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d-BLOCK ELEMENTS
1. I NT RO DU CT IO N
(a) The element lying between s- and p-block element of the periodic table are collectively known as transition or transitional elements. (T.E'.S.) (b) Their properties are transitional between the highly electropositive s- block element to least electropositive p-block element. (c) In d- block elements, the last differentiating electron is accommodated to the penultimate shell. (d) The general electronic configuration of transition element is (n-1)d 1-10^ ns 0, 1 or 2 (e) These elements either in their atomic state or in any of their common oxidation state have partly filled (n-1)d orbitals of (n-1) th^ main shell. (f) The transition elements have an incompletely filled d-level. Since Zn, Cd, Hg elements have d 10 configuration and are not considered as transition elements but they are d-block elements. ELECTRONIC CONFIGUR ATION I st^ Transat ion Serie s Symbol Sc Ti V C r Mn Fe Co Ni C u Z n Atomic No. 2 1 2 2 2 3 2 4 2 5 2 6 2 7 2 8 2 9 3 0 3d electrons 1 2 3 5 5 6 7 8 1 0 1 0 4s electrons 2 2 2 1 2 2 2 2 1 2 Ir regular electronic configurat ion Cr, Cu II nd^ Transat ion Serie s Symbol Y Z r N b M o Tc R u R h P d A g Cd Atomic No. 3 9 4 0 4 1 4 2 4 3 4 4 4 5 4 6 4 7 4 8 4d electrons 1 2 4 5 5 7 8 1 0 1 0 1 0 5s electrons 2 2 1 1 2 1 1 0 1 2 Irregular electronic configuration Nb, Mo, Ru, Rh, Pd, Ag III r d^ Transat ion Serie s Symbol L a Hf Ta W R e Os Ir P t A u Hg Atomic No. 5 7 7 2 7 3 7 4 7 5 7 6 7 7 7 8 7 9 8 0 5d electrons 1 2 3 4 5 6 7 9 1 0 1 0 6s electrons 2 2 2 2 2 2 2 1 1 2 Ir regular electronic configurat ion W, Pt, Au
The irregularities in the observed configuration of Cr (3d 5 4s 1 instead of 3d 4 4s 2 ), Cu (3d 10 4s 1 ), Mo (4d 5
5s^1 ), Pd ([Kr] 4d^10 5s^0 ), Au ( [Xe] 4f^14 5d^10 6s^1 ), Ag ([Kr] 4d^10 5s^1 ) are explained on the basis of the concept that half-filled and completely filled d-orbitals are relatively more stable than other d-orbitals.
- GENERAL PROPERTIES OF d-BLOCK ELEMENTS (a) The properties of d-block elements of any given period are not so much different from one another as those of the same period of non transtion elements.
(b) It is due to the fact that, in transition series, there is no change in number of electrons of outermost shell and only change occur in (n-1)d electron from member to member in a period.
- METALLIC CHAR ACTER
(a) All the d-block elements are metals as the numbers of electrons in the outer most shell are one or two. (b) They are hard, malleable and ductile (except Hg). IB group elements Cu, Ag and Au are most ductile and soft. (c) These are good conducter of heat and electricity (due to free e—^ ) Elements of IB group are most conductive in nature. Their order of conductivity is Ag > Cu > Au > Al (d) Covalent and metallic bonding both exist in the atom of transition metals. (e) The presence of partially filled d-subshell favour covalent bonding and metallic bonding. These bonding are favourable also due to possession of one or two electron in outermost energy shell.
REDUCING POWER (a) Reducing power of d-block elements depends on their electrode potential. (b) Standard oxidation potential (SOP) of Cu is minimum in the 3d series so it is least reducing elements in 3d series. (c) Au is the least reducing element in the d-block because of highest +ve value of Standard reduction potential. (d) The poor reducing capacity of the transition metal is due to high heats of vaporization, high ionization potential and low heat of hydration of their ions, because reduction potential depends upon all these three factors.
