Chapter 9: d & f Block

Chapter 9: d & f Block

d & f-block elements & their important compounds

 Periodic trends and chemical properties

Silver nitrate

Silver nitrate, AgNO₃ (Lunar Caustic)

Physical Properties

(i) It is a colourless crystalline compound.

(ii) Soluble  in water and alcohol.

(iii) It melts at 212°C.

Chemical Properties

(i) It possesses powerful corrosive action on organic tissues, which at turns black especially in presence of light. The blackening is due to finely divided metalic silver, reduced by organic tissure It is therefore, stored in colored Bottles. 

(ii) Solutions of halides phosphates, sulfides chromates thiocyanates, sulphates and thiosulphates salt  with silver nitrate solution.

(iii) Ammonical silver nitrate is called as Tollen's reagent and used to identify reducing sugars and aldehydes.

It is known as silver mirror test of aldehydes and reducing sugars.

Ag₂O + HCHO ¾¾®  2Ag + HCOOH.

Ag₂O + C₆H₁₂O₆ ¾¾®  2Ag + C₆H₁₂O₇.

(iv)  Reactions with Iodine :

6 AgNO₃ (excess) + 3I₂ +3H₂O ¾¾®  AgIO₃ + 5AgI + 6HNO₃

5 AgNO₃  + 3I₂ (excess) +3H₂O ¾¾®   HIO₃  + 5AgI + 5HNO₃

Uses : It is used

1. as a laboratory reagent for the identification of various acidic especially for Cl, Br and I
2. Tollen's reagent is used in organic chemistry for testing aldehydes reducing sugars etc.
3. for making AgBr, used in photography.
4. in the preparation of inks and hair dyes.
5. in preparation of silver mirror.

2. Photography :

(i) A photographic film consists of a light sensitive emulsion of fine particles (grains) of silver salts in gelatine spread on a clear celluloid strip or a glass plate. AgBr is mainly used as the light sensitive material.

(ii) The film is placed in a camera. When the photograph is exposed, light from the subject enters the camera and is focussed by the lens to give a sharp image on the film. The light starts a photochemical reaction by exciting a halide ion, which loses an electron. The electron moves in a conduction band to the surface of the grain, where it reduces a Ag+ ion to metallic silver.

 2AgBr(s)  2Ag + Br₂

(iii) In modern photography only a short exposure of perhaps 1/100th of a second is used. In this short time, only a few atoms of silver (perhaps 10–50) are produced in each grain exposed to light. Parts of the film which have been exposed to the bright parts of the subject contain a lot of grains with some silver.

(iv) Next the film is placed a developer solution. This is a mild reducing agent, usually containing quinol. Its purpose is to reduce more silver halide to Ag metal. Ag is deposited mainly where there are already some Ag atoms. Thus the developing process intensifies the latent image on the film so it becomes visible. 

(v) If the film was brought out into daylight at this stage, the unexposed parts of the emulsion would turn black and thus destroy the picture. To prevent this happening any unchanged silver halides are removed by placing the film in a fixer solution. A solution of sodium thiosulphate is used as fixer. It forms a soluble complex with silver halides.

AgBr + 2Na₂S₂O₃ ® Na₃[Ag(S₂O₃)₂] + NaBr

After fixing, the film can safely be brought out into daylight. This is called "negative".

Light is passed through the negative onto a piece of paper coated with AgBr emulsion. This is then developed and fixed in the same way as before.

Potassium Permanganatic (KMnO₄)

Physical Properties:

Purple coloured crystalline compound.

Moderately soluble in water at room temperature.

Structure of manganate and permanganate ion.

Chemical Properties

(i) Heating effect

MnO₄^2– in dilute alkaline, water and acidic solutions is unstable and disproportionates to give MnO₄^– and MnO₂.

3MnO₄^2– + 4H⁺ ¾®2MnO₄^– + MnO₂ ¯ + 2H₂O

3MnO₄^2– + 2H₂O ¾® 2MnO₄^– + MnO₂ ¯ + 4OH^–        

or

3MnO₄^2– + 3H₂O ¾® 2MnO₄^– + MnO (OH)₂ ¯+ 4OH^– 

On heating in current of H₂ , solid KMnO₄ gives MnO

2KMnO₄  + 5H₂  2KOH + 2MnO + 4H₂O

(ii) On treatment with concentrated H₂SO₄ (KMnO₄ is taken in excess), it forms manganese heptoxide via permanganyl sulphate which decomposes explosively on heating.

