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transition metal chemistry of copper complexes oxidation states +1 +2 +3 redox chemical reactions physical properties advanced inorganic chemistry of copper

3d block revision notes Transition metal chemistry of copper

for Advanced Level Inorganic Chemistry students

hexaaquacopper(II) complex ion Cu2+(aq) [Cu(H2O)6]2+ blue octahedral complex

Periodic Table - Transition Metals - 3d block Copper Chemistry - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 10. Transition Metals 3d–block:   

10.11 Copper Chemistry

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The chemistry of copper is dominated by the +2 oxidation state, e.g. copper(II) complex ions, but there is also a substantial chemistry of the +1 oxidation state and copper(I) compounds can be stabilised by particular ligands

The chemistry of the principal oxidation states of copper are described in terms of the redox reactions of copper, explaining the ligand substitution displacement reactions of copper complex ions, balanced equations of copper chemistry, formula of copper complex ions with water, ammonia and chloride ion, shapes and colours of copper complexes, formula of copper(I) and copper(II) compounds

See also the absorption spectra and colours of copper compounds   *   [WEBSITE SEARCH BOX]

10.11. Chemistry of Copper Cu, Z=29, 1s22s22p63s23p63d104s1 

Data comparison of copper with the other members of the 3d–block and transition metals

Z and symbol 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn
property\name scandium titanium vanadium chromium manganese iron cobalt nickel copper zinc
melting point/oC 1541 1668 1910 1857 1246 1538 1495 1455 1083 420
density/gcm–3 2.99 4.54 6.11 7.19 7.33 7.87 8.90 8.90 8.92 7.13
atomic radius/pm 161 145 132 125 124 124 125 125 128 133
M2+ ionic radius/pm na 90 88 84 80 76 74 72 69 74
M3+ ionic radius/pm 81 76 74 69 66 64 63 62 na na
common oxidation states +3 only +2,3,4 +2,3,4,5 +2,3,6 +2,3,4,6,7 +2,3,6 +2,3 +2,+3 +1, +2, +3 +2 only
outer electron config. 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
Electrode pot'l M(s)/M2+(aq) na –1.63V –1.18V –0.90V –1.18V –0.44V –0.28V –0.26V +0.34V –0.76V
Electrode pot'l M(s)/M3+(aq) –2.03V –1.21V –0.85V –0.74V –0.28V –0.04V +0.40 na na na
Elect. pot. M2+(aq)/M3+(aq) na –0.37V –0.26V –0.42V +1.52V +0.77V +1.87V na na na

Elect. pot. = standard electrode potential data for copper (EØ at 298K/25oC, 101kPa/1 atm.)

na = data available or not applicable to copper (less common oxidation state of copper)

Extended data table for COPPER

property of copper/unit value for Cu
melting point Cu/oC 1083
boiling point Cu/oC 2567
density Cu/gcm–3 8.92
1st Ionisation Energy Cu/kJmol–1 745
2nd IE/kJmol–1 1958
3rd IE/kJmol–1 3554
4th IE/kJmol–1 5326
5th IE/kJmol–1 7709
atomic radius Cu/pm 128
Cu2+ ionic radius/pm 69
Relative polarising power Cu2+ ion 2.9
oxidation states of Cu, less common/stable +1, +2, +3
simple electron configuration of Cu 2,8,18,1
outer electrons of Cu [beyond argon core] [Ar]3d104s1
Electrode potential Cu(s)/Cu2+(aq) +0.34V
Electronegativity of Cu 1.90

There is an apparent anomaly in the electron configuration for copper

Cu is [Ar]3d104s1 and not [Ar]3d94s2

because a fully–filled 3d sub–shell seems to be a little lower in energy, and marginally more stable.


  • copper1Uses of COPPER

    • Copper is an attractive orange–reddish coloured metal which is very ductile and malleable.

    • Copper's thermal and electrical conductivities are second only to silver, but not cheap, but cheaper than silver!

    • Copper is a relatively unreactive metal and is only slowly oxidised by moist air.

    • Copper is an important metal in many alloys e.g. brass (with zinc), bronze (with tin) and coinage metals (with nickel).

    • Copper is widely used for electrical circuits because of its excellent conducting properties and is malleable enough to easily drawn into thin wire.

    • Copper is also used for piping in plumbing, again due to its convenient malleability.

    • Copper compounds are used as catalysts in the chemical industry e.g. copper(I) chloride, CuCl, is used in the manufacture of chlorobenzene.

    • Copper(II) oxide, CuO, is used in paints and copper(II) chloride, CuCl2, in fungicides.

  • Biological role of copper

    • Copper is an essential trace element and has a role in the formation of haemoglobin.

    • It is a constituent and activator of several enzymes (in plants too) such as ascorbic acid oxidase and lactase.

    • Deficiency in copper leads to anaemia and bone disorders.

  • There are brief notes on copper extraction and purification on the GCSE/IGCSE Extraction of Metals page.