D E N S I T Y
(a) The atomic volume of the transition elements are low, compared with s-block, so their density is comparatively high (D = M/V) (b) Os (22.57 gm cm —3^ ) and Ir (22.61 gm cm —3^ ) have highest density. (c) In all the groups (except IIIB) there is normal increase in density from 3d to 4d series, and from 4d to 5d, it increases just double. Due to lanthanide contraction Ex. Ti < Zr << Hf (d) In 3d series Cu Zn Density increases (^) Density decreases
Sc Ti V Cr Mn Fe Co Ni
(e) In 3d series highest density – Cu lowest density – Sc (f) Some important orders of density Fe < Ni < Cu Fe < Cu < Au Fe < Hg < Au
- MELTING AND BOILING POINTS (a) Melting and boiling point of d-block > s-block Reason : Stronger metallic bond and presence of covalent bond formed by unpaired d-electrons.) (b) In Zn, Cd, and Hg there is no unpaired electron present in d-orbital, hence due to absence of covalent bond melting and boiling point are very low in series. (Volatile metals Zn, Cd, Hg) (c) In 3d series Sc to Cr melting and boiling point increases then Mn to Zn melting and boiling point decreases (d) As the number of d-electron increases, the number of covalent bond between the atoms are expected to increase up to Cr-Mo-W family where each of the d-orbital has only unpaired electrons and the opportunity for covalent sharing is greatest. (e) Mn and Tc have comparatively low melting point, due to weak metallic bond because of stable Half filled (d 5 ) configuration
(c) Highest oxidation state of transition elements can be calculated by n + 2 where (n = number of unpaired electrons) It is not applied for Cr and Cu.
(d) The transition metal ions having stable configuration like d 0 d 5 or d 10 are more stable.
E x. Sc +3^ , Ti +4^ , V +5, Fe +3^ , Mn +2^ , Zn +2^ etc.
(e) In aqueous medium Cr +3^ is stable.
(f) Co +3^ and Ni +2^ are stable in complexes..
(g) In aqeous medium due to disproportionation Cu +1^ is less stable than Cu +2^ while its configuration is 3d 10
(h) Most common oxidation state among the transition elements is +2.
(i) Highest oxidation state shown by transition elements of '4d' and '5d' series is +8 by Ru (44) and Os (76).
(j) The common oxidation state shown by elements of IIIB i.e., Sc, Y, La and Ac is +3 as their divalent compounds are highly unstable.
(k) In lower oxidation state transition elements form ionic compounds and in higher oxidation state their compounds are covalent.
(l) They also shows zero oxidation state in their carbonyl compounds like Ni(CO) 4.
(m) Usually transition metal ions in their lower oxidation state act as reducing agents and in higher oxidation state they are oxidising agents.
E x. Sc +2^ , Ti +2^ , V +2^ , Fe +2^ , Co +2^ etc are reducing agents
Cr +6^ , Mn +7^ , Mn +6^ , Mn +5^ , Mn +4^ etc are oxidising agents.
The relative stability of various oxidation states
(a) The relative stabilities of various oxidation states of 3d-series element can be correlated with the extra stability of 3d°,3d 5 & 3 d 10 configuration to some extent.
E x. Stability of Ti 4+^ (3d 0 ) > Ti 3 +^ (3d 1 ) Mn 2+^ (3d 5 ) > Mn 3+^ (3d 4 )
(b) The higher oxidation state of 4d and 5d series element are generally more stable than the elements of 3d series. Ex.
(i) Mo vi^ O -2 4 (oxidation state of Mo is +6), Mo vi^ O -2 4 (4d series) & W Ovi^24 ^ , Re viiO 4 ^ (5d series) are more
stable due to their maximum oxidation state.
(ii) (^) Cr Ovi^4 ^2 & (^) Mn viiO 4 ^ (3d-series) are strong oxidizing agents.
(c) Strongly reducing states probably do not form fluorides or oxides, but may well form the heavier halides Conversely, strong oxidizing state form oxides & fluoride, but not Bromide and lodide. Ex.
(i) V (Vanadium) react with halogens to form VF 5 VCl 5 , VBr 3 ,but doesn' t form VBr 5 or VI 5 because in + 5 oxidation state Vanadium is strong oxidizing agent thus convert Br –^ & I–^ to Br 2 & I 2 respectively, So VBr 3 &VI 3 are formed but not VBr 5 & VI (^) 5.
(ii) On the other hand VF 5 is formed because V 5+^ ion unable to oxidize highly electronegative & small anion F–
(iii) Similarly highly electronegative and small O 2 –^ ion formed oxides Ex. VO 4 3 –^ , CrO 4 2–^ & MnO 4 –^ etc.