2KMnO₄ + 3H₂SO₄ ¾® 2KHSO₄ + (MnO₃)₂SO₄ + 2H₂O

(MnO₃)₂SO₄ + H₂O  ¾®  Mn₂O₇ + H₂SO₄

Mn₂O₇ ¾®  2MnO₂ + O₂

 KMnO₄ + 3H₂SO₄ (conc.) ¾® K⁺ + MnO₃⁺ (green) + 3HSO₄^– + H₃O⁺.

 (iii)  Potassium permanganate is a powerful oxidising agent.

Potassium permanganate acts as an oxidising agent in alkaline, neutral or acidic solutions.

A mixture of sulphur, charcoal and KMnO₄ forms an explosive powder. A mixture of oxalic acid and KMnO₄ catches fire spontaneous after a few seconds. The same thing happens when glycerine is poured over powdered KMnO₄.

In alkaline & neutral medium :

In strongly alkaline medium KMnO₄ is reduced to manganate.

2KMnO₄ + 2KOH (conc.)  ¾® 2K₂ MnO₄ + H₂O + [O]

or    e^– + MnO₄^– ¾® MnO₄^2–

if solution is dilute then K₂MnO₄ is converted in to MnO₂ which appears as a  brownish precipitate.

2K₂MnO₄ + 2H₂O ¾® 2MnO₂ + 4KOH + 2[O]

2e^– + 2H₂O + MnO₄^2– ¾® MnO₂ + 4OH^–

This type of behaviour is shown by KMnO4 itself in neutral medium.

3e^– + 2H₂O + MnO₄^– ¾® MnO₂ + 4OH^–

In alkaline or neutral medium KMnO₄ shows following oxidising properties.

(a) It oxidises ethene to glycol.

In alkaline medium KMnO₄ solution is also known as Bayer’s reagent (1% alkaline KMnO₄ solution).

(b) It oxidises iodide into iodate.

(c)  H₂S is oxidised into sulphur :

2MnO₄^– + 3H₂S  ¾®  2MnO₂ + 2OH^– +2H₂O + 3S

(d) Manganous salt is oxidised to MnO₂ ; the presence of zinc sulphate or zinc oxide catalyses the oxidation :

2MnO₄^– + 3Mn₂⁺ + 2H₂O ¾® 5MnO₂ + 4H⁺

(e) In neutral/ faintly alkaline solution thiosulphate is quantitatively oxidised to sulphate.

8MnO₄^– + 3S₂O₃^2– + H₂O  ¾® 8MnO₂ + 6SO₄^2– + 2OH^–

(f) In the presence of sodium hydroxide, sodium sulphite is oxidised in to sodium sulphate.   

2MnO₄^– + 3SO₃^2– + 3H₂O ¾® 2MnO(OH)_2¯+ 3SO₄^2– + 2OH^–.

In acidic medium (in presence of dilute H₂SO₄) :

Manganous sulphate is formed. The solution becomes colourless.

2KMnO₄ + 3H₂SO₄ ¾® K₂SO₄ + 2MnSO₄ + 3H₂O + 5[O]

or       MnO4¯ + 8H+ + 5e^– ¾® Mn²⁺ + 4H₂O

This medium is used in quantitative (volumetric) estimations. The equivalent mass of KMnO4 in acidic medium is = .

(a) Ferrous salts are oxidised to ferric salts.

MnO₄¯ + 5Fe²⁺ + 8H+ ¾® 5Fe³⁺ + Mn²⁺ + 4H₂O

(b) Iodine is evolved from potassium iodide.

2MnO4¯ + 10I¯ + 16H⁺ ¾® 2Mn²⁺ + 5I₂ + 8H₂O

(c) H₂S is oxidised to sulphur :

2MnO₄^– + 16H⁺+ 5S^2– ¾® 2Mn²⁺ + 8H₂O + 5S

(d) SO₂ is oxidised to H₂SO₄ :

2MnO₄^– + 5SO₂ + 2H₂O ¾® 5SO₄^2– + 2Mn²⁺ + 4H⁺

(e) Nitrites are oxidised to nitrates :

2MnO₄^– + 5NO₂^– + 6H⁺ ¾® 2Mn²⁺ + 3H₂O + 5NO₃^–

(f) Oxalic acid is oxidised to CO₂ :

This reaction is slow at room temperature, but is rapid at 60ºC.

