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InorgChem – d-Block Elements

Properties
TRANSITION ELEMENTS are elements, which forms at least one compound in which the element has in
incompletely filled inner d-subshell
 Elements with at least one oxidation state in which the element has incompletely filled
inner d-subshell (i.e. w/1 to 9 e–)
 Scandium (Sc) & zinc (Zn) are not transition elements as both has only one oxidation
state Sc3+ & Zn2+, which both has no incompletely filled d-subshell
d-BLOCK ELEMENTS are elements whose highest-energy electron / last-filled electron is found in the d-
orbital
 Element atoms with electronic configuration [ ](n+1)s2nd x
 First transition series: Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn

Electronic configuration:
 Atoms of d-block elements are characterized by inner building of d-subshells
 Most contains the configuration [ ]nd x(n+1)s2
 Except: Cr (Z=24): [Ar]3d54s1 and
Cu (Z=29): [Ar]3d104s1
 The 2+ ions of d-block elements have no e– in the valence s-orbital
 Configuration: [ ]nd x

Typical properties of Transition elements:

 Properties similar to s-block elements
 Metals
 Good conductor of electricity & heat
 MP ≪ BP
 Properties different from s-block elements
 Hard, strong, high : Strong & closely packed struct. due to delocalization of 4s & 3d e–
 Similar chemical properties: Outermost subshell = 4s2
 Variable oxidation state: Successive IE are close
 Ability of formation of complexes: Availability of low-lying d-vacant orbitals
 Cpds are colored: Splitting of d-subshell into different energy levels after complex formation
 Catalytic behavior: Ability to form complexes & the existence of variable oxidation states
 Sc & Zn do not exhibit typical properties of transition elements

Variation of Properties
Ionization Enthalpy:
 d-Block elements has 4d subshell filled before 3d subshell
 Before filling, 4s has lower energy than 3d
 After filling d-orbital, 3d has lower energy than 4s
 e– in d-orbital screens 4s as it is more penetrating
 When d-block metal ionizes, the 4s e– are first removed
 Difference in energy btw 3d & 4s subshells is small, successive IE increases gradually
 Across the 1st transition series, IE increases slightly
 Additional e– is added to inner 3d subshell and screens the 4s e–, which is being ionized
 Screening effect cancels most the increase in ENC
 1st IE of Sc ~ Zn  1st IE of K & Ca as the ENC inc.
 Due to the small & gradual increase in IE, transition elements have similar chemical properties
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Metallic Radius:
 Factors affect the metallic radius of d-block metals:
 ENC on the outermost e–: inc. from Sc to Zn
 Repulsion btw the d electrons: inc. from Sc to Zn
 All d-block metals have smaller atomic radii than s-block in same period due to the inc. of ENC
 Across the period from Sc to Zn,
 Atomic radius dec. slightly from Sc to Cr: Slight increase of ENC
 Cr to Ni have nearly constant atomic radius: Increasing ENC & repulsion cancelled out
 Atomic radius inc from Ni to Zn: Repulsion btw e– increases as Z increases
 Overall variation in atomic radius is small (0.12 – 0.16 nm): e– is in inner 3d subshell

Ionic Radius:
 Variation of ionic radius of d-block elements is similar to that of atomic radius
 All d-block metals have smaller ionic radii than s-block due to the inc. of ENC
 Across the period, the ionic radius from Sc to Cu dec. slightly as the slightly inc. of ENC
 Ionic radius of Zn2+ is slightly larger than expected as the full-filled 3d10 produces a larger
electronic repulsion

Melting Point & Hardness:

 MP & Hardness of metal  Metallic bond strength
 d-Block elements usually have higher MP & hardness
 Both s & d e– are delocalized to the electron sea and take part in the metallic bond
 Strong metallic bond
 Most d-block elements are closely packed (Coordination number = 12)
 Atomic radii of d-block elements are smaller
 MP ≪ BP as only small fraction of metallic bonds has to be broken in melting but all metallic
bonds are broken in boiling (i.e. Boiling = Atomization)
 Factor affecting the MP & hardness of Sc to Zn
 Smaller atomic radius  Stronger metallic bond
 More unpaired d e–  More delocalized e– in metallic bond  Stronger metallic bond
 Closer packing of atoms  Stronger metallic bond
 Half-filled or full-filled d-subshell are extra stable
 Across the period from Sc to Zn, the MP & hardness
 Increased from Sc to V as the no. of unpaired d e– inc. & atomic radius dec.
 Metallic bond strength increases
 Decreased from V to Mn as Cr & Mn has extra stable 3d5  Metallic bond strength dec.
 Increased to a max. at Fe, Co, Ni as unpaired d e– are more involved in metallic bond
 Decreased from Cu to Zn as both has extra stable 3d10