Diffrent oxidation state of chloride & oxides compound
TiCl 2 TiCl 3 VCl 2 VCl 3 (Ionic, basic) Less ionic (Amphoteric)
TiCl 4 VCl 4 VOCl 3 Covalent and Acidic (Strong lewis acid)
TiO VO CrO MnO
Ti O 2 3 V O 2 3 Cr O 2 3 Mn O 2 3
TiO 2
MnO 2
V O 2 5
CrO 3 MnO 3 Mn O 2 7 Ionic, basic Less Ionic (Amphoteric) Acidic, covalent
(d) Such compounds are expected to be unstable except in case where vacant d-orbitals are used for accepting lone-pair from -bonding ligand. E x. [Ni(CO) 4 ], [Ag(CN) 2 ] –^ ,[Ag)(NH 3 ) 2 ] +
- COLOUR PROPERTY (a) Most of the transition metal ions exhibit colour property. (b) This is due to d-d transition of unpaired electrons in their t (^) 2g and e (^) g sets of 'd' orbitals. (c) They require less amount of energy to undergo excitation of electrons. Hence they absorb visible region of light exhibiting colour. E x. Sc +2^ : [Ar]3d 1 , Ti +2^ : [Ar]3d 2 , V +2^ : [Ar]3d 3 (d) Transition metal ions which do not have any unpaired elctrons in their 'd' orbitals like 3d 0 and 3d 10 configurations, do not exhibit any colour property. E x. Sc +3^ : [Ar]3d 0 , Cu +1^ : [Ar]3d 10 , Ti +4^ : [Ar]3d 0 etc are colourless ions. (e) A transition metal ion absorbs a part of visible region of light and emmits rest of the colours, the combination of which, is the colour of emitted light. The colour of metal ion is the colour of the emitted light. (f) In transition metal ion the 'd' orbitals split into lower energy set t2g orbitals and higher energy set eg orbitals. The electrons from t2g set get excited to higher energy set i.e., eg set. This excitation of electrons is called as 'd-d' transition. Due to this 'd -d' transition the transition metal ions exhibit colour property.
dx (^2) - (^) y 2 dz 2
dxy dyz dyz
d-orbitals (degenerate)
light
Lower energy set = t2g Higher energy set = eg Presence of ligands (^) d-d Trnsition
(t )2g
(e )g
Factors affecting the colour of complex The colour of a transition metal complex depends on- (a) The magnitude of energy difference between the two d-levels ( 0 ) (^) ,
1 0. CATALY TIC PROPERT Y
(a) Transition elements and their compounds exhibit catalytic properties. This is due to their variable valency as well as due to the free valencies on their surface. (b) When transition elements and their compounds are in powdered state, their catalytic properties exhibited will be to a greater extent. This is due to greater surface area available in the powdered state. (c) Transition metals and their compounds exhibiting catalytic properties in various processes are-
C a t a l y s t U s e d
TiCl 3 Used as the Ziegler-Natta catalyst in the production of polythene. V 2 O 5 Convert SO 2 to SO 3 in the contact process for making H 2 SO (^4) MnO 2 Used as a catalyst to decompose KClO 3 to give O (^2) Fe Promoted iron is used in the Haber-Bosch process for making NH (^3) FeCl 3 Used in the production of CCl 4 from CS 2 and Cl (^2) FeSO 4 and H 2 O 2 Used as Fenton's reagent for oxidizing alcohols to aldehydes. PdCl 2 Wacker process for converting C 2 H 4 + H 2 O + PdCl 2 to CH 3 CHO + 2HCl + Pd. Pd Used for hydrogenation (e.g. phenol to cyclohexanone). Pt/PtO Adams catalyst, used for reductions. Pt Formerly used for SO 2 SO 3 in in the contace process for making H 2 SO (^4) Pt/Rh Formerly used in the ostwald process for making HNO 3 to oxidize NH 3 to NO Cu Is used in the direct process for manufacture of (CH 3 ) 2 SiCl 2 used to make silicones. Cu/V Oxidation of cyclohexanol/cyclohexanone mixture to adipic acid which is used to make nylone- CuCl 2 Decon process of making Cl 2 from HCl Ni R ane y ni ck e l, nu m er o u s r ed u ct i on p ro c es s es ( e. g. m anu fac t ur e o f hexamethylenediamine, productiomn of H 2 from NH 3 , reducing anthraquinone to anthraquinol in the production of H 2 O (^2)
1 1. FORM ATION OF ALLOY
(a) Transition elements have maximum tendency to form alloys. (b) The reactivity of transition elements is very less and their sizes are almost similar. Due to this a transition metal atom in the lattice can be easily replaced by other transition metal atom and hence they have maximum tendency to form alloys. (c) In the alloys, ratio of component metals is fixed. (d) These are extremly hard and have high melting point.
