transition metal chemistry of chromium complexes oxidation states +2 +3 +6 redox chemical reactions physical properties advanced inorganic chemistry of chromium

Revision notes: 3d block Transition Metals chemistry of chromium for advanced A level inorganic chemistry

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

Part 10. Transition Metals 3d–block

10.6 Chromium Chemistry

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The chemistry of chromium (principal oxidation states +3 and +6) is described with particular emphasis on chromium(III) complex ions with ligands such as water, ammonia and chloride ion, and the chromium(VI) oxyanions i.e. chromate(VI) and dichromate(VI) including the latter's redox reactions. A few chromium(II) complexes exist, but very easily oxidised y dissolved atmospheric oxygen.

The chemistry of the principal oxidation states of chromium, redox reactions of chromium, ligand substitution displacement reactions of chromium, balanced equations of chromium chemistry, formula of chromium complex ions, shapes colours of chromium complexes, formula of compounds

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

10.6. Chemistry of Chromium Cr, Z=24, 1s22s22p63s23p63d54s1 

Data comparison of chromium 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 +2 only
outer electron config.[Ar]... 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
Elect. pot. M(s)/M2+(aq) na –1.63V –1.18V –0.90V –1.18V –0.44V –0.28V –0.26V +0.34V –0.76V
Elect. pot. 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 chromium (EØ at 298K/25oC, 101kPa/1 atm.)

na = data not applicable to chromium

Extended data table for CHROMIUM

property of chromium/unit value for Cr
melting point Cr/oC 1857
boiling point Cr/oC 2672
density of Cr/gcm–3 7.19
1st Ionisation Energy Cr/kJmol–1 653
2nd IE/kJmol–1 1592
3rd IE/kJmol–1 2987
4th IE/kJmol–1 4740
5th IE/kJmol–1 6690
Cr atomic radius/pm 125
Cr2+ ionic radius/pm 84
Relative polarising power Cr2+ ion 2.4
Cr3+ ionic radius/pm 69
Relative polarising power Cr3+ ion 4.3
Cr4+ ionic radius/pm 56
Polarising power M4+ ion 7.1
oxidation states of Cr, less common/stable +2, +3, +6
simple electron configuration of Cr 2,8,13,1
outer electrons of Cr [beyond argon core] [Ar]3d54s1
Electrode potential Cr(s)/Cr2+(aq) –0.90V
Electrode potential Cr(s)/Cr3+(aq) –0.74V
Electrode potential Cr2+(aq)/Cr3+(aq) –0.42V
Electronegativity of Cr 1.66

There is an apparent anomaly in the electron configuration for chromium

Cr is [Ar]3d54s1 and not [Ar]3d44s2

because an inner half–filled 3d sub–shell seem to be a little lower in energy, and marginally more stable.

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 seem to be a little lower in energy, and marginally more stable.

Some general thoughts on chromium

  • Uses of CHROMIUM

    • Chromium is a hard bluish–white metal that is extremely resistant to chemical attack at room temperature e.g. very resistant to oxidation.

    • Chromium is used in the production of extremely hard steel alloys e.g. ball bearings.

    • Chromium metal is an important component in 'stainless steel'.

    • Chromium is used to electroplate other metals like steel because of its anti–corrosion properties ('chrome/chromium plating').

    • Chromium(III) oxide, Cr2O3 is used in stained glass and a catalyst in the chemical industry.

    • Chromium(IV) oxide is used in magnetic tapes for sound/video recording.

  • Biological role of chromium

    • Chromium is an essential trace element, but its role in the body is unknown?

  • Extraction of chromium

    • Chromium ore is processed and purified into chromium(III) oxide.

    • Chromium(III) oxide is reacted, very exothermically, in a Thermit style reaction, with aluminium (see reactions of aluminium) to free the chromium metal.

    • Cr2O3(s) + 2Al(s) ===> Al2O3(s) + 2Cr(s) 

    • The chromium(III) oxide is reduced to chromium by O loss, the aluminium is oxidised to aluminium oxide by O gain, and the aluminium is the reducing agent i.e. the O remover.

  • These are examples of metal displacement reactions e.g. the less reactive chromium or titanium are displaced by the more reactive sodium, magnesium or aluminium.