The Chemistry of COPPER

Some basic reactions of copper metal, including concentrated acids, described on GCSE Reactivity Series of Metals Notes

Pd s block d blocks (3d block copper) and f blocks of metallic elements p block elements
Gp1 Gp2 Gp3/13 Gp4/14


2 3Li 4Be Part of the modern Periodic Table of Elements: ZSymbol, z = atomic or proton number

Sc to Zn are now considered the head-top elements of groups 3 to 12

3d block of metallic elements: Scandium to Zinc focus on copper

5B 6C
3 11Na 12Mg 13Al 14Si
4 19K 20Ca 21Sc







 [Ar] 3d34s2



[Ar] 3d54s1



   [Ar]   3d54s2



[Ar] 3d64s2



[Ar] 3d74s2



[Ar] 3d84s2



[Ar] 3d104s1



[Ar] 3d104s2


31Ga 32Ge
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn
6 55Cs 56Ba 57,58-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb
7 87Fr 88Ra 89,90-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Nh 114Fl

Summary of oxidation states of the 3d block metals (least important) Ti to Cu are true transition metals

Sc Ti V Cr Mn Fe Co Ni Cu Zn
                +1  (3d10)  
  (+2) (+2) (+2) +2 +2 +2 +2 +2  (3d9) +2
+3 +3 +3 +3 (+3) +3 +3 (+3) (+3) (3d8)  
  +4 +4   +4     (+4)    
      +6 (+6) (+6)        
3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
The outer electron configurations beyond [Ar] and the (ground state of the simple ion)

Note that when 3d block elements form ions, the 4s electrons are 'lost' first.

The oxidation states and electron configuration of copper in the context of the 3d block of elements

electrode potential chart diagram for the ions and oxidation states of copper + copper(I) Cu+ copper(II) Cu2+ +2

The electrode potential chart highlights the values for various oxidation states of copper.

The electrode potentials involving copper ions correspond to hydrated complex ions where the ligands are water, oxide or hydroxide.

As you can see from the chart, changing either the ligand or the oxidation state, will also change the electrode potential for that half-reaction involving a copper ion.


  • Electron configuration of Cu2+ is [Ar]3d9

  • When copper(II) salts are dissolved in water the blue tetraaquacopper(II) ion or the hexaaquacopper(II) ion is formed.

    • The scope for a variety of coloured compounds arises from the fundamental electronic configuration of the Cu2+ ion, namely [Ar]3d9, giving an incompletely filled 3d sub–shell – criteria for being a true transition metal.

      • ie there is at least one electron that can be promoted to a higher level when the 3d sub–shell is split when the central metal ion interacts with the ligands.

        • Visible light photons absorbed, colour results!

      • The Cu2+ components in the diagrams below illustrates the point.

      • For more details see Appendix 4. Electron configuration & complex ion colour theory

        • Bottom left shows the ground state of the copper(II) ion in an octahedral complex ion

  • Both the octahedral hexaaquacopper(II) ion [Cu(H2O)6]2+ and the square planar tetraaquacopper(II) ion

    • tetraaquacopper(II) complex ion [Cu(H2O)4]2+ square planar shape[Cu(H2O)4]2+ both exist, the latter is present in the blue copper(II) sulfate pentahydrate crystals of CuSO4.5H2O

      • On the near right is the copper(II) ion with a co-ordination number of 4 from 4 unidentate ligands and a square planar shape.

      • On the far right is the octahedral hydrated copper(II) ion with a co-ordination number of 6 (the number of ligand bonds) from 6 unidentate ligands.

      • (see transition metals Appendix 1 for more details on the structure of the pentahydrate crystals).

      • Solutions of copper(II) sulfate CuSO4(aq) are suitable for laboratory experiments to investigate the chemistry of the aqueous copper(II) ion.

      • I've used the hexaaquacopper(II) ion in equations involving expression of complex ion changes.

  • When alkaline aqueous ammonia or sodium hydroxide is added to a blue hexa–aqua copper(II) ion solution, initially a gelatinous palish blue precipitate of the hydroxide is formed.

    • Note it can be 4 or 6 H2O in the complex ion Cu2+(aq) i.e. [Cu(H2O)4]2+(aq)   ... or ...

    • (i)  [Cu(H2O)6]2+(aq) + 2OH(aq) ===> [Cu(H2O)4(OH)2](s) + 2H2O(l) 

      • or more simply (ii)  Cu2+(aq) + 2OH(aq) ===> Cu(OH)2(s)

      • precipitation reaction.

      • The reaction involves a ligand exchange (ligand displacement, hydroxide ion for water).

      • The copper complexes both have an octahedral shape and a co-ordination number of 6 from 6 unidentate ligands.

      • There is no change in oxidation state of copper (+2) but the overall charge on the copper complex changes from 2+ to zero (+2 + 2x-1).

      • When you add quite concentrated sodium hydroxide (or potassium hydroxide) solution to the blue gelatinous copper(II) hydroxide precipitate the blue colour deepens, so something is happening!

        • This change in colour suggest some ligand changes have taken place.

        • In fact a very small amount of the tetrahydroxocuprate(II) complex anion is formed, i.e. a tiny amount of the hydroxide precipitate dissolves, but overall copper(II) hydroxide is effectively insoluble even in strongly alkaline solutions.

        • This is why most textbooks will say the copper(II) hydroxide precipitate is insoluble in excess alkali.

        • The formation of a small amount of the soluble deeper blue complex anion can be expressed as ...

        • Cu(OH)2(s)  +  2OH(aq)  ===>  [Cu(OH)2(H2O)2]2–(aq)

        • or more correctly, since it is a ligand displacement reaction

        • [Cu(H2O)4(OH)2](s)  +  2OH(aq)  ===>  [Cu(OH)4(H2O)2]2–(aq)  +  2H2O(l) 

        • Both copper(II) complexes are octahedral, co-ordination number 6 (6 unidentate ligands), overall charge changes from 0 to 2- due to the two negative hydroxide ligands.

