Different metals

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INTRODUCTION

Different metals get ionised at different rates. For e.g., metals like sodium when exposed to air combine almost instantly with the oxygen present therein (sodium atom gives up an electron almost as soon as it is in contact with oxygen). On the other hand, metals like gold if exposed to air even for a very long period, do not react with air. Even if gold is dropped into an acid, it remains unaffected.

Based on the ease with which metals lose their electrons and form their ions, the metals are also arranged in a series called Metal Activity Series. Metals that ionise most easily are placed at the top of the Metal Activity Series, and those that ionise least easily are placed at the lower most end.

Most of the elements of the periodic table can be arranged in such a fashion, which reflects their order of activity. This arrangement of elements in order of their increasing rates of ionisation i.e. oxidising and reducing strength, is also called the activity series or the electrochemical series.

ELECTROCHEMICAL SERIES

Half-reaction

E°(V)

Ref.

N2(g) + H++e−HN3(aq)

−3.09

Li++e−Li(s)

−3.0401

N2(g) + 4H2O + 2e−2NH2OH(aq) + 2OH−

−3.04

Cs++e−Cs(s)

−3.026

Rb++e−Rb(s)

−2.98

K++e−K(s)

−2.931

Ba2++ 2e−Ba(s)

−2.912

La(OH)3(s) + 3e−La(s) + 3OH−

−2.90

Sr2++ 2e−Sr(s)

−2.899

Ca2++ 2e−Ca(s)

−2.868

Eu2++ 2e−Eu(s)

−2.812

Ra2++ 2e−Ra(s)

−2.8

Na++e−Na(s)

−2.71

La3++ 3e−La(s)

−2.379

Y3++ 3e−Y(s)

−2.372

Mg2++ 2e−Mg(s)

−2.372

ZrO(OH)2(s) + H2O + 4e−Zr(s) + 4OH−

−2.36

Al(OH)4−+ 3e−Al(s) + 4OH−

−2.33

Al(OH)3(s) + 3e−Al(s) + 3OH−

−2.31

H2(g) + 2e−2H−

−2.25

Ac3++ 3e−Ac(s)

−2.20

Be2++ 2e−Be(s)

−1.85

U3++ 3e−U(s)

−1.66

Al3++ 3e−Al(s)

−1.66

Ti2++ 2e−Ti(s)

−1.63

ZrO2(s) + 4H++ 4e−Zr(s) + 2H2O

−1.553

Zr4++ 4e−Zr(s)

−1.45

TiO(s) + 2H++ 2e−Ti(s) + H2O

−1.31

Zn(OH)42−+ 2e−Zn(s) + 4OH-

−1.199

Ti2O3(s) + 2H++ 2e−2TiO(s) + H2O

−1.23

Ti3++ 3e−Ti(s)

−1.21

Mn2++ 2e−Mn(s)

−1.185

Te(s) + 2e−Te2−

−1.143

V2++ 2e−V(s)

−1.13

Nb3++ 3e−Nb(s)

−1.099

Sn(s) + 4H++ 4e−SnH4(g)

−1.07

SiO2(s) + 4H++ 4e−Si(s) + 2H2O

−0.91

B(OH)3(aq) + 3H++ 3e−B(s) + 3H2O

−0.89

TiO2++ 2H++ 4e−Ti(s) + H2O

−0.86

Bi(s) + 3H++ 3e−BiH3

−0.8

2H2O+ 2e−H2(g) + 2OH−

−0.8277

Zn2++ 2e−Zn(Hg)

−0.7628

Zn2++ 2e−Zn(s)

−0.7618

Ta2O5(s) + 10H++ 10e−2Ta(s) + 5H2O

−0.75

Cr3++ 3e−Cr(s)

−0.74

[Au(CN)2]−+e−Au(s) + 2CN−

−0.60

Ta3++ 3e−Ta(s)

−0.6

PbO(s) + H2O + 2e−Pb(s) + 2OH−

−0.58

2TiO2(s) + 2H++ 2e−Ti2O3(s) + H2O

−0.56

Ga3++ 3e−Ga(s)

−0.53

AgI(s) +e−Ag(s) + I−

−0.15224

U4++e−U3+

−0.52

H3PO2(aq) + H++e−P(white[note 1]) + 2H2O

−0.508

H3PO3(aq) + 2H++ 2e−H3PO2(aq) + H2O

−0.499

H3PO3(aq) + 3H++ 3e−P(red)[note 1]+ 3H2O

−0.454

Fe2++ 2e−Fe(s)