SOME IMPORTANT ALLOY
(a) Bronze Cu (75 - 90 %) +Sn ( 10 - 25 %) (b) Brass Cu ( 60 - 80 %) +Zn (20 - 40 %) (c) Gun metal (Cu + Zn + Sn) (87 : 3 : 10) (d) German Silver Cu + Zn + Ni ( 2 : 1 : 1) (e) Bell metal Cu (80 %) + Sn (20 %) (f) Nichrome (Ni + Cr + Fe) (g) Alnico (Al, Ni, Co) (h) Type Metal Pb + Sn + Sb (i) Alloys of steel Vanadium steel V (0.2 - 1 %) Chromium steel Cr (2 - 4 %) Nickel steel Ni (3 -5 %) Manganese steel Mn (10 - 18 %) Stainless steel Cr (12 - 14 %) & Ni (2 - 4 %) Tunguston steel W (10 - 20 %) Invar Ni (36 %) (j) 14 Carat Gold 54 % Au + Ag (14 to 30 %) + Cu (12 - 28 %) (k) 24 Carat Gold 100 % Au (l) Solder Pb + Sn (m) Magnellium Mg (10%) + Al (90%) (n) Duralumin (Al + Mn + Cu) (o) Artificial Gold Cu (90 %) + Al (10%) (p) Constantan Cu(60%) + Ni (40%)
% of Carbon in different type of Iron N a m e % of C (a) Wrought Iron 0.1 to 0. (b) Steel 0.25 to 2. (c) Cast Iron 2.6 to 4. (d) Pig Iron 2.3 to 4.
1 2. FORM ATION OF INTERSTITIAL COMPOUNDS
(a) Transition elements form interstitial compounds with smaller sized non metal elements like hydrogen, carbon, boron, nitrogen etc. (b) The smaller sized atoms get entrapped in between the interstitial spaces of the metal lattices. These interstitial compounds are nonstoichiometric in nature and hence cannot be given any definite formula. (c) The smaller sized elements are held in interstitial spaces of transition elements by weak Vander Waals forces of attractions. (d) The interstitial compounds have essentially the same chemical properties as the parent metals but they differ in physical properties such as density and hardness. The process of adsorption of excess of H atom by the transition metals like Pd, Pt etc is called occlusion.
(d) Action of Alkalies :- On heating with alkalies the orange colour of dichromate solution changes to yellow due to the formation of chromate ions.
K 2 Cr 2 O 7 + 2KOH 2K 2 CrO 4 + H 2 O
or Cr O 2 72 ^ 2 OH ^ 2 CrO 42 H O 2
This chromate on acidifying reconverts into dichromate.
2K 2 CrO 4 + H 2 SO 4 K 2 Cr 2 O 7 + K 2 SO 4 + H 2 O
or 2 CrO 42 ^ 2 H ^ Cr O 2 72 H O 2
The interconversion is explained by the fact that dichromate ion and chromate ion exist in equilibrium at a pH of about 4.
Cr O 2 72 ^ H O 2 2 CrO 42 ^ 2 H
When alkali added, H +^ consumed so forward direction. When acid added, H +^ increases so backward direction. (e) Chromyl chloride Test :- When potassium dichromate is heated with conc. H 2 SO 4 acid and a soluble metal chloride (ex. NaCl) orange red vapours of chromyl chloride (CrO 2 Cl 2 ) are formed.
K 2 Cr 2 O 7 + 4NaCl + 6H 2 SO 4 2KHSO 4 + 4NaHSO 4 + 2CrO 2 Cl 2 + 3H 2 O
(f) Reaction with H 2 O 2 :- Acidified solution of dichromate ions give deep blue colour solution with H 2 O (^2) due to the formation of [CrO(O 2 ) 2 ] or CrO 5. The blue colour fades away gradually due to the decomposition of CrO 5 into Cr +3^ ions and oxygen. Cr
O
O
O
O
Cr O 2 7 H O H CrO H O O
2 4 2 2 2 2 5 5 2 (Butterfly stucture)
(g) Action with HCl :- Potassium dichromate reacts with hydrochloric acid and evolves chlorine.
K 2 Cr 2 O 7 + 14HCl 2KCl + 2CrCl 3 + 7H 2 O + 3Cl 2
(h) Action of con. H 2 SO (^4) (i) In cold, red crystals of chromic anhydride are formed.
K 2 Cr 2 O 7 + 2H 2 SO 4 2CrO 3 + 2KHSO 4 + H 2 O
(ii) On heating the mixture oxygen is evolved.
2K 2 Cr 2 O 7 + 8H 2 SO 4 2K 2 SO 4 + 2Cr 2 (SO 4 ) 3 + 8H 2 O + 3O 2
(i) Oxidising properties The dichromates act as powerful oxidising agent in acidic medium. In presence of dil H 2 SO 4 , K 2 Cr 2 O (^7) liberates Nascent oxygen and therefore acts as an oxidising agent.
