinate bonds
Colour: When an electron from a lower energy d orbital is excited to a higher energy d
orbital, the energy of excitation corresponds to the frequency of the light absorbed.
The colour observed corresponds to the complementary colour of the light absorbed.
The colour observed is determined by the nature of the ligand
N. Magnetic Properties: Para magnetism in transition elements arises from the
presence of unpaired electrons in atoms, ions, complex ions or molecules, with
each such electron having a magnetic moment associated with its spin angular
momentum and orbital angular momentum
For the compounds of the first series of transition metals, the contribution of orbital
angular momentum is of no significance. In terms of n (number of unpaired spins),
magnetic moment is given by the formula
sµ = n(n+2)
Thus for n=1, µ 1.73 BM; n=2, µ=2.83 BM; n=3, µ=3.87 BM; and so on
O. Catalytic Properties: Transition metals show catalytic activity due to their
ability to adopt multiple oxidation states and to form complexes. They can thus
form unstable intermediate products with various reactants in some cases.
These intermediate products decompose to give the final product, regenerating
the catalyst
Catalysts at a solid surface involve the formation of bonds between reactant
molecules and atoms of the surface of the catalyst. This results in increasing
the concentration of the reactants at the catalyst surface and also weakening
of the bonds in the reacting molecules (the activation energy is lowered)
P. Formation of Interstitial Compounds: Interstitial compounds are those in
which small-sized elements (like H, B, C and N) get entrapped inside the
interstitial spaces of the transition metal lattice. They are usually
non−Stoichiometric and are neither typically ionic nor covalent.
Non−stoichiometry is due to variable valence of these elements and due to the
defects in solid structures
The principal physical and chemical characteristics of these compounds are as
follows:
(i) They have high melting points, higher than those of pure metals
(ii) They are very hard, some borides approach diamond in hardness
(iii) They retain metallic conductivity
(iv) They are chemically inert
Q. Alloy Formation: Alloy is defined as the homogeneous solid solution containing
two or more different elements, at least one of which is a metal. Solid solution

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alloys are formed by atoms with metallic radii that are within about 15 percent
of each other. Because of similar radii and other characteristics of transition
metals, alloys are readily formed by these metals. The alloys so formed are hard
and have often high melting points. For example, stainless steel is an alloy of Fe,
Ni and Cr. Another example of solid solution alloy is bronze which an alloy of
copper and tin is
4. Oxides and Oxometal Ions of Transition Elements: All the metals except
scandium
Form MO oxides which are ionic the highest oxidation number in the oxides,
coincides with the group number and is attained in Sc2O3 to Mn2O7
As the oxidation number of a metal increases, ionic character decreases. In
higher oxides, the acidic character is predominant
5. Potassium Dichromate, K2Cr2O7: Different steps involved in the preparation of
potassium dichromate from chromite are
(i) Preparation of sodium chromate. Finely powdered chromite is mixed with soda
ash and quicklime and then roasted in a reverberatory or rotating furnace in
excess of air. The mass turns yellow due to the formation of sodium chromate
4Fe (CrO2)2+8Na2CO3+7O2  8Na2CrO4+2Fe2O3+8CO2
The yellow mass is extracted with water when the ferric oxide is left behind.
(ii) Conversion of chromate into dichromate. Sodium chromate solution obtained in
(i) is treated with concentrated sulphuric acid when it is converted into sodium
dichromate
2Na2CrO4+H2SO4  Na2Cr2O7+Na2SO4+H2O
Orange sodium dichromate Na2Cr2O7.2H2O can be crystallized from this solution
(iii) Conversion of sodium dichromate into potassium dichromate. Sodium dichromate
is treated with potassium chloride. Potassium dichromate, being the less soluble,
crystallizes out
Na2Cr2O7+2KC1  K2Cr2O7+2NaCl
6. Properties of Potassium Dichromate:
(i) It gives orange-red crystals, soluble in water (m.p. 671 K)
(ii) Action of heat. It decomposes on heating to a white heat and gives out oxygen
4K2Cr2O7  4K2CrO4+2Cr2O3+3O2
(iii) Action of alkalies. With an alkali, it gives a chromate
K2Cr2O7+2KOH  2K2CrO4+H2O
Chromates and dichromates in solution are inter-convertible. Even weak acids
cause the change from yellow chromate to orange dichromate, while this is
reversed by addition of an alkali
2CrO4
−
+2H+
 Cr2O7
2−
+H2O
Cr2O7
2−
+2OH−
 2CrO4
2−
+H2O
2HCrO4
−
+CO3
2−
 2CrO4
2−
+H2O+CO2
(iv) With concentrated sulphuric acid, in the cold, potassium dichromate gives red
crystals of chromic anhydride, CrO3