The Chemistry of CHROMIUM

Pd s block d blocks (3d block chromium) and f blocks of metallic elements p block elements
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Gp7/17 Gp0/18


2 3Li 4Be The modern Periodic Table of Elements

ZSymbol, z = atomic or proton number

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

5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc







 [Ar] 3d34s2



[Ar] 3d54s1



   [Ar]   3d54s2



[Ar] 3d64s2



[Ar] 3d74s2



[Ar] 3d84s2



[Ar] 3d104s1



[Ar] 3d104s2


31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba 57,58-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At 86Rn
7 87Fr 88Ra 89,90-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Nh 114Fl 115Mc 116Lv 117Ts 118Os
  *********** *********** ************ ************ ************** ********** ********** ********** ********** **********  

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
  (+2) (+2) (+2) (3d4) +2 +2 +2 +2 +2 +2
+3 +3 +3 +3  (3d3) (+3) +3 +3 (+3) (+3)  
  +4 +4   +4     (+4)    
      +6  (3d0) (+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 chromium in the context of the 3d block of elements

electrode potential chart for the oxidation states of chromium ions and chromium metal

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

The electrode potentials involving chromium 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 chromium ion.

Chromium(II) compounds are readily oxidised to chromium(III) compounds.

The hexaaquachromium(II) ion is a strong reducing agent.

CHROMIUM(III) oxidation state chemistry

  • diagram of the octahedral shape of the aqueous violet blue grey green hexaaquachromium(III) ion Cr3+(aq) [Cr(H2O)6]3+Chromium forms the stable hexaaquachromium(III) ion, [Cr(H2O)6]3+(aq)

    • Electron configuration of the Cr3+ ion is [Ar]3d3

    • The colour appears violet-blue-grey (some texts say red-violet?), but often looks green when produced in reactions, especially if chloride ions are present that can act as a ligand producing green coloured complexes where the chloride ion, like water, is acting as a monodentate ligand e.g..

      • [Cr(H2O)5Cl]2+(aq) is green and note how to work out the overall charge on these three chromium(III) complex ions, the substitution of a single chloride ion ligand reduces the overall complex ion charge from 3+ to 2+.

      • monochloropentaaquachrmium(III) complex ion The monochloropentaaquachormium(III) complex ion

      •  [Cr(H2O)4Cl2]+(aq), dichlorotetraaquachromium(III) ion is a dark green and again note the change in ionic charge, but no change in octahedral shape, coordination number 6 and oxidation state of +3 of either of these complex ions.

      • dichlorotetraaquachromium(III) complex ion E/Z isomers geometric example of cis/trans isomerism 

      • Z and E E/Z isomers (cis and trans isomers)
      • The 2nd Cl- ion ligand, further reduces the overall complex ion charge to a single + and there are two E/Z isomers.

    • The hexaaquachromium(III) ion has an octahedral shape and a co-ordination number of 6 from the six unidentate (monodentate) water ligands.

    • Small ligands like water, hydroxide ion or ammonia commonly form octahedral complexes with the chromium(III) ion and produce a variety of different colours.

    • You can see the red-purple Cr3+ ion colour in some solid gemstone compounds e.g. ruby in which chromium(III) ions and surrounded by an octahedral arrangement of oxygen atoms. A similar arrangement is found in blue-green emerald gemstones.

    • See the absorption spectra and colours of chromium compounds

  • Aqueous solutions of chromium(III) chloride are suitable for investigating the aqueous chemistry of the chromium(III) ion.

  • With aqueous ammonia (alkaline) or sodium hydroxide, chromium(III) ions form a green gelatinous precipitate of chromium(III) hydroxide.