Mn(II) ions produced catalyse the reaction; thus the reaction is autocatalytic.

2MnO₄^– + 16H⁺ + 5C₂O₄^2–2Mn²⁺ + 8H₂O + 10CO₂

(g) HCHO is oxidised to HCOOH

2MnO₄^– + 5HCHO + 6H⁺ ¾® 2Mn²⁺ + 5HCOOH + 3H₂O.

(h) It oxidises hydrogen halides (HCl, HBr or HI) into X₂ (halogen)

2MnO4^– + 16H⁺ + 10X^– ¾® 2Mn²⁺ + 8H₂O + 5X₂ 

(i)  H₂O₂ is oxidised to O₂.

 2MnO₄^– + 6H⁺ + 5H₂O₂ ¾® 5O₂ + Mn2⁺ + 8H₂O.

Uses :It is used

1. KMnO₄ is used as an oxidising agent in laboratory and industry.
2. Alkaline potassium permanganate is called Bayer's reagent. This reagent is used in organic chemistry for the test of unsaturation. KMnO₄ is used in the manufacture of saccharin, benzoic acid, acetaldehyde etc.
3. KMnO₄ is used in qualitative analysis for detecting halides, sulphites, oxalates, etc.

POTASSIUM DICHROMATE (K₂Cr₂O₇)

Preparation :

The chromite ore is roasted with sodium carbonate in presence of air in a reverberatory furnace.

The roasted mass is extracted with water when Na₂CrO₄ goes into the solution leaving behind
insoluble Fe₂O³. The solution is then treated with calculated amount of H₂SO₄.

The solution is concentrated when less soluble Na₂SO₄ crystallises out. The solution is further concentrated when crystals of Na₂Cr₂O₇ are obtained. Hot saturated solution of Na₂Cr₂O₇ is then treated with KCl when orange red crystals of K₂Cr₂O₇ are obtained on crystallisation. Sodium dichromate is more soluble than potassium dichromate.

Na₂Cr₂O₇ + 2KCI  ¾® K₂Cr₂O₇ + 2 NaCl

K₂Cr₂O₇ is preferred over Na₂Cr₂O₇ as a primary standard in volumetric estimation because Na₂Cr₂O₇ is

hygroscopic in nature but K₂Cr₂O₇ is not.

Properties

(a) Physical :

It is orange-red coloured crystalline compound. It is moderately soluble in cold water but freely soluble in hot water. It melts at 398°C.

Structure of Chromate and Dichromate ion

(b) Chemical :

(i) Effect of heating :

On heating strongly, it decomposes liberating oxygen.

2K₂Cr₂O₇ ¾® 2K₂CrO₄ + Cr₂O₃ + O₂

On heating with alkalies, it is converted to chromate, i.e., the colour changes from orange to yellow. On acidifying, yellow colour again changes to orange.

K₂Cr₂O₇ + 2KOH  ¾® 2K₂CrO₄ + H₂O

Cr₂O₇^2-(Orange) + 2OH¯  ¾®  2CrO₄^2- (Yellow)+ H₂O

2CrO(Yellow) + 2H⁺ ¾® Cr₂O₇^2- (Orange) + H₂O

Thus CrO₄^2– and Cr₂O₇^2– exist in equilibrium and are interconvertable by altering the pH of solution.

2CrO₄^2-  + 2H⁺2HCrO₄^–Cr₂O₇^2- + H₂O

In alkaline solution, chromate ions are present while in acidic solution, dichromate ions are present.

(ii) K₂Cr₂O₇ + 2H₂SO₄ (conc. & cold)  ¾® 2CrO₃¯ (bright orange/red) + 2KHSO₄ + H₂O

2K₂Cr₂O₇ + 8H₂SO₄ (conc. & Hot) ¾® 2K₂SO₄ + 8H₂O + 2Cr₂(SO₄)₃ + 3O₂

(iii)   Acidified K₂Cr₂O₇ solution reacts with H₂O₂ to give a deep blue solution due to the formation of CrO₅.