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K2Cr2O7+2H2SO4 2CrO3+2KHSO4+H2O
On heating the mixture, oxygen is evolved
2K2Cr2O7+8H2SO4 2K2SO4+2Cr2 (SO4)3+8H2O+3O2
(v) Oxidizing action. In neutral or in acidic solution, potassium dichromate furnishes
nascent oxygen and thus acts as an excellent oxidizing agent
Cr2O7
2−
+14H+
+6e−
 2O3+
+7H2O
For example:
(a) It oxidizes ferrous sulphate to ferric sulphate
K2Cr2O7+7H2SO4+6FeSO4  K2SO4+Cr,2(SO4)3+3Fe2(SO4)3+7H2O
or Cr2O7
2−
+14H+
+6Fe2+
 2Cr3+
+7H2O+6Fe3+
(b) Hydrogen sulphide is oxidized to sulphur.
K2Cr2O7+4H2SO4+3H2S  K2SO4+Cr2(SO4)3+7H2O+3S
or Cr2O7
2−
+8H+
+3H2S  2Cr3+
+3S+7H2O
In a similar manner it oxidizes sulphur dioxide to sulphuric acid, halogen acids to
halogens and liberates iodine from potassium iodide
7. Uses of Potassium Dichromate:
(i) As an oxidizing agent in organic chemistry
(ii) In volumetric analysis (oxidation titrations)
8. Structure of Chromate and Dichromate ions:
9. Potassium Permanganate KMnO4: Is prepared by fusing manganese dioxide
(pyrolusite) with an alkali metal hydroxide and an oxidizing agent like potassium
nitrate. The fused mass is green due to the formation of potassium magnate
2MnO2+4KOH+O2  2K2MnO4+H2O
The fused mass disproportionates in a neutral or acidic solution to give purple
permanganate
3MnO4
−
+2H2O  2MnO4
−
+MnO2+4OH−
Commercially, it is prepared by the electrolytic oxidation of the manganate
followed by the electrolytic oxidation of manganate(VI)
3
fuse with KOH, oxidize
witha.rorKNO 2
2 4MnO MnO
manganate


electrolytic oxidation
in alkaline solution2
4 4MnO MnO
permanganatemanganate
 

10. Properties of Potassium Permanganate:
(i) It is moderately soluble in water giving a purple solution
(ii) Action of heat. When heated strongly, potassium permanganate decomposes to
give potassium manganate, manganese dioxide and oxygen

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2KMnO4  K2MnO4+MnO2
(iii) Oxidizing action. Potassium permanganate acts as a powerful oxidizing agent in
neutral, alkaline or acidic solution
Medium Complete equation Ionic equation
Neutral
Alkaline
Acidic
2KMnO4+H2O→2KOH+2MnO2+3O
2KMnO4+2KOH→2K2MnO4+H2O+O
2KMnO4+3H2SO4→K2SO4+2MnSO4
+3H2O+5O
MnO4
−
+2H2O+3er→MnO2+4OH−
MnO4
−
+2H2O +3e−
→ MnO2+4OH−
MnO4
−
+8H+
+5e−
→Mn2+
+4H2O
Oxidation in neutral medium: In neutral solution, it oxidizes manganous sulphate to
manganese dioxide
2KMnO4+3MnSO4+2H2O  5MnO2+K2SO4+2H2SO4
This may be represented in the ionic form as
2MnO4"+3Mn2+
+2H2O  5MnO2+4H+
Oxidation in alkaline medium: In alkaline medium it oxidizes potassium iodide to
potassium iodate
I−
+6OH−
 IO3
−
+3H2O+6e−
Alkaline potassium permanganate is used for the oxidation of a number of organic
compounds
Oxidation in acidic medium In acidic medium it oxidizes sulphides or hydrogen sulphide
to sulphur, sulphur dioxide to sulphuric acid, sulphides to sulphates, nitrites to nitrates,
arsenites to arsenates, oxalates or oxalic acid to carbon dioxide, ferrous salts to ferric
salts, and hydrogen peroxide to water+oxygen
2KMnO4+3H2SO4 K2SO4+2MnSO4+3H2O+5O
(Effective equation)
or
MnO4
−
+8H+
+5e-
 Mn2+
+4H2O
S2−
 S+2e-
SO2+2H2O  SO4
2+
+4H+
+2e−
SO3
2−
+H2O  SO4
2−
+2H+
+2e−
NO2−
+H2O  NO3−
+2H+
+2e−
AsO33+
+H2O  AsO4
3+
+2H+
+2e−
2
2
COO
2CO + 2e
COO