K 2 Cr 2 O 7 + 4H 2 SO 4 K 2 SO 4 + Cr 2 (SO 4 ) 3 + 4H 2 O + 3[O]
In terms of electronic concept the Cr O 2 72 ^ ion takes up electrons in the acidic medium and hence acts as
an oxidising agent.
Cr O 2 72 ^ 14 H ^ 6 e ^ 2 Cr 3 7 H O 2
(i) It oxidises iodides to iodine :-
Cr O 2 72 ^ + 14H^ +^ + 6e^ –^ ^ 2Cr^ +3^ + 7H^2 O
[2I–^ I 2 + 2e–^ ] × 3
Cr O 2 72 ^ + 14H^ +^ + 6I^ –^ ^ 2Cr^ +3^ + 3I^2 + 7H^2 O
or K 2 Cr 2 O 7 + 7H 2 SO 4 + 6KI 4K 2 SO 4 + Cr 2 (SO 4 ) 3 + 7H 2 O + 3I 2
(ii) Acidified ferrous sulphate to ferric sulphate
Cr O 2 72 ^ + 14H^ +^ + 6e^ –^ ^ 2Cr^ +3^ + 7H^2 O
Fe+2^ Fe3+^ + e –^ ] × 6
Cr O 2 72 ^ + 14H^ +^ + 6Fe^ +2^ ^ 2Cr^ +3^ + 6Fe^ +3^ + 7H^2 O
or K 2 Cr 2 O 7 + 6FeSO 4 + 7H 2 SO 4 Cr 2 (SO 4 ) 3 + 3Fe 2 (SO 4 ) 3 + 7H 2 O + K 2 SO 4
(iii) Oxidises H 2 S to sulphur
Cr O 2 72 ^ + 14H^ +^ + 6e^ –^ ^ 2Cr^ +3^ + 7H^2 O
H 2 S S + 2H+^ + 2e–^ ] × 3
Cr O 2 72 ^ + 3H 2 S + 8H +^ 2Cr +3^ + 3S + 7H 2 O
or K 2 Cr 2 O 7 + 3H 2 S + 4H 2 SO 4 Cr 2 (SO 4 ) 3 + 3S + 7H 2 O + K 2 SO 4
Similarly, it oxidises sulphites to sulphates, chlorides to chlorine, nitrites to nitrates, thiosulphates to sulphates and sulphur and stannous (Sn +2^ ) salts to stannic (Sn +4^ ) salts.
3 SO 3 ^2 Cr O 2 72 ^ 8 H ^ 3 SO 42 ^ 2 Cr^3 4 H O 2
3 NO 2 Cr O 2 72 ^ 8 H ^ 3 NO 3 ^ 2 Cr 3 4 H O 2
3 S O 2 32 ^ Cr O 2 72 ^ 8 H ^ 3 SO 42 ^ 3 S 2 Cr 3 4 H O 2
6 Cl ^ Cr O 2 72 ^ 14 H ^ 3 Cl 2 2 Cr^3 7 H O 2
3 Sn ^2 Cr O 2 72 ^ 14 H ^ 3 Sn ^4 2 Cr 3 7 H O 2
It oxidises SO 2 to sulphuric acid.
K 2 Cr 2 O 7 + 4H 2 SO 4 K 2 SO 4 + Cr 2 (SO 4 ) 3 + 4H 2 O + 3O
SO 2 + O + H 2 O H 2 SO 4
U s e s (a) For volumetric estimation of ferrous salts, iodides and sulphites. (b) For preparation of other chromium compounds such as chrome alum (K 2 SO 4 , Cr 2 (SO 4 ) 3 .24H 2 O), chrome yellow (PbCrO 4 ) and chrome red (PbCrO 4 .PbO). (c) Used in photography for hardening of gelatin film. (d) It is used in leather industry (chrome tanning) (e) Chromic acid mixture is used for cleaning glassware, consist of K 2 Cr 2 O 7 and Con. H 2 SO 4. (f) In organic chemistry, it is used as an oxidising agent. (g) In dyeing and calico printing. S t r u c t u r e The chromate ion has tetrahedral structure in which four atoms around chromium atom are oriented in a tetrahedral arrangement.
P r o p e r t i e s (a) Colour and M.P. :- Dark violet crystalline solid, M.P. 523 K (b) Solubility :- Moderately soluble is room temperature, but fairly soluble in hot water giving purple solution. (c) Heating :- When heated strongly it decomposes at 746 K to give K 2 MnO 4 and O 2.
2KMnO 4 746 K K 2 MnO 4 + MnO 2 + O 2
Solid KMnO 4 gives KOH, MnO and water vapours, when heated in current of hydrogen.
