* Advanced Inorganic Chemistry Transition metals 10.11 Copper Chemistry Doc Brown's

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 Doc Brown's Chemistry  Periodic Table revision notes 10.11

Part 10. Transition Metals 3d-block:   10.11 Copper Chemistry

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses)  GCSE Periodic Table * GCSE notes Transition Metals

INORGANIC Part 10 3d block TRANSITION METALS sub-index: 10.1-10.2 Introduction 3d-block Transition Metals * 10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron * 10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt * 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

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

copper110.11. Chemistry of Copper Cu, Z=29, 1s22s22p63s23p63d104s1 

  • Cu data table 1 summary * extra copper data table 2 * Copper & electrode potential 3d-block

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

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

  • COPPER(II) CHEMISTRY

  • 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 ??? ion (ligand substitution is incomplete), the overall change can be expressed as:

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

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

    • or from the hydroxide precipitate

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

      • or more simply: Cu(OH)2](s) + 4NH3(aq) rev [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.

    • 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.

  • Sodium carbonate 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 sodium hydrogencarbonate solution.

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

  • VIEW ppts. with OH-, NH3 and CO32-, & 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) rev [CuCl4]2-(aq) + 6H2O(l)

      • Kstab =

    • 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) + 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 ==> 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.

  • COPPER(I) CHEMISTRY

  • 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) + 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) sulphate 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) 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.

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

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

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

    • 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) + H2(g)

      • The Cu2+/Cu potential is +0.34V and the Cu+/Cu potential is +0.15V, so hydrogen shouldn't be formed (EØH+/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)

  • Biochemistry of Copper

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

    • Details to do.

Advanced Inorganic Chemistry Page Index and Links

 

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 sulphide ores ...
    • These include chalcocite/chalcosine = copper(I) sulphide Cu2S and covellite = copper(II) sulphide 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 sulphide ores can be rapidly roasted in heated air enriched with oxygen to form impure copper and this extraction process is called 'flash smelting'.
      • Nasty sulphur dioxide gas is formed, this must be collected to avoid pollution and can be used to make sulphuric acid to help the economy of the process.
      • copper(I) sulphide + oxygen ==> copper + sulphur dioxide
        • Cu2S(s) + O2(g) ==> 2Cu(s) + SO2(g)
      • or copper(II) sulphide + oxygen ==> copper + sulphur dioxide
        • CuS(s) + O2(g) ==> Cu(s) + SO2(g)
  • It is also possible to dissolve an oxide or carbonate ore in dilute sulphuric 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) sulphate solution.
      • iron + copper(II) sulphate ==> iron(II) sulphate + copper
      • Fe(s) + CuSO4(aq) ==> FeSO4(aq) + Cu(s)

 

The Purification of Copper by Electrolysis

(c) doc b

copper1copper2
  • 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) sulphate 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) sulphate 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 for extraction of impure copper from an ore.

Raw materials for the electrolysis process:

  • Impure copper from a copper smelter.

  • Electrolyte of aqueous copper(II) sulphate.

  • 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)

  • Note: Reduction and Oxidation always go together, hence the use of the term redox change or reaction.
  • Electroplating is mentioned on the Industrial Chemistry and Electrochemistry pages.

 

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

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