    • Cr3+(aq) + 3OH(aq) ===> Cr(OH)3(s)   (but the structures can be quite complex)

    • or [Cr(H2O)6]3+(aq)  + 3OH(aq) ===> [Cr(OH)3(H2O)3](s) + 3H2O(l) 

      • Both chromium(III) complexes have an octahedral shape and a co-ordination number of 6 from 6 unidentate water or hydroxide ion ligands

      • The hydroxide readily dissolves in acids to form salts,

      • Cr(OH)3(s) + 3H+(aq) rev Cr3+(aq) + 3H2O(l) 

        • or more elaborately: [Cr(OH)3(H2O)3](s) + 3H3O+(aq) rev [Cr(H2O)6]3+(aq) + 3H2O(l)

        • thus showing amphoteric behaviour, since the hydroxide ppt. also dissolves in excess strong alkali to give a dark green solution and the hydroxide ppt. does not dissolve in the weak base aqueous sodium carbonate. However, it will dissolve in excess ammonia because a new green complex ion is formed. (more details on these reactions below)

    • The whole sequence of each theoretical step of chromium(III) hydroxide precipitation and its subsequent dissolving in strong base–alkali is shown the series of diagrams below.

    • All are, for simplicity, treated as octahedral complexes of 6 ligands – either water H2O or hydroxide ion OH

    • [Cr(H2O)6]3+(aq) ===> [Cr(OH)(H2O)5]2+(aq) ===> [Cr(OH)2(H2O)4]+(aq) ===> [Cr(OH)3(H2O)3](s) precipitate

    • dissolving ===> [Cr(OH)4(H2O)3](aq) ===> [Cr(OH)5(H2O)]2–(aq) ===> [Cr(OH)6]3–(aq)

    • [Cr(H2O)5OH]2+  is called the pentaaquamonohydroxochromium(III) or hydroxopentaaquachromium(III) ion.

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

The sequence of chromium(III) hydroxide precipitate formation and its subsequent dissolving in excess strong alkali. Each step is essentially one of proton removal from each complex (from 3+ to 3–).
1 2 3 4 From 1 to 7 happen as you add more alkali, increasing pH and the OH concentration, removing protons from the chromium(III) complex.
5 6 7 * From 7 back to1 represents what happens when you add acid, decreasing pH, increasing H+/H3O+ concentration and protonating the chromium(III) complex.
  • Chromium(III) ions with aqueous sodium carbonate form the green hydroxide precipitate (as above) and carbon dioxide is liberated because of the acidity of the hexaaquachromium(III) ion (see Appendix 1.):

    • *initially 2[Cr(H2O)6]3+(aq) + CO32–(aq) ===> 2[Cr(H2O)5(OH)]2+(aq) + H2O(l) + CO2(g)

    • this process of proton donation (deprotonation) continues until [Cr(OH)3(H2O)3](s) precipitate is formed

    • No Cr2(CO3)3 is formed because of the acid–base reaction above, i.e. due to the acidity of the chromium(III) ion.

    • * the acidity of a the hexa–aquachromium(III) ion can be expressed as ...

      • [Cr(H2O)6]3+(aq) + H2O(l) rev [Cr(H2O)5(OH)]2+(aq) + H3O+(aq)

  • With excess sodium hydroxide or ammonia, further complex ions are formed from chromium(III) ions by ligand displacement/replacement reactions:

    • [Cr(H2O)6]3+(aq) + 6OH(aq) ===> [Cr(OH)6]3–(aq) + 6H2O(l)  (from original hexa–aqua ion)

      • or [Cr(OH)3(H2O)3](s) + 3OH(aq) ===> [Cr(OH)6]3–(aq) + 3H2O(l) (from hydroxide ppt.)

        • or more simply: Cr(OH)3(s) + 3OH(aq) ===> [Cr(OH)6]3–(aq)

        • showing amphoteric behaviour, since the hydroxide ppt. also dissolves in acid.

    • [Cr(H2O)6]3+(aq) + 6NH3(aq) ===> [Cr(NH3)6]3+(aq) + 6H2O(l)

      • a dark green colour, some texts say purple colour?

      • (equation from the original hexa–aqua ion)

      • or [Cr(OH)3(H2O)3](s) + 6NH3(aq) ===> [Cr(NH3)6]3+(aq) + 3OH(aq) + 3H2O(l)

        • (from the hydroxide precipitate)

        • or more simply: Cr(OH)3(s) + 6NH3(aq) ===> [Cr(NH3)6]3+(aq) + 3OH(aq)

    • The uncharged ligand molecules ammonia NH3 and water H2O are similar in size and ligand exchange occurs without change in co–ordination number.

    • They are all octahedral complexes with a co–ordination number of 6 from 6 unidentate ligands.

  • