(1) Central metal atom or ion : A complex ion contains a metal atom or ion known as the central metal atom or ion. It is sometimes also called a nuclear atom.
(2) Complex ion : It is an electrically charged radical which is formed by the combination of a simple cation with one or more neutral molecules or simple anions or in some cases positive groups also.
(3) Ligands : Neutral molecules or ions that attach to central metal ion are called ligands. The donor atom associated with the ligands supplies lone pair of electrons to the central metal atom (forming dative bond) may be one or two more. Monodentate (one donor atom), bidentate (two donor atom), tridentate (three donor atom) etc.
Monodentate Ligands (with one donor site)
Anionic Ligands (Negative legands)
When solutions of two or more stable compounds are mixed in stoichiometric (simple molecular) proportions new crystalline compounds called molecular or addition compounds are formed. These are of two types,
(1) Double salts, (2) Co-ordination or Complex compounds
(1) Double salts : Addition compounds, stable in solid state. Dissociate into ions in aqueous solution as such give test for each constituent ion. Examples
(2) Co-ordination or Complex compounds : Addition compound, stable in solid state. Retain their identity even in solution. Central metal ion form dative or coordinate bond with the species surrounding it (ligands). Examples :
Lanthanides and actinides are collectively called f-block elements because last electron in them enters into f-orbitals of the antepenultimate (i.e., inner to penultimate) shell partly but incompletely filled in their elementary or ionic states. The name inner transition, elements is also given to them because they constitute transition series with in transition series (d-block elements) and the last electron enters into antepenultimate shell \[(n-2)f\]. In addition to incomplete d-subshell, their f-subshell is also incomplete. Thus, these elements have three incomplete outer shells i.e., \[(n-2),\,(n-1)\] and n shells and the general electronic configuration of f-block elements is \[(n-2){{f}^{1-14}}(n-1){{d}^{0-10}}n{{s}^{2}}\].
(1) Lanthanides : The elements with atomic numbers 58 to 71 i.e. cerium to lutetium (which come immediately after lanthanum Z = 57) are called lanthanides or lanthanones or rare earths. These elements involve the filling of 4 f-orbitals. Their general electronic configuration is, \[[Xe]4{{f}^{1-14}}5{{d}^{0-10}}6{{s}^{2}}\]. Promethium (Pm), atomic number 61 is the only synthetic (man made) radioactive lanthanide.
Properties of lanthanides
(i) These are highly dense metals and possess high melting points.
(ii) They form alloys easily with other metals especially iron. e.g. misch metal consists of a rare earth element (94–95%), iron (upto 5%) and traces of S, C, Ca and Al, pyrophoric alloys contain Ce (40–5%), La + neodymium (44%), Fe (4–5%), Al (0–5%) and the rest is Ca, Si and C. It is used in the preparation of ignition devices e.g., trace bullets and shells and flints for lighters and cigarette.
(iii) Oxidation state : Most stable oxidation state of lanthanides is +3. Oxidation states + 2 and + 4 also exist but they revert to +3 e.g. \[S{{m}^{2+}},\,E{{u}^{2+}},\,Y{{b}^{2+}}\]lose electron to become +3 and hence are good reducing agents, where as \[C{{e}^{4+}},\,{{\Pr }^{4+}}+T{{b}^{4+}}\] in aqueous solution gain electron to become + 3 and hence are good oxidizing agents. There is a large gap in energy of 4 f and 5 d subshells and thus the number of oxidation states is limited.
(iv) Colour : Most of the trivalent lanthanide ions are coloured both in the solid state and in aqueous solution. This is due to the partly filled f-orbitals which permit \[f-f\] transition. The elements with \[xf\] electrons have a similar colour to those of \[(14-x)\] electrons.
(v) Magnetic properties : All lanthanide ions with the exception of \[L{{u}^{3+}},\,Y{{b}^{3+}}\] and \[C{{e}^{4+}}\] are paramagnetic because they contain unpaired electrons in the \[4f\] orbitals. These elements differ from the transition elements in that their magnetic moments do not obey the simple “spin only” formula \[{{\mu }_{eff}}=\sqrt{n\left( n+2 \right)}\] B.M. where n is equal to the number of unpaired electrons. In transition elements, the orbital contribution of the electron towards magnetic moment is usually quenched by interaction with electric fields of the environment but in case of lanthanides the 4f-orbitals lie too deep in the atom for such quenching to occur. Therefore, magnetic moments of lanthanides are calculated by taking into consideration spin as well as orbital contributions and a more complex formula
\[{{\mu }_{eff}}=\sqrt{4S\left( S+1 \right)+L\left( L+1 \right)}\] B.M.
which involves the more...