Cr₂O₇^2–   +  2H+  + 4H₂O₂ ¾® 2CrO₅ + 5H₂O

Blue colour in aqueous solution fades away slowly due to the decomposition of CrO₅ to Cr³⁺ ions and oxygen. In less acidic solution K₂Cr₂O₇ and H₂O₂ give salt which is violet coloured and diamagnetic due to the formation of [CrO(O₂)(OH)]^–.

(iv)Potassium dichromate reacts with hydrochloric acid and evolves chlorine gas.

K₂Cr₂O₇ + 14HCl ¾® 2KCl + 2CrCl₃ + 7H₂O + 3Cl₂

(v) It acts as a powerful oxidising agent in acidic medium (dilute H₂SO₄)

Cr₂O₇^2- + 14H⁺ + 6e^– ¾® 2Cr⁺³ + 7H₂O. (Eº = 1.33 V)

The oxidation state of Cr changes from + 6 to +3.

(a) Iodine is liberated from potassium iodide :

(b) Ferrous salts are oxidised to ferric salts :

6Fe²⁺ + Cr₂O₇^2- + 14H⁺ ¾® 6Fe³⁺ + 2Cr³⁺ + 7H₂O

(c) Sulphites are oxidised to sulphates :

 Cr₂O₇^2- + 3SO₃^2- + 8H⁺¾® 3SO₄^2- + 2Cr³⁺ + 4H₂O

(d) H2S is oxidised to sulphur :

Cr₂O₇^2- + 3H₂S + 8H⁺ ¾® 2Cr³⁺ + 7H₂O + 3S

(e) SO₂ is oxidised to H₂SO₄ :

Cr₂O₇^2– + 3SO₂ + 2H+  ¾® 2Cr³⁺ + 3SO₄^2– + H₂O ;

Chrome alum is obtained when acidified K2Cr2O7 solution is saturated with SO2.

K₂Cr₂O₇ + H₂SO₄ + 3SO₂ + 23H₂O   K₂SO₄ . Cr₂(SO₄)₃ . 24H₂O

(f) It oxidises ethylalcohol to acetaldehyde and acetaldehyde to acetic acid

C₂H₅OH   CH₃CHO  CH₃COOH

 ethyl alcohol          acetaldehyde         acetic acid

(g) It also oxidises nitrites to nitrates, arsenites to arsenates, HBr to Br₂. HI to I₂, etc.

(h)  K₂Cr₂O₇ + 2C (charcoal)   Cr₂O₃ + K₂CO₃ + CO ↑­

(vi)     Chromyl chloride test :   

4Cl^– + Cr₂O₇^2– + 6H⁺ ¾® 2CrO₂Cl₂ ↑ (deep red) + 3H₂O

CrO₂Cl₂ + 4OH^– ¾® CrO₄^2– (yellow) + 2Cl^– + 2H₂O

CrO₄2^– (yellow) + Pb²⁺ ¾® PbCrO₄ ↓ (yellow)

(vii) Cr₂O₇^2– (concentrated solution) + 2Ag⁺ ¾® Ag₂Cr₂O₇↓ (reddish brown)

 Ag₂C_r2O₇ ↓ + H₂O   Ag₂CrO₄  ↓+ CrO₄^2– + 2H⁺ .

(viii)  Cr₂O₇^2– + Ba²⁺  + H₂O  2BaCrO₄ ↓+ 2H⁺

As strong acid is produced, the precipitation is only partial. But if NaOH or CH₃COONa is added, precipitate becomes quantitative.

Uses : It is used :

(i) as a volumetric reagent in the estimation of reducing agents such as oxalic acid, ferrous ions, iodide ions, etc. It is used as a primary standard.

(ii)  for the preparation of several chromium compounds such as chrome alum, chrome yellow, chrome red, zinc yellow, etc.

(iii) in dyeing, chrome tanning, calico printing, photography etc.

(iv) as a cleansing agent for glass ware in the form of chromic acid.

in leather industry and as an oxidant for preparation of azo compounds.

Periodic trends and properties of f-Block Elements

The Lanthanides

The names, symbols, electronic configurations of atomic and some ionic states and atomic and ionic radii of lanthanum and lanthanide (for which the general symbol Ln is used) are given in Table.