Density:
 d-Block elements usually have higher densities than s-block elements as the atoms are closely
packed
 The density increases from Sc to Cu as the relative atomic mass increases while atomic radii are
approximately constant
 Zn has a lower density than Cu as the 3d & 4s e– are less involved in metallic bond
 Larger atomic radius

Electronegativity:
 Across the period from Sc to Zn, ENC increases
 Electronegativity generally increases
 Ease of formation of metal ions dec.
 Metallic character dec.
 d-Block elements are less electropositive / more electronegative than s-block elements
 d-Block cations are less likely to form than s-block
 Mn has a lower electronegativity than expected
 Due to low ionization enthalpies: Extra stability of Mn2+: [Ar]3d5
 Zn has particularly low electronegativity
 Due to low ionization enthalpies: Extra stability of Zn2+: [Ar]3d10
 The standard electrode potential (E ) has a similar variation with electronegativity
 When ENC on the outermost e– increases, E becomes more positive or less negative
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Oxidation States:
 Transition metals show variable oxidation states:
Sc +1, +2, +3 Sc3+
Ti +1, +2, +3, +4 Ti2+, Ti3+, TiO2
V +1, +2, +3, +4, +5 V2+, V3+, VO2+, VO2+
Cr +1, +2, +3, +4, +5, +6 Cr2+, Cr3+, Cr4+, CrO3
Mn +1, +2, +3, +4, +5, +6, +7 Mn2+, Mn(OH)3, MnO2, MnO42–, MnO4–
Fe +1, +2, +3, +4, +5, +6 Fe2+, Fe3+
Co +1, +2, +3, +4, +5 Co2+, Co3+, [CoF6]2–
Ni +1, +2, +3, +4 Ni2+, Ni3+, Ni2O3, NiO2
Cu +1, +2, +3 Cu+, Cu2+
Zn +1, +2 Zn2+
* Complex only Most stable state Oxidation states
 The existence of variable oxidation states due to gradual increases rather than sharp rapid rise of
successive IE of transition metal
 Across the 1st transition series,
 From Sc to Mn, the highest possible ON = Total no. of e– in 3d & 4s
 Maximum ON found in Mn(VII)
 By removal of all 3d & 4s electrons from Mn
 From Mn onwards, the highest ON decreases as the no. of unpaired e– decreases
 Stability of a particular ON is related to extra stable state of half/full-filled d-orbital
 Mn2+ & Fe3+ are particularly stable as they have [Ar]3d5 config.
 Zn2+ is particularly stable as it has [Ar]3d10 config.
 Enthalpy of formation of an ion: Hf = Hatom + IE + Hhyd
 An ion of lower Hf is energetically stable
 An energetically stable ion (e.g. Cu+) may be unstable w.r.t. disproportionation:
 Cu(s)  Cu+(aq) Hf = +606 kJmol–1
Cu(s)  Cu2+(aq) Hf = +971 kJmol–1 (less stable)
 Disproportionation: 2Cu+(aq)  Cu(s) + Cu2+(aq) H = –241 kJmol–1
 Occurs spontaneously & energetically favorable

Formation of Colored Cpds:

 Most s-block metals form colorless cpds but most d-block metals form colored cpds
 Exception: Sc3+, Ti4+ Cu+, Zn2+
 When transition metal forms cpd,
 The electrons in s shell is removed
 d shell becomes the outermost shell
 The d-orbitals may split into different energy levels in cpd
 The splitting is due to the electronic repulsion btw d e– and LP e– of anion / ligand
 If the metal ion has d shell neither full-filled nor empty, e– in lower energy level can jump to a
higher orbital
 Absorbs electromagnetic radiation of a particular frequency
 Absorption in the visible light region makes the ion colored
 Color of aqueous ion:
Sc3+ Colorless
Ti3+ Purple Ti4+ Colorless
2+
V Violet V3+ Green VO2+ Blue VO2+ Yellow VO3– Red
Cr2+ Blue Cr3+ Green
Mn Pale pink Mn(OH)3 Brown MnO2 Black MnO42– Green MnO4– Purple
2+

Fe2+ Green Fe3+ Yellow

Co2+ Pink
Ni2+ Green
Cu+ Green Cu2+ Blue
2+
Zn Colorless
 Some transition metals show colored flame when burned as the excitation of e–
 Radiation in visible spectrum is emitted when the e– falls back