(1) Occurrence and extraction of mercury : Cinnabar (HgS) is the only important ore of Hg. It is concentrated by froth floatation method and mercury is extracted from this ore by heating it in air at 773-873 K (auto reduction).
\[HgS+{{O}_{2}}\xrightarrow{273-873\,K}Hg+S{{O}_{2}}\]
The mercury vapours thus obtained are condensed to give liquid metal. Hg thus obtained contains impurities of Zn, Sn and Pb. These are removed by treating the impure metal with dil \[HN{{O}_{3}}\], mercurous nitrate, \[H{{g}_{2}}{{(N{{O}_{3}})}_{2}}\] thus formed react with metals present as impurities forming their nitrates which pass into solution leaving behind pure mercury. However, it is best purified by distillation under reduced pressure.
\[6Hg+8HN{{O}_{3}}(dil.)\xrightarrow{\text{warm}}\]\[3H{{g}_{2}}{{(N{{O}_{3}})}_{2}}+4{{H}_{2}}O+2NO\]
\[Zn+H{{g}_{2}}{{(N{{O}_{3}})}_{2}}\to Zn{{(N{{O}_{3}})}_{2}}+2Hg\]
Similar reaction is given by Pb and Sn.
Properties of mercury : Mercury is less reactive than Zn. It is a liquid at room temperature and has low thermal and electrical conductivity. Mercury forms dimeric mercury (I) ions, \[Hg_{2}^{+2}\] in which the two atoms are bonded by a covalent bond. It is slowly oxidised to HgO at about its boiling point. Hg does not react with dil. HCl or dil.\[{{H}_{2}}S{{O}_{4}}\] but reacts with hot concentrated \[{{H}_{2}}S{{O}_{4}}\] to form \[HgS{{O}_{4}}\], it reacts with both warm dil. and conc. \[HN{{O}_{_{3}}}\] evolving NO and \[N{{O}_{2}}\] respectively.
\[Hg+2{{H}_{2}}S{{O}_{4}}(\text{hot, conc}\text{.})\to HgS{{O}_{4}}+S{{O}_{2}}+2{{H}_{2}}O\]
\[Hg+4HN{{O}_{3}}(\text{conc}\text{.})\to Hg{{(N{{O}_{3}})}_{2}}+2N{{O}_{2}}+2{{H}_{2}}O\]
Hg does not react with steam or water hence can’t form any hydroxide.
Compounds of mercury
(1) Mercuric oxide, HgO : It is obtained as a red solid by heating mercury in air or oxygen for a long time at 673 K
\[2Hg+{{O}_{2}}\xrightarrow{673\,K}2HgO(\text{red})\]
or by heating mercuric nitrate alone or in the presence of Hg
\[2Hg{{(N{{O}_{3}})}_{2}}\xrightarrow{\text{Heat}}\underset{\text{red}}{\mathop{2HgO}}\,+4N{{O}_{2}}+{{O}_{2}}\]
When NaOH is added to a solution of \[HgC{{l}_{2}}\], yellow precipitate of HgO are obtained.
\[H{{g}_{2}}C{{l}_{2}}+2NaOH\xrightarrow{{}}\underset{\text{(yellow)}}{\mathop{HgO\downarrow }}\,+{{H}_{2}}O+2NaCl\]
Red and yellow forms of HgO differ only in their particle size. On heating to 673 K, yellow form changes to red form.
\[\underset{\text{yellow}}{\mathop{HgO}}\,\xrightarrow{673\,K}\underset{\text{red}}{\mathop{HgO}}\,\]
It is used in oil paints or as a mild antiseptic in ointments.
(2) Mercuric chloride, HgCl2 : It is obtained by treating Hg with \[C{{l}_{2}}\] or by heating a mixture of NaCl and \[HgS{{O}_{4}}\] in presence of small amount of \[Mn{{O}_{2}}\] (which oxidises any Hg(I) salts formed during the reaction).
\[HgS{{O}_{4}}+2NaCl\underset{Mn{{O}_{2}}}{\mathop{\xrightarrow{\text{Heat}}}}\,HgC{{l}_{2}}+N{{a}_{2}}S{{O}_{4}}\]
It is a white crystalline solid and is commonly known as corrosive sublimate. It is a covalent compound since it dissolves in organic solvents like ethanol and ether.
It is extremely poisonous and causes death. Its best antidote is white of an egg.
When treated with stannous chloride, it is first reduced to white ppt. of mercurous chloride and then to mercury (black).
\[2HgC{{l}_{2}}+SnC{{l}_{2}}\to \underset{\text{white ppt}\text{.}}{\mathop{H{{g}_{2}}C{{l}_{2}}}}\,+SnC{{l}_{4}}\]
\[H{{g}_{2}}C{{l}_{2}}+SnC{{l}_{2}}\to \underset{\text{grey}}{\mathop{2Hg}}\,+SnC{{l}_{4}}\]
With ammonia it gives a white ppt. known as infusible white ppt.
\[HgC{{l}_{2}}+2N{{H}_{3}}\to Hg(N{{H}_{2}})Cl+N{{H}_{4}}Cl\]
A dilute solution of \[HgC{{l}_{2}}\] is used as an antiseptic.