      • This lack significant amphoteric character means copper(II) oxide is essentially a basic oxide (insoluble in water) that readily dissolves in acids to form soluble copper(II) salts e.g. the formation of copper(II) sulfate, copper(II) chloride and copper(II) nitrate from the appropriate mineral acid solution.

        • CuO(s) + H2SO4(aq) ===> CuSO4(aq) + H2O(l)

        • CuO(s) + 2HCl(aq) ===> CuCl2(aq) + H2O(l)

        • CuO(s) + 2HNO3(aq) ===> Cu(NO3)2(aq) + H2O(l)

  • diaquatetraamminecopper(II) ion tetraamminediaquacopper(II) ion complex formed between ammonia and the aqueous copper(II) ion [Cu(NH3)4(H2O)2]2+(aq)Excess sodium hydroxide has no significant effect (see above), BUT, after an initial copper(II) hydroxide precipitate, with excess ammonia, a deep blue solution is formed of the ammine complex ion (ligand substitution is incomplete), the overall changes can be expressed as:

    • (i) [Cu(H2O)6]2+(aq) +  2NH3(aq) rev [Cu(H2O)4(OH)2](s) + 2NH4+(aq)

      • Initially a 'turquoise' precipitate of the hydrated copper(II) hydroxide is formed.

    • (ii) [Cu(H2O)4(OH)2](s) + 4NH3(aq) rev [Cu(NH3)4(H2O)2]2+(aq) +  2OH(aq) +  2H2O(l)

      • Then the hydroxide precipitate dissolves in excess ammonia to give the soluble deep blue complex ion (shown on the right).

      • diaquatetraamminecopper(II) ion  or  tetraamminediaquacopper(II) ion ?

      • There are other possible copper(II) ions with the ligand ammonia, depending on pH and ammonia concentration,

        • including square planar [Cu(NH3)4]2+

        • and octahedral [Cu(NH3)5OH]+, [Cu(NH3)(H2O)5]2+ 

        • and  [Cu(NH3)5H2O]2+ ... and others!

      • You can show the overall change as equation (iii)

    • (iii) [Cu(H2O)6]2+(aq) +  4NH3(aq) rev [Cu(NH3)4(H2O)2]2+(aq) + 4H2O(l)

      • Kc = [[Cu(NH3)4(H2O)2]2+(aq)] / [[Cu(H2O)6]2+(aq)] [NH3(aq)]4

        • = 1.2 x 1013 mol-4dm12

      • Sometimes shown as (iv) [Cu(H2O)6]2+(aq) +  4NH3(aq) rev [Cu(NH3)4]2+(aq) + 6H2O(l) 

        • but I think (iii) is more correct.

      • This is a ligand displacement reaction you can carry out directly by adding aqueous ammonia to pale blue copper(II) sulfate solution, the ammonia ligand displaces the water ligand.

      • Both copper(II) complexes are octahedral, co-ordination number 6, overall charge remains at 2+ because both water and ammonia ligands are electrically neutral.

      • The octahedral complex [Cu(NH3)4(H2O)2]2+ exhibits E/Z isomerism (cis/trans geometrical isomerism) and the more stable form is the E isomer (trans in 'old' terms for the two H2O molecules) shown in the diagram above.

      • or can be [Cu(H2O)4]2+(aq) + 4NH3(aq) rev [Cu(NH3)4]2+(aq) + 4H2O(l)

      • In this representation of the reaction both of these copper(II) complex ions are square planar with a co-ordination number of 4.

      • Unfortunately in many transition metals reactions, there are often several possibilities of the structure of the complex formed and many co-exist in solution depending on ion concentrations and pH!

      • All the equations show the eventual formation of the diaquatetraamminecopper(II) ion.

      • [Cu(NH3)6]2+(aq) hexaamminecopper(II) complex ion deep blue-violet octahedral complexIf you add a large excess of conc. ammonia ('0.880') to a pale blue copper(II) sulfate solution, (ignoring the intermediate hydroxide precipitate), you get a stepwise ligand exchange reaction to give a whole series of copper(II) ion ammine complexes as each water molecule is replaced by an ammonia molecules, the final product is [Cu(NH3)6]2+(aq), the hexaamminecopper(II) ion - shown on the right.

      • See also the absorption spectra and colours of copper compounds

    • Note: These are ligand exchange reactions, not a redox change, co–ordination number remains at 6, both octahedral complexes, both ligands electrically neutral so the overall charge of the complex remains at 2+, both the ligands are of similar size but the substitution by ammonia is incomplete.

    • Kstab = [ [Cu(NH3)4(H2O)2]2+(aq) ] / [ [Cu(H2O)6]2+(aq) ] [ NH3 (aq) ]4

      • = 1.0 x 1012 mol–4 dm12

    • by convention the term [ H2O(l) ]4 is omitted from the equilibrium expression because water is the medium and the bulk of the solution, therefore it effectively remains constant.

    • At very high concentrations of ammonia the darker blue-violet hexaamminecopper(II) ion can be formed (shown on the right.

      • [Cu(NH3)6]2+(aq)

    • The charge on this copper(II) complex remains at 2+ and retains its octahedral shape and co-ordination number 6 from the six unidentate ligands.