2KMnO 4 + 5H 2 2KOH + 2MnO + 4H 2 O
(d) Action of alkali :- On heating with alkali, potassium permanganate changes into potassium manganate and oxygen gas is evolved.
4KMnO 4 + 4KOH 4K 2 MnO 4 + 2H 2 O + O 2
(e) Action of con. H 2 SO 4 :- With cold H 2 SO 4 , it gives Mn 2 O 7 which on heating decomposes into MnO 2.
2KMnO 4 + 2H 2 SO 4 Mn 2 O 7 + 2KHSO 4 + H 2 O
2Mn 2 O 7 4MnO 2 + 3O 2
(f) Oxidising character :- KMnO 4 acts as powerful oxidising agent in neutral, alkaline or acidic solution because it liberates nascent oxygen as :- Acidic solution :- Mn +2^ ions are formed
2KMnO 4 + 3H 2 SO 4 K 2 SO 4 + 2MnSO 4 + 3H 2 O + 5[O]
or MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O equal wt^
. M
^
Neutral solution :- MnO 2 is formed
2KMnO 4 + H 2 O 2KOH + 2MnO 2 + 3[O]
or MnO 4 –^ + 2H 2 O + 3e –^ MnO 2 + 4OH –^ equal wt. M
^
During the reaction the alkali produced generates the alkaline medium even if we start from neutral medium. Alkaline medium :- Manganate ions are formed.
2KMnO 4 + 2KOH 2K 2 MnO 4 + H 2 O + [O]
Reactions in Acidic Medium : In acidic medium KMnO 4 oxidizes –
(a) Ferrous salts to feric salts
MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O
Fe+2^ Fe+3^ + e –^ ] × 5
MnO 4 –^ + 5Fe +2^ + 8H +^ Mn 2+^ + 5Fe +3^ + 4H 2 O
(b) Oxalates to CO 2 :
MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O] × 2
C 2 O 4 2–^ 2CO 2 + 2e –^ ] × 5
2MnO 4 –^ + 5C 2 O 4 2–^ + 16H +^ 2Mn +2^ + 10CO 2 + 8H 2 O
(c) Iodides to Iodine
MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O] × 2
2I–^ I 2 + 2e–^ ] × 5
10I –^ + 2MnO 4 –^ + 16H +^ 2Mn +2^ + 5I 2 + 8H 2 O
(d) Sulphites to sulphates
MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O] × 2
SO 3 2–^ + H 2 O SO 4 2–^ + 2H +^ + 2e –^ ] × 5
5SO 3 2–^ + 2MnO 4 –^ + 6H +^ 2Mn +2^ + 5SO 4 2–^ + 3H 2 O
(e) It oxidizes H 2 S to S
MnO 4 –^ + 8H +^ + 5e –^ Mn +2^ + 4H 2 O] × 2
S2–^ S + 2e–^ ] × 5
2MnO 4 –^ + 16H +^ + 5S –2^ 2Mn +2^ + 5S + 8H 2 O
(f) It oxidizes SO 2 to sulphuric acid
2KMnO 4 + 3H 2 SO 4 K 2 SO 4 + 2MnSO 4 + 3H 2 O + 5[O]
SO 2 + H 2 O + [O] H 2 SO 4 ] × 5
2KMnO 4 + 5SO 2 + 2H 2 O K 2 SO 4 + 2MnSO 4 + 2H 2 SO 4
(g) It oxidizes Nitrites to nitrates
2KMnO 4 + 3H 2 SO 4 K 2 SO 4 + 2MnSO 4 + 3H 2 O + 5[O]
KNO 2 + O KNO 3 ] × 5
2KMnO 4 + 5KNO 2 + 3H 2 SO 4 K 2 SO 4 + 2MnSO 4 + 5KNO 3 + 3H 2 O
Reactions in Neutral Medium :
(a) It oxidizes H 2 S to sulphur :
2KMnO 4 + H 2 O 2KOH + 2MnO 2 + 3 [O]
H 2 S + O H 2 O + S] × 3
2KMnO 4 + 3H 2 S 2KOH + 2MnO 2 + 2H 2 O + 3S
(b) It oxidizes Manganese sulphate (MnSO 4 to MnO 2 ) manganese dioxide :
2KMnO 4 + H 2 O 2KOH + 2MnO 2 + 3 [O]
MnSO 4 + H 2 O + O MnO 2 + H 2 SO 4 ] × 3
2KOH + H 2 SO 4 K 2 SO 4 + 2H 2 O
2KMnO 4 + 3MnSO 4 + 2H 2 O 5MnO 2 + K 2 SO 4 + 2H 2 SO 4
P r e p a r a t i o n It is obtained by dissolving scrap iron in dilute sulphuric acid.