Reaction with water:

 d-Block metals usually react slowly with water
 The low reactivity of d-block metals is due to its high IE & Hatom
 Exception: Sc + hot water  Sc(OH)3(aq) + H2(g)
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Vanadium & Manganese

Vanadium:
 Vanadium is a soft & silvery-white metal
 Common oxidation states: +2, +3, +4, +5
 Yellow VO2+, Blue VO2+, Green V3+, Violet V2+ respectively
 Vanadium can react w/oxidizing acids HNO3 & H2SO4
 Vanadium(V):
 Dioxovanadium(V) ion VO2+ is produced by dissolving ammonium polytrioxovanadate(V):
NH4VO3(aq) + 2H+(aq) VO2+(aq) + NH4+(aq) + H2O()
– +
VO3 (aq) + 2H (aq) VO2+(aq) + H2O()
 Neither V5+ nor [V(H2O)6]5+ exists as the charge density is extremely high
 Hydrolysis occurs: V5+(aq) + 2H2O(aq)  VO2+(aq) + 4H+(aq)
 It is a stable oxidation state of V
 Vanadium(V) oxide, V2O5, is an amphoteric oxide:
V2O5(s) + 2H+(aq) 2VO2+(aq) + H2O()
–
V2O5(s) + 6OH (aq) 2VO43–(aq) + 3H2O()
 Vanadium(IV):
 Oxovanadium(IV) ion VO2+ can be produced by reacting VO2+ with metal:
Zn(s) Zn2+(aq) + 2e–
+ + –
VO2 (aq) + 2H (aq) + e VO2+(aq) + H2O()
 Neither V4+ nor [V(H2O)6]4+ exists as the charge density is extremely high
 Vanadium(III) & Vanadium(II):
 V3+ and V2+ ions can be produced by reducing VO2+ by metal:
VO2+(aq) + 2H+(aq) + e– V3+(aq) + H2O()
V3+(aq) + e– V2+(aq)
 By adding excess Zn granules or Zn amalgam in acidic condition (conc. HCl),
VO3–(aq)  VO2+(aq)  VO2+(aq)  V3+(aq)  V2+(aq)
 Both V3+ and V2+ ions are reducing and readily oxidized by air or water

Manganese:
 Manganese is a hard, pinkish-grey metal
 Chemically active and rapidly attacked by hot water, steam or acid:
Mn(s) + 2H2O() Mn(OH)2(s) + H2(g)
Mn(s) + H2O(g) MnO(s) + H2(g)
+
Mn(s) + 2H (aq) Mn2+(aq) + H2(g)
 Common oxidation state: +2, +4, +7
 Manganese(II):
 The pale pink Mn2+ ion is very stable in acids (partly due to 3d5 config.)
 It forms complex [Mn(H2O)6]2+ in water and undergoes slight hydrolysis
[Mn(H2O)6]2+(aq) [Mn(OH)(H2O)5]+(aq) + H+(aq)
 In alkaline medium, Mn would becomes white Mn(OH)2:
2+

[Mn(H2O)6]2+(aq) + 2OH–(aq) [Mn(OH)2(H2O)5](s) + 2H2O()

 The Mn(OH)2 in alkaline is readily oxidized by O2, H2O2, OCl– to unstable Mn(OH)3(s)
and finally to MnO2(s):
4Mn(OH)2(s) + O2(g) + 2H2O()  4Mn(OH)3(s)
2Mn(OH)3(s) + (n–3)H2O(g)  Mn2O3nH2O(s)
2Mn2O3nH2O(s) + O2(g)  4MnO2(s) + (2n)H2O()
 2Mn(OH)2(s) + O2(g)  2MnO2(s) + 2H2O()
 Manganese(IV):
 MnO2 is a dark brown amphoteric solid w/strong oxidizing property:
MnO2(s) + 4H+(aq) + 2e– Mn2+(aq) + 2H2O()
 MnO4 can oxidize Cl in HCl(aq):
–

MnO2(s) + 4HCl(aq)  MnCl2(aq) + 2H2O() + Cl2(g)

 MnO4 can be oxidized by strong oxidizing agent in basic medium:
3MnO2(s) + 6OH–(aq) + ClO3–(aq)  3MnO42–(aq) + 3H2O() + Cl–(aq)
2MnO2(s) + 4OH–(aq) + O2(g)  2MnO42–(aq) + 2H2O()
 Manganese(VII):
 Mn2O7 is acidic covalent compound
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 Purple manganate(VII) ion MnO4– is a powerful oxidizing agent:

Acidic medium: MnO4–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O()
Alkaline medium: MnO4–(aq) + 2H2O() + 3e– MnO2(s) + 4OH–(aq)
 MnO4– can react w/iron(III) ion:
MnO4–(aq) + 8H+(aq) + 5Fe2+(aq)  Mn2+(aq) + 4H2O() + 5Fe3+(aq)

Complexes
Complex:
 A complex (coordination compound) is formed when a central metal ion/atom is attached by
ligands in which the no. of ligands is greater than the normal valance of the central ion/atom
 The metal ion/atom acts as an electron pair acceptor (Lewis acid or electrophile)
 The metal ion/atom ought to be has high positive charge density to attract the ligands
 The ligand acts as an electron pair donor (Lewis base or nucleophile)
 COMPLEX: A central metal ion/atom surrounded by and datively bonded to other ligands
 LIGAND: An ion or molecule containing at least one atom having a lone pair of electrons which
can be donated to the central cation or atom to form a dative covalent bond
 COORDINATION NUMBER of the central ion is the number of ligands bonded to the central ion
 d-Block metal has strong tendency towards complex formation as the availability of low-lying
vacant orbitals enables metal ions to accept LP e– from ligands

Nomenclature:
 Coordination compounds are named according to the nomenclature recommended by IUPAC
 Ionic coordination compounds:
 Cation is name before the anion
 The ligands and central metal are named together as one word
 Order of ligands: Anionic, Neutral, Cationic
 Example: [PtCl2(NH3)4]2+ = Dichlorotetraammineplatium(IV)
 Names of anionic ligands end in –o
 Names of neutral ligands are the names of the molecules, expt NH3, H2O, CO, NO
 Names of common ligands:
LIGAND PREFIX LIGAND PREFIX
–
Bromide Br Bromo Ammonia NH3 Ammine
Chloride Cl– Chloro Water H2O Aqua
–
Cyanide CN Cyano Carbon monoxide CO Carbonyl
Fluoride F– Fluoro
–
Hydroxide OH Hydroxo
Sulphate(VI) SO42– Sulphato
Amide NH2– Amido
 If the no. of a particular ligand is more than one, the number is indicated w/Greek prefix:
 Within each type of ligand, the ligands are arranged in alphabetical order, ignoring the
numbering prefixes
NUMBER PREFIX NUMBER PREFIX
Two Di- Five Penta-
Three Tri- Six Hexa-
Four Tetra-
 If the complex is anionic, the name of the metal ends in –ate
 Example: [Fe(CN)6]3– = Hexacyanoferrate(III)
 Names of common metals in anionic complexes:
METAL NAME METAL NAME
Titanium Titanate Nickel Nickelate
Chromium Chromate Copper Cuprate
Manganese Manganate Zinc Zincate
Iron Ferrate Platinum Platinate
Cobalt Cobaltate
 If the complex is cationic or neutral, the name of the metal is unchanged
 Example: [CrCl2(H2O)4]+ = Dichlorotetraaquachromium(III)
 Neutral coordination compounds:
 The name of the complex is the name of the cpd
Bonding & Stability:
 Ligand contains LP e– which forms dative bonds with the central metal
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 Different ligand has different tendency to donate their unshared electrons to form dative bond
 The strength of the dative covalent bond of differs as the ligands differ
 A ligand can be replaced by another ligand which form more strong dative bond
 i.e. the ligand that of higher tendency to donate the electrons
 When a stronger ligands present, the weaker ligand would be replaced:
[Fe(H2O)6]2+(aq) + 6CN–(aq)  [Fe(CN)6]4–(aq) + 6H2O()
[Ni(H2O)6]2+(aq) + 6NH3(aq)  [Ni(NH3)6]2+(aq) + 6H2O()
 In the displacement of ligands of the complex, usually there would be color change
 The formation of complex or displacement of ligand is a stepwise reaction
 In [Cu(H2O)6]2+(aq) + 4NH3(aq) [Cu(NH3)4(H2O)2]2+(aq) + 4H2O():
[Cu(H2O)6]2+(aq) + NH3(aq) [Cu(NH3)(H2O)5]2+(aq) + H2O() K1 = 1.41104mol–1dm3
[Cu(NH3)(H2O)5]2+(aq) + NH3(aq) [Cu(NH3)2(H2O)4]2+(aq) + H2O() K2 = 3.16103mol–1dm3
[Cu(NH3)2(H2O)4]2+(aq) + NH3(aq) [Cu(NH3)3(H2O)3]2+(aq) + H2O() K3 = 7.77102mol–1dm3
[Cu(NH3)3(H2O)3]2+(aq) + NH3(aq) [Cu(NH3)4(H2O)2]2+(aq) + H2O() K4 = 1.35102mol–1dm3
 K1, K2, K3, K4 are called stepwise stability constant
 The equilibrium constant of overall reaction:

 K is called the stability constant of the complex and it gives a measure of the stability
of the complex ion
 The complexes with large stability constant is stable

Stereostructure:
 The spatial arrangement of ligands around the central metal is related to the coordination number
 The coord. no. is determined by: Size of the central metal
No. and nature of vacant orbitals for forming dative bonds
 Generally the coord. no. is constant for a particular metal atom / ion
 Common coordination numbers: 2, 4, 6

 Coordination number = 6:
 In complexes with coord. no. of 6, the structure usually is octahedral:
 Example: [Cr(NH3)6]3+ and [Fe(CN)6]3–

 Coordination number = 4
 In complexes with coord. no. of 4, the structure usually is tetrahedral:
 Example: [Zn(NH3)4]2+ and [CoCl4]2–
 A few four-coordinated complexes are having square planar structure:
 Example: [Cu(NH3)4]2+ and [CuCl4]2–
 The [Cu(NH3)4]2+ and [CuCl4]2– complex should be six-coordinated and in octahedral
structure, but the two H2O is loosely bonded and ignored usually

 Isomerism occurs in the complexes:

 Structural isomers: Different ligands coordinated to the central metal
 Example: [Co(NH3)5Br]2+SO42– and [Co(NH3)5SO4]+Br–
 Geometrical isomer: Different geometrical arrangement of ligands
 Square planar complexes:

 [Ma2b2]: and

 [Ma2bc]: and

 [Mabcd]: , and
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 Octahedral complexes:

 [Ma4b2]: and

 [Ma3b3]: and

Catalysis
Catalytic action:
 d-Block elements and their cpds are common catalysts in industry & biological systems
 By providing suitable reaction surfaces, or
 By forming unstable intermediates
 Common catalysts:
 V2O5 or VO3–: Contact process, 2SO2(g) + O2(g) 3SO3(g)
 Fe or Fe2O3: Haber process, N2(g) + 3H2(g) 2NH3(g)
 MnO2: Decomposition of hydrogen peroxide, 2H2O2()  2H2O() + O2(g)
 Ni: Manufacture of margarine, RCH=CH2 + 2H2(g)  RCH2CH3
 Pt: Manufacture of nitric(V) acid, 4NH3(g) + 5O2(g)  4NO(g) + 6H2O()

Heterogeneous catalysis:
 In heterogeneous catalysis, the catalyst and reactants are in different phases
 Usually, the catalyst are finely divided solids
 The heterogeneous catalyst provides a suitable reaction surface for reactants to come close
together and react
 Example, catalysis of gaseous reaction on solid surfaces like Haber process
 Outlined mechanism:
 The gaseous reactant diffuse to the catalyst
 The catalytic surface adsorbs one reactant (Cat + R1  Cat–R1)
 The catalytic surface adsorbs another reactant (Cat + R2  Cat–R2)
 Bond breaking & forming btw catalyst & reactants (Cat–R1 + Cat–R2  Cat–R2–R1–Cat)
 The products desorbed and diffused away (Cat–R2–R1–Cat  R1–R2)
 The abundance of valence e– and availability of vacant orbitals in d-block elements facilitate the
surface adsorption of reactants
 The adsorption brought the reactants within close proximity to catalyze reaction

Homogeneous catalysis:
 In homogeneous catalysis, the catalyst and reactants are in same phase
 The catalyst form an intermediate with the reactants to catalyze reaction
 Reaction mechanism is changed such that EA is lower
 Example: S2O82–(aq) + 2I–(aq)  2SO42–(aq) + I2(aq) is slow due to kinetic factors
 Fe3+ can catalyze the reaction as Fe2+(aq) + e– Fe2+(aq):
2I–(aq) + 2Fe3+(aq)  I2(aq) + 2Fe2+(aq)
2Fe2+(aq) + S2O82–(aq)  2SO42–(aq) + 2Fe3+(aq)
 Fe oxidizes iodide ions and gives Fe2+, but Fe2+ reduces peroxodisulphate(VI) and gives Fe3+
3+

 The catalyst is unchanged at last