    • You can form similar complexes if an excess of a primary aliphatic amine is added to copper(II) salt solution (illustrated below).

primary aliphatic ammine complexes with the copper(II) ion

  • With sodium carbonate solution, copper(II) ions gives the turquoise? precipitate of copper(II) carbonate,

    • Cu2+(aq) + CO32–(aq) ===> CuCO3(s) 

    • Its actually a basic carbonate, a mixture of the hydrated hydroxide, Cu(OH)2, and carbonate, CuCO3.

      • You can make the pure carbonate by using the less alkaline sodium hydrogencarbonate solution.

      • Cu2+(aq) + 2HCO3(aq) ===> CuCO3(s) + H2O(l) + CO2(g)

  • VIEW more on ppts. with OH, NH3 and CO32–, and complexes, if any, with excess reagent.

  • If e.g. sodium chloride or hydrochloric acid is added to copper(II) sulfate solution the pale yellow–brown tetrachlorocuprate(II) complex ion is formed (can be greenish colour due to residual blue from the original Cu2+ ion).

    • pale yellow–brown tetrachlorocuprate(II) complex ion terahedral shape [CuCl4]2+(aq)[Cu(H2O)6]2+(aq) + 4Cl(aq) rev [CuCl4]2–(aq) + 6H2O(l)

      • Kstab = [ [CuCl4]2+(aq) ] / [ [Cu(H2O)6]2+(aq) ] [Cl(aq) ]4  

        • = ? mol–4 dm12

    • This particular ligand substitution/exchange reaction involves several changes (L to R):

      • the larger chloride ion ligand leads to a change in co–ordination number from 6 to 4,

      • the complex ion shape changes from octahedral to tetrahedral or square planar (not sure?),

      • it is likely that the more bulky chloride ion (radius Cl > O) 'forces' the formation of the tetrahedral shape of this copper complex ion, rather than a square planar shaped complexes.

      • the colour of the complex changes from blue to yellow–brown (green due to residual blue),

      • the complex changes from a cationic complex ion (2+) to an anionic complex ion 2- (+2 + 4x-1).

    • There is no oxidation state change at all, copper is in the +2 state throughout the reaction.

    • This is quite a good reaction to demonstrate Le Chatelier's equilibrium principle:

      • If you dissolve copper(II) chloride in water you get a greenish–blue solution as both copper(II) complexes are present in equilibrium.

      • By adding water i.e. dilution, it shifts the equilibrium to the left, more blue.

      • Increasing the chloride ion concentration by adding hydrochloric acid or sodium chloride solution shifts the equilibrium to the right, more green ==> yellowish brown.

  • The reaction between copper(II) salts and iodide ion salts:

    • i.e. the redox reaction between the copper(II) ion and the iodide ion.

    • This is a way of preparing insoluble copper(I) iodide.

    • On mixing solutions of a copper(II) salt e.g. blue copper(II) sulfate and an iodide salt e.g. colourless potassium iodide the dark colour of iodine formation is seen. Unseen, because it is masked by the iodine, is the formation of a white copper(I) iodide precipitate. This can be made visible by adding sodium thiosulfate solution which reduces the iodine back to the colourless iodide ion.

    • Cu2+(aq) + 4I(aq) ===> 2CuI(s) + I2(aq/s)

      • In terms of oxidation states:

        • copper is reduced (+2 to +1) by electron gain by the copper(II) ion

        • iodine is oxidised (–1 to 0) by electron loss by the iodide ion.

    • 2S2O32–(aq)  +  I2(aq)  ===>  S4O62–(aq) + 2I(aq) (black/brown/blue ==> colourless)

    • This reaction between the released iodine and sodium thiosulfate can be used to estimate oxidising agents like copper(II) ions. The iodine is titrated with standardised sodium thiosulfate (e.g. 0.10 mol dm–3) using a few drops of starch solution as an indicator. Iodine gives a blue colour with starch, so, the end–point is very sharp change from the last hint of blue to colourless.

    • Copper analysis eg. in brass

      • Brass can be dissolved in acid and potassium iodide solution added.

      • The resulting iodine formed can be titrated with sodium thiosulfate using starch indicator.

      • Need more details and an example calculation.

  • Summary of some complexes–compounds & oxidation states of copper compared to other 3d–block elements


  • Electron configuration of Cu+ is [Ar]3d10

  • The colour of copper(I) compounds

    • Many copper(I) compounds and copper(I) complex ions do not show the same variety of colour you see in copper(II) compounds and complex ions..

      • The lack of scope for a variety of coloured compounds arises from the fundamental electronic configuration of the Cu+ ion, namely [Ar]3d10, giving a completely filled 3d sub–shell (identical to that of the zinc ion Zn2+, and zinc compounds and complex ions tend to be white or colourless).

      • ie there is no electron that can be promoted to a higher level when the 3d sub–shell is split when the central metal ion interacts with the ligands e.g. in an octahedral complex.

      • Bottom right (4) shows the ground state of the zinc(II) ion which is electronically identical to the copper(I) ion, and clearly, no electron can be promoted, so no absorption, no colour!

      • For more details see Appendix 4. Electron configuration & complex ion colour theory

  • Disproportionation reactions:

    • Where an element in one oxidation state simultaneously changes into two species with different oxidation states

    • If solid copper(I) oxide is dissolved in dil. sulfuric acid a pinky–brown precipitate of copper and a blue solution of copper(II) sulfate solution is obtained.