Fe + H 2 SO 4 2FeSO 4 + H 2
The solution is crystallised by the addition of alcohol as ferrous sulphate is sparingly soluble in it. P r o p e r t i e s (a) Action of heat : At 300°C, it becomes anhydrous. The anhydrous ferrous sulphate is colourless. The anhydrous salt when strongly heated, breaks up to form ferric oxide with the evolution of SO 2 and SO 3. FeSO 4 ·7HO 2 Green
– 7H O300°C 2 2FeSO 4 White
High temperature Fe^2 O^3 + SO^2 + SO^3
(b) The aqueous solution of ferrous sulphate is slightly acidic due to its hydrolysis. FeSO 4 + 2H 2 O Fe(OH) 2 + H 2 SO (^4) Weak base Strong acid (c) It reduces gold chloride to gold.
AuCl 3 + 3FeSO 4 Au + Fe 2 (SO 4 ) 3 + FeCl 3
(d) It reduces mercuric chloride to mercurous chloride.
[2HgCl 2 Hg 2 Cl 2 + 2Cl] × 3
[3FeSO 4 + 3Cl Fe 2 (SO 4 ) 3 + FeCl 3 ] × 2
6HgCl + 6FeSO 4 3Hg 2 Cl 2 + 2Fe2(SO 4 ) 3 + 2FeCl 3
(e) A cold solution of ferrous sulphate absorbs nitric oxide forming dark brown addition compound, nitroso ferrous sulphate.
FeSO 4 + NO FeSO 4 · NO
Nitroso ferrous sulphate(Brown) The NO gas is evolved when the solution is heated. U s e s
(a) Ferrous sulphate is used for making blue black ink. (b) It is used as a mordant in dyeing. (c) It is also used as an insecticide in agriculture. (d) It is employed as a laboratory reagent and in the preparation of Mohr's salt.
Ferrous-oxide FeO (Black)
Preparation : FeC 2 O (^4) In absence of air^ FeO + CO + CO (^2) Properties : It is stable at high temperature and on cooling slowly disproportionates into Fe 3 O 4 and iron
Ferrous chloride (FeCl 2 )
Preparation : Fe + 2HCl (^) a current of HClheated in^ FeCl 2 + H (^2)
Properties : 2FeCl 3 + H 2 ^ 2FeCl 2 + 2HCl (a) It is deliquescent in air like FeCl (^3) (b) It is soluble in water, alcohol and ether also because it is sufficiently covalent in nature
(c) It volatilises at about 1000°C and vapour density indicates the presence of Fe 2 Cl 4. Above 1300°C density becomes normal (d) It oxidises on heating in air
12FeCl 2 + 3O 2 2Fe 2 O 3 + 8FeCl 3
(e) H 2 evolves on heating in steam
3FeCl 2 + 4H 2 O Fe 3 O 4 + 6HCl + H 2
(f) It can exist as different hydrated form
FeCl 2 · 2H 2 O colourless
FeCl 2 · 4H 2 O pale green
FeCl 2 · 6H 2 O green
1 7. COMPOUND OF ZINC
Zinc oxide (ZnO) zinc white
P r e p a r a t i o n (a) ZnO is formed when ZnS is oxidised
2ZnS + 3O 2 2ZnO + 2SO 2
(b) Zn(OH) 2 on strongly heating gives ZnO
Zn(OH) 2 ^ ZnO + H 2 O
(c) Zinc on burning in air gives ZnO (commercial method)
2Zn + O 2 2ZnO
P r o p e r t i e s (a) ZnO is white when it is cold, a property that has given it a use as a pigment in paints. However, it changes colour, when hot, to a pale yellow. This is due to change in the structure of lattice. (b) ZnO is soluble both in acid and alkali and is thus amphoteric in nature.