(3) Mercuric iodide, \[\mathbf{Hg}{{\mathbf{I}}_{\mathbf{2}}}\] : It is obtained when a required amount of KI solution is added to a solution of \[HgC{{l}_{2}}\].
\[HgC{{l}_{2}}+2KI\to \underset{\text{(red)}}{\mathop{Hg{{I}_{2}}}}\,+2KCl\]
Below 400 K, \[Hg{{I}_{2}}\] is red but above 400 K, it turns yellow
\[\underset{\text{(red)}}{\mathop{Hg{{I}_{2}}}}\,\underset{\text{(yellow)}}{\mathop{Hg{{I}_{2}}}}\,\]
\[Hg{{I}_{2}}\] readily dissolves in excess of KI solution to form the \[{{(Hg{{I}_{4}})}^{2-}}\] complex ion.
\[\underset{\text{Red more...
(1) Occurrence of zinc: Zinc does not occur in the native form since it is a reactive metal. The chief ores of zinc are (i) Zinc blende \[(ZnS)\] (ii) Calamine or zinc spar \[(ZnC{{O}_{3}})\] and (iii) Zincite \[(ZnO)\]
(2) Extraction of zinc : Zinc blende, after concentration by Froth floatation process, is roasted in air to convert it into \[ZnO\]. In case of calamine, ore is calcined to get \[ZnO\]. The oxide thus obtained is mixed with crushed coke and heated at 1673 K in fire clay retorts (Belgian Process) when \[ZnO\] gets reduced to metallic zinc. Being volatile at this temperature, the metal distils over and is condensed leaving behind Cd, Pb and Fe as impurities. The crude metal is called spelter. The metal may be refined either by electrolysis or by fractional distillation.
Properties of Zn : Zinc is more reactive than mercury. It is a good conductor of heat and electricity. Zinc readily combines with oxygen to form \[ZnO\]. Pure zinc does not react with non-oxidising acids (\[HCl\] or \[{{H}_{2}}S{{O}_{4}})\] but the impure metal reacts forming \[Z{{n}^{2+}}\] ions and evolving \[{{H}_{2}}\] gas.
\[Zn+2HCl\to ZnC{{l}_{2}}+{{H}_{2}}\uparrow \]
Hot and conc. \[{{H}_{2}}S{{O}_{4}}\] attacks zinc liberating \[S{{O}_{2}}\] gas
\[Zn+2\,{{H}_{2}}S{{O}_{4}}\to ZnS{{O}_{4}}+S{{O}_{2}}+2{{H}_{2}}O\]
Zinc also reacts with both dilute (hot and cold) \[HN{{O}_{3}}\] and conc. \[HN{{O}_{3}}\] liberating nitrous oxide \[({{N}_{2}}O)\], ammonium nitrate \[(N{{H}_{4}}N{{O}_{3}})\] and nitrogen dioxide \[(N{{O}_{2}})\] respectively.
\[4Zn+10HN{{O}_{3}}\] (warm, dilute) \[\to \] \[4\,Zn{{(N{{O}_{3}})}_{2}}+{{N}_{2}}O+5\,{{H}_{2}}O\]
\[4Zn+10HN{{O}_{3}}\] (coldvery dilute)\[\to \] \[4Zn{{(N{{O}_{3}})}_{2}}+N{{H}_{4}}N{{O}_{3}}+3{{H}_{2}}O\]
\[Zn+4HN{{O}_{3}}\](hot and conc.)\[\to Zn{{(N{{O}_{3}})}_{2}}+2N{{O}_{2}}+2{{H}_{2}}O\]
Zinc dissolves in hot concentrated \[NaOH\] forming the soluble sod. Zincate
\[Zn+2NaOH+2{{H}_{2}}O\to N{{a}_{2}}[Zn{{(OH)}_{4}}]+{{H}_{2}}\]
or \[Zn+2NaOH\to N{{a}_{2}}Zn{{O}_{2}}+{{H}_{2}}\]
(3) Special varieties of zinc. (i) Zinc dust : It is prepared by melting zinc and then atomising it with a blast of air.
(ii) Granulated zinc : It is prepared by pouring molten zinc into cold water.
Both these varieties of zinc are used as reducing agents in laboratory.
Compounds of zinc
(1) Zinc oxide (Zinc white or Chinese white), ZnO : It is obtained by burning zinc in air or by heating zinc carbonate or zinc nitrate.
\[2Zn+{{O}_{2}}\xrightarrow{\text{Heat}}2ZnO\]
\[ZnC{{O}_{3}}\xrightarrow{\text{Heat}}ZnO+C{{O}_{2}}\]
\[2Zn{{(N{{O}_{3}})}_{2}}\xrightarrow{\text{Heat}}2ZnO+4N{{O}_{2}}+{{O}_{2}}\]
It is a white powder but becomes yellow on heating and again white on cooling.
It is insoluble in water and is very light and hence commonly known as philosopher’s wool.
It is amphoteric in nature.
\[\underset{\text{(Basic)}}{\mathop{ZnO}}\,+2HCl\to ZnC{{l}_{2}}+{{H}_{2}}O\]
\[\underset{\text{(Acidic)}}{\mathop{ZnO}}\,+2NaOH\to \underset{\text{Sod}\text{.zincate}}{\mathop{N{{a}_{2}}Zn{{O}_{2}}}}\,+{{H}_{2}}O\]
or \[ZnO+2NaOH+{{H}_{2}}O\to \underset{\text{Sod}\text{. tetrahydroxozincate (II)}}{\mathop{N{{a}_{2}}[Zn{{(OH)}_{4}}]}}\,\]
It is reduced both by carbon and \[{{H}_{2}}\] and is used as a white paint
\[ZnO+C\to Zn+CO\]; \[ZnO+{{H}_{2}}\to Zn+{{H}_{2}}O\]
(2) Zinc chloride, \[\mathbf{ZnC}{{\mathbf{l}}_{\mathbf{2}}}\] : It is obtained when \[Zn\] metal, \[ZnO\] or \[ZnC{{O}_{3}}\] is treated with dil. HCl. It crystallizes as \[ZnC{{l}_{2}}.2{{H}_{2}}O\] and becomes anhydrous on heating. \[ZnC{{l}_{2}}\] is highly deliquescent and is highly soluble in \[{{H}_{2}}O\] and also readily dissolves in organic solvents like acetone, alcohol, ether etc. its aqueous solution is acidic due to hydrolysis.
\[ZnC{{l}_{2}}+{{H}_{2}}O\to Zn(OH)Cl+HCl\]
Anhydrous \[ZnC{{l}_{2}}\] is used as a Lewis acid catalyst in organic reactions. Mixed with moist zinc oxide, it is used for filling teeth and its solution is used for preserving timber. Anhydrous \[ZnC{{l}_{2}}\] used as a Lucas reagent with more...
(1) Occurrence of gold : Gold is mainly found in native state either as vein gold, placer gold or alluvial gold. It is also present to a small extent in the combined state as sulphide, telluride and arsenosulphide. It is considered to be the king of metal.
Some important ores of gold are:
(i) Calaverite, \[AuT{{e}_{2}}\] (ii) Sylvanite, \[AuAgT{{e}_{2}}\] and
(iii) Bismuth aurite, \[BiA{{u}_{2}}\]
(2) Extraction of gold : (i) Mac-Arthur-Forest Cyanide process : The powdered gold ore, after concentration by Froth-floatation process, is roasted to remove easily oxidisable impurities of tellurium, arsenic and sulphur. The roasted ore is then treated with a dilute solution of KCN in presence of atmospheric oxygen when gold dissolves due to the formation of an aurocyanide complex.
\[4\,Au+8\,KCN+2\,{{H}_{2}}O+{{O}_{2}}\to \underset{\text{Solution}}{\mathop{4K[Au{{(CN)}_{2}}]}}\,+4KOH\]
The metal is then extracted by adding zinc dust.
\[2\,K\,[Au{{(CN)}_{2}}]+Zn\to {{K}_{2}}[Zn{{(CN)}_{4}}]+\underset{\text{ppt}\text{.}}{\mathop{2Au\downarrow }}\,\]
(ii) Plattner’s chlorine process : The roasted ore is moistened with water and placed in wooden vats with false perforated bottoms. It is saturated with current of chlorine, gold chloride thus formed is leached with water and the solution is treated with a reducing agent such as \[FeS{{O}_{4}}\] or \[{{H}_{2}}S\] to precipitate gold.
\[AuC{{l}_{3}}+3FeS{{O}_{4}}\to Au\downarrow +FeC{{l}_{3}}+F{{e}_{2}}{{(S{{O}_{4}})}_{3}}\]
\[2\,AuC{{l}_{3}}+3{{H}_{2}}S\to 6HCl+3S+2Au\downarrow \]
The impure gold thus obtained contains impurities of Ag and Cu. The removal of Ag and Cu from gold is called parting. This is done by heating impure gold with conc. \[{{H}_{2}}S{{O}_{4}}\](or \[HN{{O}_{3}})\] when Ag and Cu dissolve leaving behind Au.
\[Cu+2{{H}_{2}}S{{O}_{4}}\to CuS{{O}_{4}}+S{{O}_{2}}+2{{H}_{2}}O\]
\[2Ag+2{{H}_{2}}S{{O}_{4}}\to A{{g}_{2}}S{{O}_{4}}+S{{O}_{2}}+2{{H}_{2}}O\]
Properties of Gold: Gold is a yellow, soft and heavy metal. Gold and Ag are called noble metals since they are not attacked by atmospheric oxygen. However, Ag gets tarnished when exposed to air containing traces of \[{{H}_{2}}S\]. Gold is malleable, ductile and a good conductor of heat and electricity.
Pure gold is soft. It is alloyed with Ag or Cu for making jewellery. Purity of gold is expressed in terms of carats. Pure gold is 24 carats. Gold ’14 carats’ means that it is an alloy of gold which contains 14 parts by weight of pure gold and 10 parts of copper per 24 parts by weight of the alloy. Thus the percentage of gold in ’14 carats” of gold is = \[\frac{100}{24}\times 14=58.3%\].
Most of the jewellery is made from 22 carat gold (91.66% pure gold). Gold is quite inert. It does not react with oxygen, water and acids but dissolves in aqua regia
\[3HCl+HN{{O}_{3}}\to NOCl+2\,{{H}_{2}}O+2Cl]\times 3\]
\[Au+3\,\,Cl\to AuC{{l}_{3}}]\times 2\]
\[2\,Au+9\,HCl+3\,HN{{O}_{3}}\to \underset{\text{Auric chloride }\,\,\,\,\,\,\,\text{Nitrosyl chloride}}{\mathop{2\,AuC{{l}_{3}}+6\,{{H}_{2}}O+3NOCl\,\,\,\,\,\,\,\,\,\,\,}}\,\]
Oxidation states of gold: The principal oxidation states of gold are + 1 and + 3 though + 1 state is more stable than + 3.
