Doc Brown's advanced A level inorganic chemistry exam revision notes

Inorganic chemistry Part 10. Transition Metals 3d–block

The chemistry of copper  - a 3d block transition element

10.11 Revision notes on Copper Chemistry including oxidation states & complex ions

Sub-index for this page on the chemistry of copper - a 3d block transition element

1. Introduction to the chemistry of copper

2. Data tables for the chemistry of copper

3. Uses of copper and its biological role

4. Electron configuration, oxidation states and electrode potentials

5. Copper(II) chemistry, complexes and a few aspects of copper(I) chemistry

6. More on copper(I) chemistry

7. The original extraction of copper from copper ores and purification via electrolysis

8. The vertical connection of copper with the other d-block elements of Group 11 (IUPAC designation)

9. Learning objectives for the chemistry of copper

1. Introduction to the chemistry of copper and its complexes - a 3d block transition element

Copper is the head-top element of Group 11 of the periodic table (modern IUPAC assignment).

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]

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

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

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

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

Extended data table for COPPER

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.

3. Uses of COPPER and the biochemistry 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.

4. The Chemistry of COPPER - electron configuration, oxidation states and electrode potentials

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

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

Electrode potential chart for copper the context of the 3d block transition metals

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.

PLEASE note: The electrode potentials (E^Ø) for copper chemistry in the text sometimes vary from the chart above - apologies, but data sources for copper chemistry can vary! Any discrepancy shouldn't significantly affect any electrode potential calculation outcomes e.g feasibility (email if concerned?).

5. COPPER(II) CHEMISTRY, complexes and a few aspects of copper(I) chemistry

• 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.

  • i.e. 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

  • 

  • The right of the diagram shows the excited state of the copper(II) ion in an octahedral complex

  • See also the absorption spectra and colours of copper compounds for a detailed discussion.

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

  • [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 CuSO_4(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 pale 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)

    • A  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) + H₂SO_4(aq) ===> CuSO_4(aq) + H₂O_(l)

      • CuO_(s) + 2HCl_(aq) ===> CuCl_2(aq) + H₂O_(l)

      • CuO_(s) + 2HNO_3(aq) ===> Cu(NO₃)_2(aq) + H₂O_(l)

• 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:

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

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

  • (i) [Cu(H₂O)₆]²⁺_(aq) +  2NH_3(aq) [Cu(H₂O)₄(OH)₂]_(s) + 2NH₄⁺_(aq)

  • (ii) [Cu(H₂O)₄(OH)₂]_(s) + 4NH_3(aq)  [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) +  2OH^–_(aq) +  2H₂O_(l)

    • 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) +  4NH_3(aq) [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) + 4H₂O_(l)

    • K_c = [[Cu(NH₃)₄(H₂O)₂]²⁺_(aq)] / [[Cu(H₂O)₆]²⁺_(aq)] [NH_3(aq)]⁴

      • = 1.2 x 10¹³ mol^-4dm¹²

    • Sometimes shown as

      • (iv) [Cu(H₂O)₆]²⁺_(aq) +  4NH_3(aq) [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) + 4NH_3(aq) [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.

    • If 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.

  • K_stab = [ [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) ] / [ [Cu(H₂O)₆]²⁺_(aq) ] [ NH_{3 (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).

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

  • Cu²⁺_(aq) + CO₃^2–_(aq) ===> CuCO_3(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) ===> CuCO_3(s) + H₂O_(l) + CO_2(g)

• VIEW more on ppts. with OH^–, NH₃ and CO₃^2–, 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).

  • [Cu(H₂O)₆]²⁺_(aq) + 4Cl^–_(aq) [CuCl₄]^2–_(aq) + 6H₂O_(l)

    • K_stab = [ [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) + I_2(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.

  • 2S₂O₃^2–_(aq)  +  I_2(aq)  ===>  S₄O₆^2–_(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 e.g.. 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

6. More on COPPER(I) CHEMISTRY

• 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).

  • i.e. 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) + H₂SO_4(aq) ===> Cu_(s) + CuSO_4(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.

    • Cu₂SO_4(s) + aq ===> Cu_(s) +  CuSO_4(aq)

    • Cu₂SO_4(s) + aq ===> Cu_(s) +  Cu²⁺_(aq) + SO₄^2–_(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   (E^Ø_Cu+/Cu = +0.52V)

    • (ii) Cu²⁺ + e^– Cu⁺   (E^Ø_Cu2+/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.

    • E^Ø_reaction = E^Ø_reduction – E^Ø_oxidation = (+0.52) – (+0.15) = +0.37V

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

      • i.e. 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) + I_2(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 e.g. 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 E^Ø 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−		Cu+(aq)	+0.16
      Cu2+(aq) + 2 e−		Cu(s)	+0.34
      Cu+(aq) +  e−		Cu(s)	+0.52
      I3−(aq) + 2 e−		3 I−(aq)	+0.53
      I2(s) + 2 e−		2 I−(aq)	+0.54
      Cu2+(aq) + I –(aq) + e–		CuI(s)	+0.86

    • E^Ø_reaction = E^Ø_reduction (most +ve half-cell) – E^Ø_oxidation (least +ve half cell), therefore ...

    • E^Ø_reaction 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 E^Ø/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) + H_2(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 (E^Ø_H+/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^–    [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)  +  CuCl_2(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 tetrahedral 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

8. The vertical connection of copper with the other d-block elements of Group 11 (IUPAC designation)