−0.44

2CO2(g) + 2H++ 2e−HOOCCOOH(aq)

−0.43

Cr3++e−Cr2+

−0.42

Cd2++ 2e−Cd(s)

−0.40

GeO2(s) + 2H++ 2e−GeO(s) + H2O

−0.37

Cu2O(s) + H2O + 2e−2Cu(s) + 2OH−

−0.360

PbSO4(s) + 2e−Pb(s) + SO42−

−0.3588

PbSO4(s) + 2e−Pb(Hg) + SO42−

−0.3505

Eu3++e−Eu2+

−0.35

In3++ 3e−In(s)

−0.34

Tl++e−Tl(s)

−0.34

Ge(s) + 4H++ 4e−GeH4(g)

−0.29

Co2++ 2e−Co(s)

−0.28

H3PO4(aq) + 2H++ 2e−H3PO3(aq) + H2O

−0.276

V3++e−V2+

−0.26

Ni2++ 2e−Ni(s)

−0.25

As(s) + 3H++ 3e−AsH3(g)

−0.23

MoO2(s) + 4H++ 4e−Mo(s) + 2H2O

−0.15

Si(s) + 4H++ 4e−SiH4(g)

−0.14

Sn2++ 2e−Sn(s)

−0.13

O2(g) + H++e−HO2•(aq)

−0.13

Pb2++ 2e−Pb(s)

−0.13

WO2(s) + 4H++ 4e−W(s) + 2H2O

−0.12

P(red) + 3H++ 3e−PH3(g)

−0.111

CO2(g) + 2H++ 2e−HCOOH(aq)

−0.11

Se(s) + 2H++ 2e−H2Se(g)

−0.11

CO2(g) + 2H++ 2e−CO(g) + H2O

−0.11

SnO(s) + 2H++ 2e−Sn(s) + H2O

−0.10

SnO2(s) + 2H++ 2e−SnO(s) + H2O

−0.09

WO3(aq) + 6H++ 6e−W(s) + 3H2O

−0.09

P(white) + 3H++ 3e−PH3(g)

−0.063

HCOOH(aq) + 2H++ 2e−HCHO(aq) + H2O

−0.03

2H++ 2e−H2(g)

0.0000

≡0

AgBr(s) +e−Ag(s) + Br−

+0.07133

S4O62−+ 2e−2S2O32−

+0.08

Fe3O4(s) + 8H++ 8e−3Fe(s) + 4H2O

+0.085

N2(g) + 2H2O + 6H++ 6e−2NH4OH(aq)

+0.092

HgO(s) + H2O + 2e−Hg(l) + 2OH−

+0.0977

Cu(NH3)42++e−Cu(NH3)2++ 2NH3

+0.10

Ru(NH3)63++e−Ru(NH3)62+

+0.10

N2H4(aq) + 4H2O + 2e−2NH4++ 4OH−

+0.11

H2MoO4(aq) + 6H++ 6e−Mo(s) + 4H2O

+0.11

Ge4++ 4e−Ge(s)

+0.12

C(s) + 4H++ 4e−CH4(g)

+0.13

HCHO(aq) + 2H++ 2e−CH3OH(aq)

+0.13

S(s) + 2H++ 2e−H2S(g)

+0.14

Sn4++ 2e−Sn2+

+0.15

Cu2++e−Cu+

+0.159

HSO4−+ 3H++ 2e−SO2(aq) + 2H2O

+0.16

UO22++e−UO2+

+0.163

SO42−+ 4H++ 2e−SO2(aq) + 2H2O

+0.17

TiO2++ 2H++e−Ti3++ H2O

+0.19

SbO++ 2H++ 3e−Sb(s) + H2O

+0.20

AgCl(s) +e−Ag(s) + Cl−

+0.22233

H3AsO3(aq) + 3H++ 3e−As(s) + 3H2O

+0.24

GeO(s) + 2H++ 2e−Ge(s) + H2O

+0.26

UO2++ 4H++e−U4++ 2H2O

+0.273

Re3++ 3e−Re(s)

+0.300

Bi3++ 3e−Bi(s)

+0.308

VO2++ 2H++e−V3++ H2O

+0.34

Cu2++ 2e−Cu(s)

+0.340

[Fe(CN)6]3−+e−[Fe(CN)6]4−

+0.36

O2(g) + 2H2O + 4e−4OH−(aq)