Compounds of gold
(1) Auric chloride, \[\mathbf{AuC}{{\mathbf{l}}_{\mathbf{3}}}\]: It is prepared by passing dry \[C{{l}_{2}}\] over finely divided gold powder at 573 K
\[2\,Au+3C{{l}_{2}}\xrightarrow{573\,K}2\,AuC{{l}_{3}}\]
It is a red coloured crystalline solid soluble in water and decomposes on heating to give gold (I) chloride and \[C{{l}_{2}}\]
\[AuC{{l}_{3}}\xrightarrow{\text{Heat}}AuCl+C{{l}_{2}}\]
It dissolves in conc. \[HCl\] forming chloroauric acid
\[AuC{{l}_{3}}+HCl\to H[AuC{{l}_{4}}]\]
Chloroauric acid is used in photography for toning silver more...
(1) Ores : Copper pyrites (chalcopyrite) \[CuFe{{S}_{2}},\] Cuprite (ruby copper) \[C{{u}_{2}}O,\] Copper glance \[(C{{u}_{2}}S)\], Malachite \[[Cu{{(OH)}_{2}}.\,CuC{{O}_{3}}],\] Azurite \[[Cu{{(OH)}_{2}}.\,2CuC{{O}_{3}}]\]
(2) Extraction : Most of the copper (about 75%) is extracted from its sulphide ore, copper pyrites.
Concentration of ore : Froth floatation process.
Roasting : Main reaction :
\[2CuFe{{S}_{2}}+{{O}_{2}}\to C{{u}_{2}}S+2FeS+S{{O}_{2}}\].
Side reaction : \[2C{{u}_{2}}S+3{{O}_{2}}\to 2C{{u}_{2}}O+2S{{O}_{2}}\]
\[2FeS+3{{O}_{2}}\to 2FeO+2S{{O}_{2}}\].
Smelting : \[FeO+Si{{O}_{2}}\to FeSi{{O}_{3}}(\text{slag)}\]
\[C{{u}_{2}}O+FeS\to FeO+C{{u}_{2}}S\]
The mixture of copper and iron sulphides melt together to form 'matte' \[(C{{u}_{2}}S+FeS)\] and the slag floats on its surface.
Conversion of matte into Blister copper (Bessemerisation) : Silica is added to matte and a hot blast of air is passed \[FeO+Si{{O}_{2}}\to FeSi{{O}_{3}}(\text{slag})\]. Slag is removed. By this time most of iron sulphide is removed. \[C{{u}_{2}}S+2C{{u}_{2}}O\to 6Cu+S{{O}_{2}}\]
Blister copper : Which contain about 98% pure copper and 2% impurities (Ag, Au, Ni, Zn etc.)
Properties of copper : It has reddish brown colour. It is highly malleable and ductile. It has high electrical conductivity and high thermal conductivity. Copper is second most useful metal (first being iron). It undergoes displacement reactions with lesser reactive metals e.g. with Ag. It can displace Ag from \[AgN{{O}_{3}}\]. The finally divided Ag so obtained is black in colour.
Copper shows oxidation states of +1 and +2. Whereas copper (I) salts are colourless, copper (II) salts are blue in colour. Cu (I) salts are less stable and hence are easily oxidised to Cu (II) salts \[(2C{{u}^{+}}\to C{{u}^{2+}}+Cu)\]. This reaction is called disproportionation.
(1) In presence of atmospheric \[C{{O}_{2}}\] and moisture, copper gets covered with a green layer of basic copper carbonate (green layer) which protects the rest of the metal from further acton.
\[Cu+{{O}_{2}}+C{{O}_{2}}+{{H}_{2}}O\to \underset{\text{(green layer)}}{\mathop{Cu{{(OH)}_{2}}CuC{{O}_{3}}}}\,\]
(2) In presence of oxygen or air, copper when heated to redness (below 1370K) first forms red cuprous oxide which changes to black cupric oxide on further heating. If the temperature is too high, cupric oxide changes back to cuprous oxide
\[4Cu+{{O}_{2}}\xrightarrow{\text{Below 1370}\,\text{K}}\underset{\text{(Red)}}{\mathop{2C{{u}_{2}}O}}\,\underset{\text{Above 1370 }K}{\mathop{\xrightarrow{{{\text{O}}_{\text{2}}}}}}\,\underset{\text{(Black)}}{\mathop{4CuO}}\,\]
\[CuO+Cu\]\[\xrightarrow{\text{High temp}\text{.}}\]\[C{{u}_{2}}O\]
(3) Action of acids. Non oxidising dil. acids such as \[HCl,{{H}_{2}}S{{O}_{4}}\] have no action on copper. However, copper dissolves in these acids in presence of air.