 
Fe2+
 Fe3+
+e−
H2O2  O2+2H+
+2e−
11. Uses of Potassium Permanganate:
(i) As an oxidizing agent in the laboratory
(ii) As a disinfectant and bleaching agent
(iii) In volumetric analysis

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(iv) Alkaline KMnO4 is used in organic chemistry under the nameBaeyer's reagent
12. Structure of Manganate and Permanganate ions:
13. Inner Transition or F-Block: Elements comprise of two series, Lanthanoids and
actinoids in which f orbitals are progressively filled. As the f-orbitals lie
comparatively deep within the kernel, they are also termed as inner-transition
elements
14. Lanthanoids: Have electronic configuration 6s2
4fn
(n=1 to 14)
15. Lanthanoids Contraction: In lanthanoid atoms (or ions), the nuclear charge
increases with atomic number while the differentiating electron is being added
in an inner orbital (4f). The shielding of one 4f electron by another from the
increasing nuclear charge is much more imperfect than that those encountered
in the case of d-electrons. This imperfect shielding is due to the difference in
the shapes of the orbitals. Thus, as atomic number increases, the effective
nuclear charge experienced by each 4f electron also increases and hence causes
a slight reduction in the entire 4fn
shell. The successive contractions accumulate
and the total effect for all the lanthanum’s is known as the lanthanoid
contraction Lanthanoids contraction causes the radii of the members of the
third transition series to be very similar to those of the corresponding members
of the second series. The almost identical radii of Zr and Hf, a consequence of
the lanthanoid contraction, account for their occurrence together in nature and
for the difficulty faced in their separation
16. Oxidation States: Whatever the electronic configurations of lanthanons in the
ground state, all of them form ions in oxidation state +3. The ionization and
hydration energies are such that the tripositive state is more stable than the
di- or tetrapositive states in aqueous solutions
In the solid state too the combination of ionization and lattice energies is more
negative for the tripositive state than for the di- or tetrapositive states.
Consequently, the tripositive state is also the most common in the solid
compounds
In addition to the common tripositive state, some lanthanons form stable
compounds in +2 and +4 states
17. Magnetic Properties: Paramagnetic behaviour of an ion (or an atom) is
associated with its unpaired electrons. Thus all the lanthanoid ions other than f°
type (La3+
and Ce4+
) and f14
type (Yb2+
and Lu3+
) are paramagnetic. The
paramagnetism is maximum in neodymium
18. Colour: Most of the tripositive ions of lanthanons are coloured due to the
presence of f electrons. The absorption bands are narrow probably because of
the excitation within f level. La3+
and Lu3+
ions do not exhibit any colour

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19. Properties of Lanthanoids:
(i) In general, the earlier members of the series are quite reactive (similar to
calcium) but with increasing atomic number, they behave more like aluminium
(ii) Values for E^ for the half-reaction
Ln3+
(aq)+3e−
 Ln(s)
Are in the range of−2.2 to −2.4 V except for Ln=Eu for which the value is −2.0 V
(this being a small variation)
(iii) The metals combine with hydrogen when gently heated in the gas
(iv) (The carbides, Ln3C, Ln2C3 and LnC2 are formed when the metals are heated with
carbon
(v) They liberate hydrogen from dilute acids and burn in halogens to form halides
(vi) (They form oxides and hydroxides−M2O3 and M(OH)3. The hydroxides are
definite compounds, not just hydrated oxides, basic like alkaline earth metal
oxides and hydroxides Chemical reactions of the lanthanoids
20. Applications of Lanthanoids and their Compounds:
(i) Lanthanoid are used for making alloys steels
(ii) Mischmetall, an alloy containing nearly 95% lanthanoid metal is used to produce
bullets and shells
(iii) Mixed lanthanum oxides are used as phosphors in television screens
(iv) Compounds of lanthanoids are good catalysts, e.g., in cracking of petroleum
21. Actinoids. The 14 f-block elements of the seventh period following actinium (at
no. 89) are termed as actinoids. Actinoids are radioactive. Elements with atomic
numbers higher than that of uranium are called transuranium elements. All of
them have been prepared artificially
22. Electronic Configuration of Actinoids
(i) All the actinoids have the electronic configuration of 7s2
and variable occupancy
of the 5f and 6d subshells
(ii) The irregularities in the electronic configuration of the actinoids are related to
the stabilities of the f°, f7
and f14
occupancies of the 5f orbitals
23. Ionization enthalpies of actinoids are lower than lanthanoids. 5f orbitals of
actinoids penetrate less into the corresponding lanthanoids
Because the outer electrons are less firmly held, they are available for bonding
in the actinoids
24. Oxidation States. The actinoids show in general+3 oxidation states. Maximum
oxidation state is+7. The greater range of oxidation states is due to the fact
that the 5f, 6d and 7s levels are of comparable energies
25. Magnetic Properties. Although the variation in the magnetic susceptibility of the
actinoids with the number of unpaired 5f electrons is roughly parallel to the
corresponding results for the lanthanoids, the latter have higher values
26. Ionic Sizes. The ionic radius decreases regularly along the series. The decrease
in ionic radius is because of the poor screening of the nuclear charge by the f
electrons. This decrease being analogous to lanthanoid contraction is termed as