+0.40

H2MoO4+ 6H++ 3e−Mo3++ 2H2O

+0.43

CH3OH(aq) + 2H++ 2e−CH4(g) + H2O

+0.50

SO2(aq) + 4H++ 4e−S(s) + 2H2O

+0.50

Cu++e−Cu(s)

+0.520

CO(g) + 2H++ 2e−C(s) + H2O

+0.52

I2(s) + 2e−2I−

+0.54

I3−+ 2e−3I−

+0.53

[AuI4]−+ 3e−Au(s) + 4I−

+0.56

H3AsO4(aq) + 2H++ 2e−H3AsO3(aq) + H2O

+0.56

[AuI2]−+e−Au(s) + 2I−

+0.58

MnO4−+ 2H2O + 3e−MnO2(s) + 4OH−

+0.59

S2O32−+ 6H++ 4e−2S(s) + 3H2O

+0.60

H2MoO4(aq) + 2H++ 2e−MoO2(s) + 2H2O

+0.65

+ 2H++ 2e−

+0.6992

O2(g) + 2H++ 2e−H2O2(aq)

+0.70

Tl3++ 3e−Tl(s)

+0.72

PtCl62−+ 2e−PtCl42−+ 2Cl−

+0.726

H2SeO3(aq) + 4H++ 4e−Se(s) + 3H2O

+0.74

PtCl42−+ 2e−Pt(s) + 4Cl−

+0.758

Fe3++e−Fe2+

+0.77

Ag++e−Ag(s)

+0.7996

Hg22++ 2e−2Hg(l)

+0.80

NO3−(aq) + 2H++e−NO2(g) + H2O

+0.80

[AuBr4]−+ 3e−Au(s) + 4Br−

+0.85

Hg2++ 2e−Hg(l)

+0.85

MnO4−+ H++e−HMnO4−

+0.90

2Hg2++ 2e−Hg22+

+0.91

Pd2++ 2e−Pd(s)

+0.915

[AuCl4]−+ 3e−Au(s) + 4Cl−

+0.93

MnO2(s) + 4H++e−Mn3++ 2H2O

+0.95

[AuBr2]−+e−Au(s) + 2Br−

+0.96

Br2(l) + 2e−2Br−

+1.066

Br2(aq) + 2e−2Br−

+1.0873

IO3−+ 5H++ 4e−HIO(aq) + 2H2O

+1.13

[AuCl2]−+e−Au(s) + 2Cl−

+1.15

HSeO4−+ 3H++ 2e−H2SeO3(aq) + H2O

+1.15

Ag2O(s) + 2H++ 2e−2Ag(s) + H2O

+1.17

ClO3−+ 2H++e−ClO2(g) + H2O

+1.18

Pt2++ 2e−Pt(s)

+1.188

ClO2(g) + H++e−HClO2(aq)

+1.19

2IO3−+ 12H++ 10e−I2(s) + 6H2O

+1.20

ClO4−+ 2H++ 2e−ClO3−+ H2O

+1.20

O2(g) + 4H++ 4e−2H2O

+1.23

MnO2(s) + 4H++ 2e−Mn2++ 2H2O

+1.23

Tl3++ 2e−Tl+

+1.25

Cl2(g) + 2e−2Cl−

+1.36

Cr2O72−+ 14H++ 6e−2Cr3++ 7H2O

+1.33

CoO2(s) + 4H++e−Co3++ 2H2O

+1.42

2NH3OH++ H++ 2e−N2H5++ 2H2O

+1.42

2HIO(aq) + 2H++ 2e−I2(s) + 2H2O

+1.44

Ce4++e−Ce3+

+1.44

BrO3−+ 5H++ 4e−HBrO(aq) + 2H2O

+1.45

β-PbO2(s) + 4H++ 2e−Pb2++ 2H2O

+1.460

α-PbO2(s) + 4H++ 2e−Pb2++ 2H2O

+1.468

2BrO3−+ 12H++ 10e−Br2(l) + 6H2O

+1.48

2ClO3−+ 12H++ 10e−Cl2(g) + 6H2O

+1.49

MnO4−+ 8H++ 5e−Mn2++ 4H2O

+1.51

HO2•+ H++e−H2O2(aq)

+1.51

Au3++ 3e−Au(s)