      • Cu2O(s) + H2SO4(aq) ===> Cu(s) + CuSO4(aq) + H2O(l)

        • Cu2O(s) + 2H+(aq) ===> Cu(s) + Cu2+(aq) + H2O(l)

        • Oxidation number changes: 2Cu(I) ==> Cu(0) + Cu(II)

    • If solid copper(I) sulfate is dissolved in water the observations and oxidation number changes are identical to the reaction above.

      • Cu2SO4(s) + aq ===> Cu(s) +  CuSO4(aq)

      • Cu2SO4(s) + aq ===> Cu(s) +  Cu2+(aq) + SO42–(aq)

      • Oxidation state changes: 2Cu(+1) ===> Cu(0) + Cu(+2)

    • These two reactions suggest that Cu+(aq) has no stability in aqueous media and spontaneously undergoes a redox change and an electrode potential argument predicts this potential for instability and therefore the observations.

      • Note: A chemical change in which a species in one oxidation state spontaneously and simultaneously changes into two species of different oxidation states, one higher and one lower in oxidation number, is called a disproportionation reaction.

      • The argument is as follows ....

      • (i) Cu+ + e Cu   (EØCu+/Cu = +0.52V)

      • (ii) Cu2+ + e Cu+   (EØCu2+/Cu+ = +0.15V)

      • (i) the more positive redox potential, so equation (i) represents the reduction half–cell reaction and ...

      • (ii) this half–cell equation is reversed, with the less positive potential, will represent the oxidation change.

      • EØreaction = EØreduction – EØoxidation = (+0.52) – (+0.15) = +0.37V

      • showing the disproportionation is thermodynamically feasible, i.e. EØreaction must be greater than zero.

        • ie if a copper(I) compound is potentially soluble in water, the following disproportionation reaction of the copper(I) occurs

        • (iii)  2Cu+(aq) ==> Cu2+(aq) + Cu(s)

        • This ionic equation is derived from adding together the ...

        • reduction Cu+ + e  ==>  Cu and the oxidation Cu+  ==>  Cu2+  +  e

        • that is (i) + (ii) reversed = (iii) ...

        • (i)  Cu+(aq) + e ====> Cu(s) EØ = +0.52V reduction
          (ii)  Cu+(aq) ====> Cu2+(aq) + e EØ = +0.15V oxidation
          (iii) 2Cu+(aq) ====> Cu2+(aq) + Cu(s) EØ = +0.37V redox
    • See manganese(VI) chemistry for another example of disproportionation.

  • Formation of copper(I) compounds and examples of copper(I) complexes

    • How to prepare copper(I) iodide?

    • The preparation of copper(I) iodide from copper(II) sulfate and potassium iodide solutions.

    • Copper(I) iodide is formed on mixing solutions of a soluble copper(II) salt with potassium iodide solution.

      • 2Cu2+(aq) + 4I(aq) ===> 2CuI(s) + I2(aq/s)

      • It is unfortunate, from a preparation point of view, that iodine is also formed – completely obscuring the 'white' copper(I) iodide!

      • The temporary stabilisation of a copper(I) compound is facilitated by its immediate precipitation from the aqueous media so that the disproportionation of the copper(I) ion (described above) cannot happen.

      • Copper(I) iodide, like copper(I) chloride, is white when pure, when left out in air, they will slowly oxidise to the copper(II) compound eg copper(I) chloride slowly turns green as copper(II) compounds are formed.

      • I'm not quite sure how you can isolate the copper(I) iodide from this mixture?

        • I think you can remove the iodine with sodium thiosulfate and then filter off, wash and dry the CuI. Try it? You will be lucky if its white, and, if left out in air, CuI will discolour further due to aerial oxidation.

      • If you look at the basic half−cell potential data for copper/copper ions, this reaction between soluble copper(II) salts and iodide ions to form copper(I) iodide, wouldn't seem to be possible.

      • BUT, the half−cell potential for the reaction between copper(II) ions and iodide ions to form copper(I) iodide is rarely (if ever?) quoted in textbooks or data books at pre-university level.

      • Therefore when you use the appropriate EØ data (in table below) its quite obvious why the reaction described above is perfectly feasible.

      • Half−reaction (half-cell equation and standard potential)

        EØ (V)

        Oxidant (reduced)

        Reductant (oxidised)

        Cu2+(aq) +  e



        Cu2+(aq) + 2 e



        Cu+(aq) +  e



        I3(aq) + 2 e

        3 I(aq)


        I2(s) + 2 e

        2 I(aq)


        Cu2+(aq) + I(aq) + e CuI(s) +0.86
      • EØreaction = EØreduction (most +ve half-cell) – EØoxidation (least +ve half cell), therefore ...

      • EØreaction for the formation of copper(I) iodide is (+0.86) − (+0.54) = +0.32 V, therefore very feasible!

      • Details of the argument below! (with the appropriate EØ/V)

      • (i)   2Cu2+(aq)  + 2I(aq) + 2e

        ===> 2CuI(s) +0.86 V

        (ii)   2I(aq)


        I2(s) + 2 e

        +0.54 V

        (i) + (ii) gives:   2Cu2+(aq) + 4I(aq)

        ===> 2CuI(s) + I2(s) +0.32 V
      • See below for more examples of things that happen with copper/copper(I) chemistry you might not expect!