Electronic Configurations : It may be noted that atoms of these elements have electronic configuration with 6s² common but with variable occupancy of 4f level (Table). However, the electronic configurations of all the tripositive ions (the most stable oxidation state of all the lanthanides) are of the form 4fⁿ (n = 1 to 14 with increasing atomic number).

Atomic and Ionic Sizes : The overall decrease in atomic and ionic radii from lanthanum to lutetium (the lanthanide contraction). The shielding of one 4 f electron by another is less than one d electron by another with the increase in nuclear charge along the series. There is fairly regular decrease in the sizes with increasing atomic number. The cumulative effect of the contraction of the lanthanide series, known as lanthanide 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 (160 pm) and Hf (159 pm), a consequence of the lanthanide contraction.

Oxidation States : In the lanthanides, La(III) and Ln(III) compounds are predominant species. However, occasionally +2 and +4 ions in solution or in solid compounds are also obtained.

This irregularity (as in ionisation enthalpies) arises mainly from the extra stability of empty, half-filled or filled f subshell. Thus, the formation of Ce^IV is favoured by its noble gas configuration, but it is a strong oxidant reverting to the common +3 state. The E^Θ value for Ce⁴⁺/ Ce³⁺ is + 1.74 V which suggests that it can oxidise water. However, the reaction rate is very slow and hence Ce(IV) is a good analytical reagent.

Pr, Nd, Tb and Dy also exhibit +4 state but only in oxides, MO₂. Eu²⁺ is formed by losing the two s electrons and its f⁷ configuration accounts for the formation of this ion.

However, Eu²⁺ is a strong reducing agent changing to the common +3 state. Similarly Yb²⁺ which has f 14 configuration is a reductant.

Tb^IV has half-filled f-orbitals and is an oxidant. The behaviour of samarium is very much like europium, exhibiting both +2 and +3 oxidation states.

General Characteristics :

All the lanthanides are silvery white soft metals and tarnish rapidly in air.

The hardness increases with increasing atomic number, samarium being steel hard.

Their melting points range between 1000 to1200K but samarium melts at 1623 K.

They have typical metallic structure and are good conductors of heat and electricity. Density and other properties change smoothly except for Eu and Yb and occasionally for Sm and Tm.

Many trivalent lanthanide ions are coloured both in the solid state and in aqueous solutions. Colour of these ions may be attributed to the presence of f electrons. Neither La³⁺ nor Lu³⁺ ion shows any colour but the rest do so. However, absorption bands are narrow, probably because of the excitation within f level.

The lanthanide ions other than the f°type (La³⁺ and Ce⁴⁺) and the f¹⁴ type (Yb²⁺ and Lu³⁺) are all paramagnetic. The paramagnetism rises to maximum in neodymium.

The first ionisation enthalpies of the lanthanides are around 600 kJ mol^–1, the second about 1200 kJ mol^–1 comparable with those of calcium.

A detailed discussion of the variation of the third ionisation enthalpies indicates that the exchange enthalpy consideration (as in 3d orbitals of the first transition series), appear to impart a certain degree of stability to empty, half-filled and completely filled orbitals f level. This is indicated from the abnormally low value of the third ionization enthalpy of lanthanum, gadolinium and lutetium.
In their chemical behaviour, in general, the earlier members of the series are quite reactive similar to calcium but, with increasing atomic number, they behave more like aluminium.

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 Eu for which the value is – 2.0 V. This is, of course, a small variation.

The metals combine with hydrogen when gently heated in the gas.

The carbides, Ln₃C, Ln₂C₃ and LnC₂ are formed when the metals are heated with carbon.

They liberate hydrogen from dilute acids and burn in halogens to form halides.

They form oxides M₂O₃ and hydroxides M(OH)₃. The hydroxides are definite compounds, not just hydrated oxides.

They are basic like alkaline earth metal oxides and hydroxides.

The best single use of the lanthanides is for the production of alloy steels for plates and pipes. A well known alloy is mischmetall which consists of a lanthanide metal (~ 95%) and iron (~ 5%) and traces of S, C, Ca and Al. A good deal of mischmetall is used in Mg-based alloy to produce bullets, shell and lighter flint. Mixed oxides of lanthanides are employed as catalysts in petroleum cracking. Some individual Ln oxides are used as phosphors in television screens and similar fluorescing surfaces.