+1.52

NiO2(s) + 4H++ 2e−Ni2++ 2OH−

+1.59

2HClO(aq) + 2H++ 2e−Cl2(g) + 2H2O

+1.63

Ag2O3(s) + 6H++ 4e−2Ag++ 3H2O

+1.67

HClO2(aq) + 2H++ 2e−HClO(aq) + H2O

+1.67

Pb4++ 2e−Pb2+

+1.69

MnO4−+ 4H++ 3e−MnO2(s) + 2H2O

+1.70

H2O2(aq) + 2H++ 2e−2H2O

+1.78

AgO(s) + 2H++e−Ag++ H2O

+1.77

Co3++e−Co2+

+1.82

Au++e−Au(s)

+1.83

BrO4−+ 2H++ 2e−BrO3−+ H2O

+1.85

Ag2++e−Ag+

+1.98

S2O82−+ 2e−2SO42−

+2.010

O3(g) + 2H++ 2e−O2(g) + H2O

+2.075

HMnO4−+ 3H++ 2e−MnO2(s) + 2H2O

+2.09

F2(g) + 2e−2F−

+2.87

F2(g) + 2H++ 2e−2HF(aq)

+3.05

Oxidation / reduction and the electrochemical series

Oxidation and reduction in terms of electron transfer

Reducing agents and oxidising agents

A reducing agent reduces something else. That must mean that it gives electrons to it.

Magnesium is good at giving away electrons to form its ions. Magnesium must be a good reducing agent.

An oxidising agent oxidises something else. That must mean that it takes electrons from it.

Copper doesn't form its ions very readily, and its ions easily pick up electrons from somewhere to revert to metallic copper. Copper(II) ions must be good oxidising agents.

Metals at the top of the series are good at giving away electrons. They are good reducing agents. The reducing ability of the metal increases as you go up the series.

Metal ions at the bottom of the series are good at picking up electrons. They are good oxidising agents. The oxidising ability of the metal ions increases as you go down the series.

The more negative the E° value, the more the position of equilibrium lies to the left - the more readily the metal loses electrons. The more negative the value, the stronger reducing agent the metal is.

The more positive the E° value, the more the position of equilibrium lies to the right - the less readily the metal loses electrons, and the more readily its ions pick them up again. The more positive the value, the stronger oxidising agent the metal ion is.

REACTIVITY SERIES

Thereactivity seriesoractivity seriesis an empirical series ofmetals, in order of "reactivity" from highest to lowest.[7][8][9]It is used to summarize information about the reactions of metals withacidsandwater,single displacement reactionsand the extraction of metals from theirores.

Going from bottom to top, the metals:

  • increase in reactivity;
  • lose electrons more readily to form positive ions;
  • corrode or tarnish more readily;
  • require more energy (and different methods) to be separated from their ores
  • become stronger reducing agents.

Reaction with water and acids

The most reactive metals (for example,sodium) will react with hot water to producehydrogenand the metalhydroxide:

2Na(s) + 2H2O(l) → 2NaOH(aq) +H2(g)

Metals in the middle of the reactivity series (for example,lead) will react with acids, but not water, to give hydrogen and a metalsalt:

Fe(s) +H2SO4(aq) →FeSO4(aq) + H2(g)

There is some ambiguity at the borderlines between the groups.Magnesium,aluminiumandzinccanreact with water, but the reaction is usually very slow unless the metal samples are specially prepared to remove the surface layer of oxide which protects the rest of the metal.Copperandsilverwill react withnitric acid, but not by the simple equation shown for iron.

Single displacement reactions

An iron nail placed in a solution ofcopper sulfatewill quickly change colour as metallic copper is deposited. The iron is converted intoiron(II) sulfate:

Fe(s) +CuSO4(aq) →Cu(s) +FeSO4(aq)

In general, a metal can displace any of the metals which are lower in the reactivity series: the higher metalreducesthe ions of the lower metal. This is used in thethermite reactionfor preparing small quantities of metallic iron, and in theKroll processfor preparingtitanium(Ti comes at about the same level as Al in the reactivity series).

Al(s) +Fe2O3(s) → Fe (s) +Al2O3(s)

2Mg(s) +TiCl4(l) →Ti(s) + 2MgCl2(s)

However, other factors can come into play, as in the preparation of metallicpotassiumby the reduction ofpotassium chloridewith sodium at 850ºC: although sodium is lower than potassium in the reactivity series, the reaction can proceed because potassium is more volatile, and is preferentially distilled off from the mixture.