    • Copper(I)/Cu+(aq) can be stabilised by forming complexes from suitable ligands

      • e.g. copper(I) chloride dissolves in conc. hydrochloric acid to form the stable dichlorocuprate(I) complex ion (NOT a redox reaction).

      • CuCl(s) + Cl(aq) ===> [CuCl2](aq)

        • I think this is a linear shape, co-ordination number 2.

        • As well as the dichlorocuprate(I) ion, with excess concentrated chloride ion, you can get further chloro complexes formed, trichlorcuprate(I) ion [CuCl3]2– and tetrachlorocuprate(I) ion [CuCl4]3–.

        • Not sure on the shape of the latter two, but, maybe trigonal pyramid (co-ordination number 3) and tetrahedral (co-ordination number 4).

        • The same complex ions are formed if copper metal is boiled with conc. hydrochloric acid when the redox reaction, ' surprisingly' produces hydrogen! e.g.

        • 2Cu(s) + 2H+(aq) + 4Cl(aq) ===> 2[CuCl2](aq) + H2(g)

        • When it comes to complex ion formation, the usual 'reactivity series' protocol doesn't always apply!

        • This is a redox reaction however, although the Cu2+/Cu potential is +0.34V and the Cu+/Cu potential is +0.15V, on both counts hydrogen shouldn't be formed (EØH+/H2 = 0.00V),

        • BUT the actual redox potential involved is for the [CuCl2]/Cu half–cell system and this must be <0.00V ? (couldn't find it on the internet), for the half–cell oxidation reaction:

          • Cu(s) + 2Cl(aq)  – e  rev  [CuCl2](aq),

        • AND one important principal often overlooked in reactions that don't seem at first sight to be very feasible ...

        • ... the position of the equilibrium

        • It is the formation of the chlorocuprate(I) complex ions that moves the equilibrium from left (copper) to the right (complexes).

        • How to prepare copper(I) chloride?

          • The preparation of copper(I) chloride

          • The same chlorocuprate(I) complex ions are also formed when concentrated copper(II) chloride solution is boiled with excess copper turnings.

          • Cu(s)  +  CuCl2(aq)  +  2Cl(aq) ===>  2[CuCl2](aq)

          • or more accurately and ionically:   Cu(s)  +  Cu2+(aq)  +  4Cl(aq) ===>  2[CuCl2](aq)

          • This redox reaction is the opposite of disproportionation, Cu(0) + Cu(II) ==> 2Cu(I).

          • When the solution is decanted or filtered to remove the excess copper metal, and then diluted with water the white copper(I) chloride is precipitated ...

            • [CuCl2](aq)  + aq ===>  CuCl(s)  +  Cl(aq)

          • ... and must be quickly filtered, washed, dried and sealed from air because it rapidly starts to turn greenish is air as copper(I) chloride is readily oxidised by oxygen to copper(II) compounds.

      • Copper(I) compounds dissolve in an excess of potassium cyanide solution to give the tetracyanocuprate(I) complex ion.

        • CuCl(s) + 4CN(aq) ===> [Cu(CN)4]3–(aq) +  Cl(aq)

        • This shows that you can stabilise copper(I) compounds in solution using an appropriate ligand, in this case the cyanide ion, CN.

    • How to prepare copper(I) oxide?

      • The preparation of copper(I) oxide.

      • Copper(I) oxide Cu2O is formed as a dark red–brown precipitate when an aldehyde (reducing agent) or reducing sugar reacts with Fehling's solution, a copper(II) complex with a carboxylic acid).

      • In principle the half−cell reduction is: 2Cu2+(aq) +  H2O(l) + 2e ===> Cu2O(s) + 2H+(aq)

      • Note that this copper(I) compound seems stable because it is insoluble and produced in a 'reducing environment'.

        • Over time, the copper(I) oxide slowly oxidises to black copper(II) oxide CuO.

        • It must be quickly filtered, washed and dried and sealed from the oxidising action of oxygen in air.

  • Biochemistry of Copper

    • Copper ions play a vital role in electron transport/transfer reactions in cytochrome chemistry.

    • A few details to do?

    • Explaining why horseshoe crabs have blue blood instead of red blood!

      • In many mammals such as ourselves, the oxygen carrying haemoglobin (hemoglobin) is a complex of iron and is a dark red colour, hence the red colour of blood cells.
      • Horseshoe crabs don't use haemoglobin to transport oxygen, but use haemocyanin (hemocyanin) in their blood streams, which is a complex ion based on copper, hence the blue coloured blood!
      • In the complex hemocyanin molecule, two copper(II) ions in the complex can reversibly hold one oxygen molecule between them