\[Cu+2HCl+\frac{1}{2}{{\text{O}}_{\text{2}}}\text{(air)}\to CuC{{l}_{2}}+{{H}_{2}}O\]
With dil. \[HN{{O}_{3}}\], \[Cu\] liberates \[NO\] (nitric oxide)
\[3Cu+8HN{{O}_{3}}\to 3Cu{{(N{{O}_{3}})}_{2}}+2NO+4{{H}_{2}}O\]
With conc. \[HN{{O}_{3}}\], copper gives \[N{{O}_{2}}\]
\[Cu+4HN{{O}_{3}}\to Cu{{(N{{O}_{3}})}_{2}}+2N{{O}_{2}}+2{{H}_{2}}O\]
With hot conc. \[{{H}_{2}}S{{O}_{4}}\], copper gives \[S{{O}_{2}}\]
\[Cu+2{{H}_{2}}S{{O}_{4}}\to CuS{{O}_{4}}+S{{O}_{2}}+2{{H}_{2}}O\]
Compounds of Copper
(1) Halides of copper : Copper (II) chloride, \[CuC{{l}_{2}}\] is prepared by passing chlorine over heated copper. Concentrated aqueous solution of \[CuC{{l}_{2}}\] is dark brown but changes first to green and then to blue on dilution.
On heating, it disproportionates to copper (I) chloride and chlorine
\[2CuC{{l}_{2}}\xrightarrow{\text{Heat}}2CuCl+C{{l}_{2}}\]
It is used as a catalyst in the Daecon’s process for the manufacture of chlorine.
Copper (I) chloride, \[CuCl\] is a white solid insoluble in water. It is obtained by boiling a solution of \[CuC{{l}_{2}}\] with excess of copper turnings and conc. \[HCl\].
\[CuC{{l}_{2}}+Cu\xrightarrow{\text{HCl}}2CuCl\]
It dissolves in conc. \[HCl\] due to the formation of complex \[H[CuC{{l}_{2}}]\]
\[CuCl+HCl\to H[CuC{{l}_{2}}]\]
It is used as a catalyst alongwith \[N{{H}_{4}}Cl\] in the preparation of synthetic rubber.
(2) Cuprous oxide more...
(1) Ores of iron : Haematite \[F{{e}_{2}}{{O}_{3}}\], Magnetite \[(F{{e}_{3}}{{O}_{4}}),\] Limonite \[(F{{e}_{2}}{{O}_{3}}.3{{H}_{2}}O)\], Iron pyrites \[(Fe{{S}_{2}}),\] Copper pyrities \[(CuFe{{S}_{2}})\] etc.
(2) Extraction : Cast iron is extracted from its oxides by reduction with carbon and carbon monoxide in a blast furnace to give pig iron.
Roasting : Ferrous oxide convert into ferric oxide.
\[F{{e}_{2}}{{O}_{3}}.\,3{{H}_{2}}O\to F{{e}_{2}}{{O}_{3}}+3{{H}_{2}}O\];\[2FeC{{O}_{3}}\to 2FeO+2C{{O}_{2}}\]
\[4FeO+{{O}_{2}}\to 2F{{e}_{2}}{{O}_{3}}\]
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)
\[2C+{{O}_{2}}\to 2CO\]
(ii) \[F{{e}_{2}}{{O}_{3}}+3CO2Fe+3C{{O}_{2}}\]. It is a reversible and exothermic reaction. Hence according to Le-chatelier principle more iron will be produced in the furnace at lower temp.
\[\underset{\text{(it is not reversible)}}{\mathop{F{{e}_{2}}{{O}_{3}}+CO\to 2FeO+C{{O}_{2}}}}\,\]
(iii) \[FeO+C\underset{\begin{smallmatrix}
\text{endothermic } \\
\text{ reaction}
\end{smallmatrix}}{\mathop{\xrightarrow{1073\,K}}}\,\] \[Fe+CO\]
The gases leaving at the top of the furnace contain up to 28% CO and are burnt in cowper's stove to pre-heat the air for blast
Varieties of iron : The three commercial varieties of iron differ in their carbon contents. These are;
(1) Cast iron or Pig-iron : It is most impure form of iron and contains highest proportion of carbon (2.5–4%).
(2) Wrought iron or Malleable iron : It is the purest form of iron and contains minimum amount of carbon (0.12–0.25%).
(3) Steel : It is the most important form of iron and finds extensive applications. Its carbons content (Impurity) is mid-way between cast iron and wrought iron. It contains 0.2–1.5% carbon. Steels containing 0.2–0.5% of carbon are known as mild steels, while those containing 0.5–1.5% carbon are known as hard steels.
Steel is generally manufactured from cast iron by three processes, viz, (i) Bessemer Process which involves the use of a large pear-shaped furnace (vessel) called Bessemer converter, (ii) L.D. process and (iii) open hearth process, Spiegeleisen (an alloy of Fe, Mn and C) is added during manufacture of steel.
Heat treatment of steels : Heat treatment of steel may be defined as the process of carefully heating the steel to high temperature followed by cooling to the room temperature under controlled conditions. Heat treatment of steel is done for the following two purposes,
(i) To develop certain special properties like hardness, strength, ductility etc. without changing the chemical composition.
