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

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

Extended data table for COPPER

• Uses 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 cooper 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, are described on the GCSE Reactivity Series of Metals Notes

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

• COPPER(II) CHEMISTRY

• 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

    • 

      • On the right the diagram shows the excited state of the copper(II) ion

• 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 in the blue pentahydrate crystals of CuSO₄.5H₂O

    • (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..

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

  • [Cu(H₂O)₆]²⁺_(aq) + 2OH^–_(aq) ==> [Cu(H₂O)₄(OH)₂]_(s) + 2H₂O_(l) 

    • or more simply: Cu²⁺_(aq) + 2OH^–_(aq) ==> Cu(OH)_2(s)

    • A  precipitation reaction involving ligand displacement.

• Excess sodium hydroxide has no significant effect, BUT with excess ammonia, a deep blue solution is formed of the ammine complex ion (ligand substitution is incomplete), the overall change can be expressed as:

  • [Cu(H₂O)₆]²⁺_(aq) + 4NH_3(aq) [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) + 4H₂O_(l)

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

  • or from the hydroxide precipitate

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

    • or more simply: Cu(OH)₂]_(s) + 4NH_3(aq)  [Cu(NH₃)₄]²⁺_(aq) + 2OH^–_(aq)

      • All showing the formation of the diaquatetramminecopper(II) ion/tetraamminecopper(II) ion.

  • Note: ligand exchange reaction, 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 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 concentration the dark violet hexaamminecopper(II) ion can be formed.

• 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 sodium hydrogencarbonate solution.

    • Cu²⁺_(aq) + 2HCO₃^–_(aq) ==> CuCO_3(s) + H₂O_(l) + CO_2(g)

• VIEW ppts. with OH^–, NH₃ and CO₃^2–, & complexes, if any, with excess reagent.

• If e.g. sodium chloride or hydrochloric acid is added to copper(II) sulphate solution the pale yellow–brown tetrachlorocuprate(II) complex ion is formed (seen as green due to the 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

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

    • the complex changes from a cationic complex ion to an anionic complex ion.

  • 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 principles:

    • 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 salts:

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

  • On mixing solutions of a copper(II) salt e.g. blue copper(II) sulphate 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 thiosulphate 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 ==> 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 thiosulphate (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

• COPPER(I) CHEMISTRY

• 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.

      • 

      • Bottom right 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:

  • If solid copper(I) oxide is dissolved in dil. sulphuric acid a pinky–brown precipitate of copper and a blue solution of copper(II) sulphate 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) sulphate 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) with the more positive redox potential represents the reduction half–cell reaction and (ii), reversed, with the less positive potential, will represent the oxidation half–cell reaction.

    • 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.

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

        • 2Cu⁺_(aq) ==> Cu²⁺_(aq) +Cu_(s)

  • See manganese(VI) chemistry for another example of disproportionation.

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

  • 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!

    • 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 left out in air will discolour further due to aerial oxidation.

  • Copper(I)/Cu+(aq) can be stabilised by making 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)

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

      • 2Cu_(s) + 2H⁺_(aq) + 4Cl^–_(aq) ==> 2[CuCl₂]^–_(aq) + H_2(g)

      • The Cu²⁺/Cu potential is +0.34V and the Cu+/Cu potential is +0.15V, so 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 which is <0.00V.

    • 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^–.

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

    • In principle the 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.

• Biochemistry of Copper

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

  • Details to do.

Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

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

Advanced Level Inorganic Chemistry of Copper – A level Revision notes to help revise for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Advanced Level Chemistry OCR GCE Advanced Level Chemistry Edexcel GCE Advanced Level Chemistry Salters AS A2 Chemistry CIE Chemistry, WJEC GCE AS A2 Chemistry, CCEA/CEA GCE AS A2 Chemistry revising courses for pre–university students (equal to US grade 11 and grade 12 and AP Honours/honors level courses)

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Scandium * Titanium * Vanadium * Chromium * Manganese * 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, E^ø * 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