The original extraction of copper from copper ores

  • From copper carbonate ores ...
    • The ore can be roasted to concentrate the copper as its oxide.
    • Water is driven off and the carbonate thermally decomposed.
    • copper(II) carbonate ==> copper oxide + carbon dioxide
    • CuCO3(s) ==> CuO(s) + CO2(g)
    • The oxide can be smelted by heating with carbon (coke, charcoal) to reduce the oxide to impure copper, though this method isn't really used much these days (the 'bronze age' method archaeologically!).
    • copper(II) oxide + carbon ==> copper + carbon dioxide
    • 2CuO(s) + C(s) ==> 2Cu(s) + CO2(g)
    • The carbon acts as the reducing agent – the 'oxygen remover'.
  • From copper sulfide ores ...
    • These include chalcocite/chalcosine = copper(I) sulfide Cu2S and covellite = copper(II) sulfide CuS
      • and chalcopyrite CuFeS2. which is one of the most important ores for the extraction of copper.
        • This can be roasted in air to produce copper(I) sulfide which is roasted again in a controlled amount of air so as not to form a copper oxide (see below).
        • 2CuFeS2 +  4O2 ==> Cu2S + 3SO2 + 2FeO
    • Copper sulfide ores can be rapidly roasted in heated air enriched with oxygen to form impure copper and this extraction process is called 'flash smelting' and is the most widely used and efficient method of copper extraction.
      • Nasty sulfur dioxide gas is formed, this must be collected to avoid pollution and can be used to make sulfuric acid to help the economy of the process.
      • copper(I) sulfide + oxygen ==> copper + sulfur dioxide
        • Cu2S(s) + O2(g) ==> 2Cu(s) + SO2(g)
      • or copper(II) sulfide + oxygen ==> copper + sulfur dioxide
        • CuS(s) + O2(g) ==> Cu(s) + SO2(g)
  • It is also possible to dissolve an oxide or carbonate ore in dilute sulfuric acid and extracting copper by ....
    • (1) using electrolysis see purification by electrolysis below, or
    • (2) by adding a more reactive metal to displace it e.g. scrap iron or steel is used by adding it to the resulting copper(II) sulfate solution.
      • iron + copper(II) sulfate ==> iron(II) sulfate + copper
      • Fe(s) + CuSO4(aq) ==> FeSO4(aq) + Cu(s)
    • It is possible to spray acid onto copper ore waste and leach out the copper compounds prior to electrolysing the solution or displacing the copper with a cheap metal like iron AND this can also be achieved with the help of bacteria for particular ores – see below.
  • No industrial process is ever 100% efficient, and metal extraction processes of create lots of waste material AND, crucially, that waste may contain some of the desired metal, or indeed other potentially valuable metals.
    • So, any method that can extract the small percentages of valuable metals from waste will aid the economy of production of the main product.
    • For example, 10% of the copper produced in the US is derived from bacteria which feed of chalcopyrite CuFeS2 (this could be waste or very low grade ores ??).
      • The bacteria use the Fe2+ ion and S2– ion to obtain energy needed to live.
        • The redox chemistry of the bacteria via the oxygen from air involves the ...
        • oxidation of Fe2+ to Fe3+ and S2– to SO42–
        • reduction of O2 to H2O
        • overall an energy releasing process to sustain the bacteria.
        • We use sugar, these bacteria use chalcopyrite! – I don't think we'd like the taste!
      • The bacteria effectively break down the chalcopyrite, releasing copper(II) ions into an acid solution.
        • The optimum conditions for this 'bacterial leaching' are pH 2–3 and 20oC–55oC.
      • Extracting copper in this way is cheaper, quieter, and less polluting than conventional smelting processes.
      • BUT, it is much slower, so it is primarily being used on waste dumps by spraying dilute acid on them and the aerated water slowly percolates through rock fragments and their naturally occurring bacterial colonies.
      • The leached solution of copper(II) ions is very dilute.
        • The solution is concentrated and Cu2+ ions separated from other ions e.g. Fe3+
        • The copper is displaced using cheap scrap iron.
        • The copper is then further purified by electrolysis – described below.

The Purification of Copper by Electrolysis

(c) doc b


  • The impure copper from a smelter is cast into a block to form the positive anode. The cathode is made of previously purified copper. These are dipped into an electrolyte of copper(II) sulfate solution. 
  • When the d.c electrical current is passed through the solution electrolysis takes place.  The copper anode dissolves forming blue copper(II) ions Cu2+.
  • These positive ions are attracted to the negative cathode and become copper atoms. The mass of copper dissolving at the anode exactly equals the mass of copper deposited on the cathode. The concentration of the copper(II) sulfate remains constant.
  • Any impurities present in the impure copper anode fall to the bottom of the electrolysis cell tank. This 'anode sludge' is not completely mineral waste, it can contain valuable metals such as silver!
  • See section below on extraction of impure copper from an ore.

Raw materials for the electrolysis process:

  • Impure copper from a copper smelter.

  • Electrolyte of aqueous copper(II) sulfate.

  • A pure copper cathode.

Electrolysis is using d.c. electrical energy to bring about chemical changes at the electrolyte connections called the anode and cathode  electrodes.

An electrolyte is a conducting melt or solution of ions which carry the electric charge as part of the circuit.

Scrap copper can be recycled and purified this way too ,and is cheaper than starting from copper ore AND saves valuable mineral resources.

The redox details of the electrode processes:
  • At the positive (+) anode, the process is an oxidation, electron loss, as the copper atoms dissolve to form copper(II) ions.

Cu(s) ===> Cu2+(aq) + 2e

  • at the negative () cathode, the process is a reduction, electron gain by the attracted copper(II) ions to form neutral copper atoms.