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actinoid contraction. Ionic radii of actinoids are very similar to those of
lanthanoids. Their chemical properties are, therefore, alike
27. General Properties:
The actinoids are reactive metals, especially when divided. Some important
points regarding their chemical reactivity are:
a) Actinoids react with boiling water to give a mixture of oxide and hydride
b) Actinoids combine with most non-metals at moderate temperatures
c) All actinoids are attacked by hydrochloric acid, but most are slightly affected
by nitric acid owing to the formation of protective oxide layer
d) Alkalies do not react with actinoids
28. Differences Between Lanthanoids and Actinoids:
Lanthanoids Actinoids
(i) Besides +3 oxidation states they
show +2 and +4 oxidation states
only in few cases
(ii) most of their ions are colourless
(iii) They have less tendency towards
complex formation
(iv) Lanthanoid compounds are less
basic
(v) Do not form oxocation
(vi) Except promethum, they are non
radioactive
(vii) Their magnetic properties can be
explained Easily
(i) Besides +3 oxidation states the
show higher oxidation states of
+4, +5, +6, +7 also
(ii) Most of their ions are coloured
(iii) They have greater tendency
towards complex formation
(iv) Actinoid compounds are more
basic.
(v) Form oxocations e.g. UO2+2,
PUO22+ and UO.
(vi) (They are radioactive
(vii) Their magnetic properties
cannot be explained easily

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EXERCISE
01. On what ground can you say that scandium (Z = 21) is a transition element but
zinc (Z = 30) is not?
02. Why do the transition elements exhibit higher enthalpies of atomisation?
03. Name a transition element which does not exhibit variable oxidation states.
04. Calculate the magnetic moment of a divalent ion in aqueous solution if its atomic
number is 25.
05. Which of the 3d series of the transition metals exhibits the largest number of
oxidation state and why?
06. Why is the highest oxidation state of a metal exhibited in its oxide of fluoride
only?
07. Calculate the ‘spin only’ magnetic moment of M2+
(aq) ion (Z = 27).
08. Explain why Cu+ ion is not stable in aqueous solution?
09. Write down the electronic configuration of :
(i) Cr3+
(ii) Cu+
(iii) Mn2+
10. Why are Mn2+
compounds more stable than Fe2+
towards oxidation to their + 3
state?
11. Explain giving reason:
(i) Transition metals and many of their compounds show paramagnetic
behaviour.
(ii) The transition metals generally form coloured compounds.
(iii) Transition metals in their many compounds act as good catalyst.
12. What are interstitial compounds? Why are such compounds well known for
transition metals?
13. How is the variability in oxidation states of transition metals different from
that of the non-transition metals? Illustrate with examples.
14. Predict which of the following will be coloured in aqueous solution?

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Page126
Ti3+
,V3+
, Cu+
, Sc3+
, Mn2+
, Fe3+
, Ca2+
and MnO4
−
. Give reasons for each.
15. Compare the stability of 2+
oxidation state for the elements of the first
transition series.
16. Name the oxometal anions of the first series of the transition metals in which
the metal exhibits the oxidation state equal to its group number.
17. What are the characteristics of the transition elements and why are they called
transition elements? Which of the d-block elements may not be regarded as the
transition elements?
18. What are the different oxidation states exhibited by the lanthanoids?
19. Describe the preparation of potassium permanganate. How does the acidified
permanganate solution react with (i) iron(II) ions (ii) SO2 and (iii) oxalic acid?
Write the ionic equations for the reactions.
20. Predict which of the following will be coloured in aqueous solution? Ti3+
, V3+
, Cu+
,
Sc3+
, Mn2+
, Fe3+
and Co2+
. Give reasons for each.
21. Name a transition metal which does not exhibit variation in oxidation state in its
compounds.
22. Explain why transition metals form alloys with other transition metals easily.
23. Explain hoe the colour of K2Cr2O7 solution depends on pH of the solution.
24. Compare the chemistry of actinoids with that of the lanthatnoids in reference
to
(i) Atomic and ionic sizes.
(ii) Oxidation states
25. Write the complete chemical equation for each of the following
(i) An alkaline solution of KMnO4 reacts with an iodide
(ii) An excess of SnCl2 solution is added to a solution of mercury (II)
Chloride.

8. d and f block elements 01

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