(ii) To remove some undesirable properties or gases like entrapped gases, internal stresses and strains. The various methods of heat treatment are,
(a) Annealing : It is a process of heating steel to redness followed by slow cooling.
(b) Quenching or hardening : It is a process of heating steel to redness followed by sudden cooling by plunging the red hot steel into water or oil.
(c) Tempering : It is a process of heating the hardened or quenched steel to a temperature much below redness (473–623K) followed by slow cooling.
(d) Case-hardening : It is a process of giving a thin coating of hardened steel to wrought iron or to a strong and flexible mild steel by heating it in contact with charcoal followed by quenching more...
Potassium Permanganate, \[(KMn{{O}_{4}})\]
Potassium permanganate is a salt of an unstable acid \[HMn{{O}_{4}}\] (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.
\[Mn{{O}_{2}}+2KOH\xrightarrow{fused}\underset{potassium\,manganite}{\mathop{{{K}_{2}}Mn{{O}_{3}}}}\,+4{{H}_{2}}O]\times 2\]
\[2{{K}_{2}}Mn{{O}_{3}}+{{O}_{2}}\to 2{{K}_{2}}Mn{{O}_{4}}+\,2\,{{H}_{2}}O\]
\[\underset{pyrolusite}{\mathop{2Mn{{O}_{2}}}}\,+4KOH+{{O}_{2}}\xrightarrow{fused}\underset{\underset{\left[ dark-green\,\,mass \right]}{\mathop{potassium\,manganate}}\,}{\mathop{2{{K}_{2}}Mn{{O}_{4}}}}\,+2{{H}_{2}}O\]
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 \[C{{O}_{2}}\] is passed
\[\underset{potassium\,manganate}{\mathop{3{{K}_{2}}Mn{{O}_{4}}+}}\,2{{H}_{2}}O\to \underset{potassium\,permanganate}{\mathop{2KMn{{O}_{4}}}}\,+Mn{{O}_{2}}\downarrow +4KOH\]
\[2C{{O}_{2}}+4KOH\to 2{{K}_{2}}C{{O}_{3}}+2{{H}_{2}}O\]
When chlorine or ozone is passed
\[2{{K}_{2}}Mn{{O}_{4}}+C{{l}_{2}}\to 2KMn{{O}_{4}}+2KCl\]
\[2{{K}_{2}}Mn{{O}_{4}}+{{O}_{3}}+{{H}_{2}}O\to 2KMn{{O}_{4}}+2KOH+{{O}_{2}}\left( g \right)\]
The purple solution so obtained is concentrated and dark purple, needle-like crystals having metallic lustre are obtained.
Electrolytic method : Presently, potassium manganite \[({{K}_{2}}Mn{{O}_{4}})\] is oxidised electrolytically. The electrode reactions are,
At anode: \[\underset{green}{\mathop{2MnO_{4}^{2-}}}\,\to \underset{purple}{\mathop{2MnO_{4}^{-}}}\,+2{{e}^{-}}\]
At cathode: \[2{{H}^{+}}+2{{e}^{-}}\to {{H}_{2}}\left( g \right)\]
The purple solution containing KMnO4 is evaporated under controlled condition to get crystalline sample of potassium permanganate.
Physical properties
\[KMn{{O}_{4}}\] 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 \[KMn{{O}_{4}}\] has a purple colour.
Chemical properties : Some important chemical reactions of \[KMn{{O}_{4}}\] are given below,
Action of heat : \[KMn{{O}_{4}}\] is stable at room temperature, but decomposes to give oxygen at higher temperatures.
\[2KMn{{O}_{4}}\left( s \right)\xrightarrow{heat}{{K}_{2}}Mn{{O}_{4}}\left( s \right)+Mn{{O}_{2}}+{{O}_{2}}\left( g \right)\]
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 \[KMn{{O}_{4}}\] (to be more specific, of \[MnO_{4}^{-}\]) is indicated by high positive reduction potentials for the following reactions.
Acidic medium
\[MnO_{4}^{-}+8{{H}^{+}}+5{{e}^{-}}\rightleftharpoons M{{n}^{2+}}+4{{H}_{2}}O\,\,\,\,\,\,{{E}^{o}}=1.51\,V\]
Alkaline medium
\[MnO_{4}^{-}+2{{H}_{2}}O+3{{e}^{-}}\rightleftharpoons Mn{{O}_{2}}+4O{{H}^{-}}\,\,\,{{E}^{o}}=1.23\,V\,\]
In strongly alkaline solutions and with excess of \[MnO_{4}^{-}\], the reaction is \[MnO_{4}^{-}+{{e}^{-}}\rightleftharpoons MnO_{4}^{2-}\,\,\,\,\,\,\,\,{{E}^{o}}=0.56\,V\]
There are a large number of oxidation-reduction reactions involved in the chemistry of manganese compounds. Some typical reactions are,
In the presence of excess of reducing agent in acidic solutions permanganate ion gets reduced to manganous ion, e.g., \[5F{{e}^{2+}}+MnO_{4}^{-}+8{{H}^{+}}\to 5F{{e}^{3+}}+M{{n}^{2+}}+4{{H}_{2}}O\]
An excess of reducing agent in alkaline solution reduces permanganate ion only to manganese dioxide e.g.,
\[3NO_{2}^{-}+MnO_{4}^{-}+2O{{H}^{-}}\to 3NO_{3}^{-}+Mn{{O}_{2}}+{{H}_{2}}O\]
In faintly acidic and neutral solutions, manganous ion is oxidised to manganese oxidised to manganese dioxide by permanganate.