Cu2+(aq) + 2e ===> Cu(s)

keywords redox reactions ligand substitution displacement balanced equations formula complex ions complexes ligand exchange reactions redox reactions ligands colours oxidation states: copper ions Cu(0) Cu+ Cu(+1) Cu(I) Cu2+ Cu(+2) Cu(II) Cu(+3) Cu(III) CuSO4 Cu2O CuSO4.5H2O [Cu(H2O)4]2+ [Cu(H2O)6]2+ + 2OH– ==> [Cu(H2O)4(OH)2] + 2H2O [Cu(H2O)6]2+ + 4 NH3 [Cu(NH3)4(H2O)2]2+  + 4 H2O [Cu(H2O)4]2+ + 4NH3 [Cu(NH3)4]2+ + 4H2O [Cu(H2O)4(OH)2] + 4NH3  [Cu(NH3)4(H2O)2]2+ + 2 OH– + 4H2O Cu(OH)2] + 4NH3  [Cu(NH3)4]2+ + 2OH– Kstab = [ [Cu(NH3)4(H2O)2]2+ ] / [ [Cu(H2O)6]2+ ] [ NH3  ]4 = 1.0 x 1012 mol–4 dm12 Cu2+ + CO32– ==> CuCO3  Cu2+ + 2HCO3– ==> CuCO3 + H2O + CO2 [Cu(H2O)6]2+ + 4Cl– [CuCl4]2–  + 6H2O units of Kstab = [ [CuCl4]2+ ] / [ [Cu(H2O)6]2+ ] [ Cl–  ]4  = ? mol–4 dm12 Cu2+ + 4I– ==> 2 CuI + I2 Cu2O + H2SO4 ==> Cu + CuSO4 + H2O Cu2O + 2H+ ==> Cu + Cu2+ + H2O Oxidation number changes: 2 Cu(I) ==> Cu(0) + Cu(II) Cu2SO4 + aq ==> Cu + CuSO4 Cu2SO4 + aq ==> Cu + Cu2+ + SO42– Oxidation state changes: 2Cu(+1) ==> Cu (0) + Cu (+2) 2Cu+ ==> Cu2+ +Cu 2Cu2+ + 4 I– ==> 2 CuI + I2 CuCl + Cl– ==> [CuCl2]– 2Cu + 2H+ + 4Cl– ==> 2[CuCl2]– + H2 CuCl + 4CN– ==> [Cu(CN)4]3– + Cl–  oxidation states of copper, redox reactions of copper, ligand substitution displacement reactions of copper, balanced equations of copper chemistry, formula of copper complex ions, shapes colours of copper complexes  Na2CO3 NaOH NH3 transition metal chemistry of copper for AQA AS chemistry, transition metal chemistry of copper for Edexcel A level AS chemistry, transition metal chemistry of copper for A level OCR AS chemistry A, transition metal chemistry of copper for OCR Salters AS chemistry B, transition metal chemistry of copper for AQA A level chemistry, transition metal chemistry of copper for A level Edexcel A level chemistry, transition metal chemistry of copper for OCR A level chemistry A, transition metal chemistry of copper for A level OCR Salters A level chemistry B transition metal chemistry of copper for US Honours grade 11 grade 12 transition metal chemistry of copper for pre−university chemistry courses pre−university A level revision notes for transition metal chemistry of copper  A level guide notes on transition metal chemistry of copper for schools colleges academies science course tutors images pictures diagrams for transition metal chemistry of copper A level chemistry revision notes on transition metal chemistry of copper for revising module topics notes to help on understanding of transition metal chemistry of copper university courses in science careers in science jobs in the industry laboratory assistant apprenticeships technical internships USA US grade 11 grade 11 AQA A level chemistry notes on transition metal chemistry of copper Edexcel A level chemistry notes on transition metal chemistry of copper for OCR A level chemistry notes WJEC A level chemistry notes on transition metal chemistry of copper CCEA/CEA A level chemistry notes on transition metal chemistry of copper for university entrance examinations physical and chemical properties of the 3d block transition metal copper, oxidation and reduction reactions of copper ions, outer electronic configurations of copper, principal oxidation states of copper, shapes of copper's complexes, octahedral complexes of copper, tetrahedral complexes of copper, square planar complexes of copper, stability data for copper's complexes, aqueous chemistry of copper ions, redox reactions of copper ions, physical properties of copper, melting point of copper, boiling point of copper, electronegativity of copper, density of copper, atomic radius of copper, ion radius of copper, ionic radii of copper's ions, common oxidation states of copper, standard electrode potential data for copper, ionisation energies of copper, polarising power of copper ions, industrial applications of copper compounds, chemical properties of copper compounds, why are copper complexes coloured?, isomerism in the complexes of copper, chemistry of the copper(II) ion, complexes of the copper(II) ion


GCSE Level Notes on Transition Metals (for the basics)

The chemistry of Scandium * Titanium * Vanadium * Chromium * Manganese

The chemistry of Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations Some pages have a matching sub-index

Advanced Level Inorganic Chemistry Periodic Table Index: Part 1 Periodic Table history Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr AND important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots All 11 Parts have their own sub-indexes near the top of the pages

Group numbering and the modern periodic table

The original group numbers of the periodic table ran from group 1 alkali metals to group 0 noble gases. To account for the d block elements and their 'vertical' similarities, in the modern periodic table, groups 3 to group 0 are numbered 13 to 18. So, the p block elements are referred to as groups 13 to group 18 at a higher academic level, though the group 3 to 0 notation is still used, but usually at a lower academic level. The 3d block elements (Sc to Zn) are now considered the head (top) elements of groups 3 to 12.

Website content © Dr Phil Brown 2000+. All copyrights reserved on revision notes, images, quizzes, worksheets etc. Copying of website material is NOT permitted. Doc Brown's Chemistry theoretical-physical chemistry revision notes for pre-university level students on d-block elements including the physical and chemical properties reactions equations and trends explained for the 3d-block of transition metals series