Na(g) +KCl(l) →K(g) +NaCl(l)

Comparison with standard electrode potentials

The reactivity series is sometimes quoted in the strict reverse order ofstandard electrode potentials, when it is also known as the "electrochemical series":

Li > K > Sr > Ca > Na > Mg > Al > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H > Cu > Ag > Hg > Pt > Au

The positions oflithiumandsodiumare changed on such a series: gold and platinum are also inverted, although this has little practical significance as both metals are highly unreactive.

Standard electrode potentials offer a quantitative measure of the power of a reducing agent, rather than the qualitative considerations of other reactivity series. However, they are only valid forstandardconditions: in particular, they only apply to reactions in aqueous solution. Even with this proviso, the electrode potentials of lithium and sodium - and hence their positions in the electrochemical series - appear anomalous. The order of reactivity, as shown by the vigour of the reaction with water or the speed at which the metal surface tarnishes in air, appears to be

potassium > sodium > lithium > alkaline earth metals,

the same as the reverse order of the (gas-phase)ionisation energies. This is borne out by the extraction of metallic lithium by the electrolysis of aeutecticmixture oflithium chlorideandpotassium chloride: lithium metal is formed at the cathode, not potassium.[10]

Asaltis a chemical containing ametal ionand anegative ionbonded together. The metal ions might consist of copper, sodium or zinc etc. The negative ions could be sulphate, chloride or oxide etc.

Some metals are stronger than others. If a strong metal is mixed with a salt from a weaker metal, the strong metal grabs the negative ion from the weaker one. Here is an example:

  1. Lead is stronger than Gold, so if Lead is heated with Gold chloride, it reacts to give Lead chloride and Gold
  2. Calcium is stronger than Zinc, so if Calcium is heated with Zinc chloride, it reacts to give Calcium chloride and Zinc
  3. Iron is stronger than Silver, so if Iron is heated with Silver sulphide, it reacts to give Iron sulphide and Silver
  4. Zinc is stronger than Lead, so if Zinc is heated with Lead phosphate, it reacts to give Zinc phosphate and Lead
  5. Sodium is stronger than Iron, so if Sodium is heated with Iron bromide, it reacts to give Sodium bromide and Iron

Features of the electrochemical series (Activity series)

In the electrochemical series the elements that are lower in the series get discharged (lose their charge to become neutral) more easily than the ones above them.

Hydrogen is also included as a reference point in the series.

The electropositive power and the reducing power of the elements regularly decrease downwards while the electronegative power and the oxidizing power of the elements regularly increase upwards.

Significance of the electrochemical series (Activity series)

This series is an important tool that helps in predicting many electrochemical reactions.

All metals placed above hydrogen will displace hydrogen from acids while those below it do not displace hydrogen from acids.

Elements with high electropositive or electronegative power are highly reactive elements.

Each element in the series will displace any other element below it from a solution of its salt. For e.g., when we add zinc turnings in copper sulphate solution, copper is replaced by zinc because zinc is in higher position as compared to copper.

In the replacement of one metal ion from its solution by another metal, the element that gives up electrons most easily to become an ion will be in the solution. This is because it will accept the electrons back with greatest difficulty.

Application of electrochemical series

  1. Higher the SRP, greater is the tendency to accept e-, higher is the tendency to get reduced and greater is the oxidizing power. Fluorine system (F2/F)-has the highest SRP and hence it possesses highest oxidising power and this increases down the group.
  2. Lower the SRP, lesser is the tendency to accept e-, higher is the tendency to donate electrons, higher is the tendency to get oxidized and greater is the reducing power.
  3. In the ECS, Lithium has the lowest SRP. Hence, it has the highest reducing power and this goes on decreasing down the series.
  4. Higher the SRP, greater is the tendency to accept e-to form anions, higher is the electro negative nature. In the ECS, Fluorine system has the highest SRP and hence it is most electro negative and this goes on increasing down the series.
  5. Lower the SRP, lesser is the tendency to accept the electrons, greater is the tendency to donate the e-to form cations and higher is the electro positive nature. In the ECS, Li system has the lowest SRP and is highly electro positive which goes on decreasing down the series.
  6. A metal system occurring above H2, displaces H2from dilute acids, from water steam depending on its position and gets tarnished.
  7. A metal system occurring below H2does not displace H2from dilute acids, water or steam does not get tarnished.

In general, a metal system with highest SRP, has highest oxidizing power, more electro negative, gets displaced by all other system above it and always acts as positive electrode.

A metal system with lowest SRP has highest reducing power, more electro positive, displaces all other systems below it and always acts as negative electrode.

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