ZnO + 2H +^ Zn 2+^ + H 2 O
ZnO + 2OH –^ + H 2 O [Zn(OH) 4 ] 2–^ or ZnO 22
zincate ion
ZnO + 2HCl ZnCl 2 + H 2 O
ZnO + 2NaOH Na 2 ZnO 2 + H 2 O
sodium zincate
(c) ZnO + C >1000°C Zn + CO
ZnO + CO ^ Zn + CO 2
It is preferred to white lead as it is not blackened by H 2 S. It is also used in medicine and in the perparation of Rinman's green (ZnCo 2 O 4 ) Zinc Sulphate (ZnSO 4 ) P r e p a r a t i o n (a) ZnSO 4 · 7H 2 O (also called white vitriol) is formed by decomposing ZnCO 3 with dil. H 2 SO (^4)
ZnCO 3 + H 2 SO 4 ZnSO 4 + H 2 O + CO 2
1 8. COMPOUND OF SILVER
Silver Nitrate (Lunar Caustic) AgNO 3
P r e p a r a t i o n (a) When Ag is heated with dil HNO 3 , AgNO 3 is formed. Crystals separate out on cooling the concentrated solution of AgNO (^3) 3Ag + 4HNO 3 ^ 3 AgNO 3 + NO + 2H 2 O Colourless crystalline compound soluble in H 2 O and alcohol ; m.p. 212°C (b) When exposed to light, it decomposes hence, stored in a brown coloured bottle:
2Ag + 2NO 2 + O 2 , red hot^ 2AgNO 3 , T > 212°C^ 2AgNO 2 + O 2
P r o p e r t i e s (a) It is reduced to metallic Ag by more electropositive metals like Cu, Zn, Mg and also by PH 3.
2AgNO 3 + Cu Cu(NO 3 ) 2 + 2Ag
6AgNO 3 + PH 3 + 3H 2 O 6Ag + 6HNO 3 + H 3 PO 3
(b) It dissolves in excess of KCN: AgNO (^3) ^ KCN ^ AgCN (^) ^ KCN ^ K[Ag(CN) 2 ] white ppt soluble potassium argentocyanide AgNO 3 gives white precipitate with Na 2 S 2 O 3 ; white precipitate changes to black.
2AgNO 3 + Na 2 S 2 O 3 Ag 2 S 2 O 3 + 2NaNO 3
white ppt
Ag 2 S 2 O 3 + H 2 O Ag 2 S + H 2 SO 4
black (c) Ammoniacal AgNO 3 is called Tollen's reagent and is used to identify reducing sugars (including aldehydes): RCHO + 2Ag +^ + 3OH –^ ^ RCOO –^ + 2Ag + 2H 2 O It is called 'silver mirror test' of aldehydes and reducing sugar (like glucose, fructose). Some important reaction of AgNO (^3)
AgNO 3 ^ ,T>212°C^ ^ Ag
Ag
KCN
HNO 3
RCHO,OH
— Na S O^2
3 PH 3
Cu Cu2+ + Ag Ag + H PO + HNO 3 3 3
O + Ag + NO 2 2
Ag (^) K[Ag(CN) ] 2
Ag S O 2 2 3 white H O 2 Ag S black^2
1 9. COMPOUND OF COPPER
Cupric oxide (CuO)
It is called black oxide of copper and is found in nature as tenorite. P r e p a r a t i o n (a) By heating Cu 2 O in air or by heating copper for a long time in air (the temperature should not exceed above 1100°C) Cu 2 O + 12 O 2 2CuO 2Cu + O 2 2CuO (b) By heating cupric hydroxide, Cu(OH) 2 CuO + H 2 O (c) By heating copper nitrate, 2Cu(NO 3 ) 2 2CuO + 4NO 2 + O (^2) (d) On a commercial scale, it is obtained by heating molachite which is found in nature. CuCO 3 · Cu(OH) 2 2CuO + CO 2 + H 2 O P r o p e r t i e s (a) It is black powder and stable to moderate heating. (b) The oxide is insoluble in water but dissolves in acids forming corresponding salts. CuO + 2HCl CuCl 2 + H 2 O CuO + H 2 SO 4 CuSO 4 + H 2 O CuO + 2HNO 3 Cu(NO 3 ) 2 + H 2 O (c) When heated to 1100 – 1200°C, it is converted into cuprous oxide with evolution of oxygen. 4CuO 2Cu 2 O + O (^2) (d) It is reduced to metallic copper by reducing agents like hydrogen, carbon and carbon monoxide. CuO + H 2 Cu + H 2 O CuO + C Cu + CO CuO + CO Cu + CO (^2) U s e s It is used to impart green to blue colour to glazes and glass.
Cupric Chloride, (CuCl 2 · 2H 2 O)
P r e p a r a t i o n
(a) 2Cu + 4HCl + O 2 2CuCl 2 + 2H 2 O
CuO + 2HCl CuCl 2 + H 2 O
Cu(OH) 2 CuCO 3 + 4HCl 2CuCl 2 + 3H 2 O + CO 2
(b) Cu + Cl 2 CuCl 2
CuCl 2 · 2H 2 O HCl gas^ 150°C CuCl 2 + 2H 2 O