The Actinides

The actinides include the fourteen elements from Th to Lr. The names, symbols and some properties of these elements are given in Table.

The actinides are radioactive elements and the earlier members have relatively long half-lives, the latter ones have half-life values ranging from a day to 3 minutes for lawrencium (Z =103). The latter members could be prepared only in nanogram quantities. These facts render their study more difficult.

Electronic Configurations :

All the actinides are believed to have the electronic configuration of 7s² and variable occupancy of the 5f and 6d subshells.

The fourteen electrons are formally added to 5f, though not in thorium (Z = 90) but from Pa onwards the 5f orbitals are complete at element 103.

The irregularities in the electronic configurations of the actinides, like those in the lanthanides are related to the stabilities of the f° , f⁷ and f¹⁴ occupancies of the 5f orbitals. Thus, the configurations of Am and Cm are [Rn] 5f⁷7s² and [Rn] 5f⁷ 6d¹ 7s² .

Although the 5f orbitals resemble the 4f orbitals in their angular part of the wave-function, they are not as buried as 4f orbitals and hence 5f electrons can participate in bonding to a far greater extent.

Ionic Sizes :

The general trend in lanthanides is observable in the actinides as well. There is a gradual decrease in the size of atoms or M³⁺ ions across the series. This may be referred to as the actinide contraction (like lanthanide contraction). The contraction is, however, greater from element to element in this series resulting from poor shielding by 5f electrons.

Oxidation States :

There is a greater range of oxidation states, which is in part attributed to the fact that the 5f, 6d and 7s levels are of comparable energies. The known oxidation states of actinides are listed in Table.

The actinides show in general +3 oxidation state.

The elements, in the first half of the series frequently exhibit higher oxidation states. For example, the maximum oxidation state increases from +4 in Th to +5, +6 and +7 respectively in Pa, U and Np but decreases in succeeding elements.

The actinides resemble the lanthanides in having more compounds in +3 state than in the +4 state. However, +3 and +4 ions tend to hydrolyse.

Because the distribution of oxidation states among the actinides is so uneven and so different for the earlier and latter elements, it is unsatisfactory to review their chemistry in terms of oxidation states.

General Characteristics and Comparison with Lanthanides :

The actinide metals are all silvery in appearance but display a variety of structures. The structural variability is obtained due to irregularities in metallic radii which are far greater than in lanthanides.

The actinides are highly reactive metals, especially when finely divided. The action of boiling water on them, for example, gives a mixture of oxide and hydride and combination with most non metals takes place at moderate temperatures. Hydrochloric acid attacks all metals but most are slightly affected by nitric acid owing to the formation of protective oxide layers; alkalies have no action.

The magnetic properties of the actinides are more complex than those of the lanthanides. Although the variation in the magnetic susceptibility of the actinides with the number of unpaired 5 f electrons is roughly parallel to the corresponding results for the lanthanides, the latter have higher values.

It is evident from the behaviour of the actinides that the ionisation enthalpies of the early actinides, though not accurately known, but are lower than for the early lanthanides. This is quite reasonable since it is to be expected that when 5f orbitals are beginning to be occupied, they will penetrate less into the inner core of electrons. The 5f electrons, will therefore, be more effectively shielded from the nuclear charge than the 4f electrons of the corresponding lanthanides. Because the outer electrons are less firmly held, they are available for bonding in the actinides.

A comparison of the actinides with the lanthanides, with respect to different characteristics as discussed above, reveals that behaviour similar to that of the lanthanides is not evident until the second half of the actinide series. However, even the early actinides resemble the lanthanides in showing close similarities with each other and in gradual variation in properties which do not entail change in oxidation state. The lanthanide and actinide contractions, have extended effects on the sizes, and therefore, the properties of the elements succeeding them in their respective periods. The lanthanide contraction is more important because the chemistry of elements succeeding the actinides are much less known at the present time.

Some Applications of d- and f-Block Elements

Iron and steels are the most important construction materials. Their production is based on the reduction of iron oxides, the removal of impurities and the addition of carbon and alloying metals such as Cr, Mn and Ni.

TiO for the pigment industry and MnO₂ for use in dry battery cells. The battery industry also requires Zn and Ni/Cd.

The ‘silver’ UK coins are a Cu/Ni alloy.