\[2MnO_{4}^{-}+3M{{n}^{+2}}+2{{H}_{2}}O\to 5Mn{{O}_{2}}+4{{H}^{+}}\]
In strongly basic solutions, permangante oxidises manganese dioxide to manganate ion.
\[Mn{{O}_{2}}+2MnO_{4}^{-}+4O{{H}^{-}}\to 3MnO_{4}^{2-}+2{{H}_{2}}O\]
In acidic medium, \[KMn{{O}_{4}}\] oxidises,
Ferrous salts to ferric salts
\[2KMn{{O}_{4}}+3{{H}_{2}}S{{O}_{4}}\to {{K}_{2}}S{{O}_{4}}+2MnS{{O}_{4}}+3{{H}_{2}}O+5\left[ O \right]\]
\[\underset{\underline{2KMn{{O}_{4}}+8{{H}_{2}}S{{O}_{4}}+10FeS{{O}_{4}}\to {{K}_{2}}S{{O}_{4}}+2MnS{{O}_{4}}+5F{{e}_{2}}{{\left( S{{O}_{4}} \right)}_{3}}+8{{H}_{2}}O}}{\mathop{\underline{2FeS{{O}_{4}}+{{H}_{2}}S{{O}_{4}}+\left[ O \right]\to F{{e}_{2}}{{\left( S{{O}_{4}} \right)}_{3}}+{{H}_{2}}O]\times 5}}}\,\]
Ionic equation
\[2MnO_{4}^{-}+16{{H}^{+}}+10F{{e}^{2+}}\to 2M{{n}^{2+}}+10F{{e}^{3+}}+8{{H}_{2}}O\]
The reaction forms the basis of more...
Lanthanides and
actinides are collectively called f-block elements because last electron
in them enters into f-orbitals of the antepenultimate (i.e.,
inner to penultimate) shell partly but incompletely filled in their elementary
or ionic states. The name inner transition, elements is also given to them
because they constitute transition series with in transition series (d-block
elements) and the last electron enters into antepenultimate shell (n-2)f.
In addition to incomplete d-subshell, their f-subshell is also
incomplete. Thus, these elements have three incomplete outer shells i.e., (n?2),
(n?1) and n shells and the general electronic configuration of f-block
elements is (n?2)\[{{f}^{1-14}}\,(n-1){{d}^{0-10}}n{{s}^{2}}\].
(1)
Lanthanides : The elements with atomic numbers 58 to 71 i.e.
cerium to lutetium (which come immediately after lanthanum Z = 57) are called
lanthanides or lanthanones or rare earths. These elements involve
the filling of 4 f-orbitals. Their general electronic configuration is, \[[Xe]4{{f}^{1-14}}5{{d}^{0-10}}6{{s}^{2}}\].
Promethium (Pm), atomic number 61 is the only synthetic (man made)
radioactive lanthanide.
Properties of lanthanides
(i)
These are highly dense metals and possess high melting points.
(ii)
They form alloys easily with other metals especially iron. e.g. misch
metal consists of a rare earth element (94?95%), iron (upto 5%) and
traces of S, C, Ca and Al, pyrophoric alloys
contain Ce (40?5%), La + neodymium (44%), Fe (4?5%), Al
(0?5%) and the rest is Ca, Si and C. It is used in the
preparation of ignition devices e.g., trace bullets and shells and flints for
lighters and cigarette.
(iii)
Oxidation state : Most stable oxidation state of lanthanides is
+3. Oxidation states + 2 and + 4 also exist but they revert to +3 e.g. \[S{{m}^{2+}},\,E{{u}^{2+}},\,Y{{b}^{2+}}\]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.
(iv)
Colour : Most of the trivalent lanthanide ions are coloured both in
the solid state and in aqueous solution. This is due to the partly filled f-orbitals
which permit f?f transition. The elements with xf electrons have
a similar colour to those of (14 ? x) electrons.
(v) Magnetic
properties : All lanthanide ions with the exception of Lu3+,
Yb3+ and Ce 4+ are paramagnetic because
they contain unpaired electrons in the 4 f orbitals. These elements
differ from the transition elements in that their magnetic moments do not obey
the simple ?spin only? formula \[{{\mu }_{eff}}=\sqrt{n\left( n+2
\right)}\] B.M. where n is equal to the number of unpaired electrons. In
transition elements, the orbital contribution of the electron towards magnetic
moment is usually quenched by interaction with electric fields of the
environment but in case of lanthanides the 4f-orbitals lie too deep in
the atom for such quenching to occur. Therefore, magnetic moments of
lanthanides are calculated by taking into consideration spin as well as orbital
contributions and a more complex formula
more...
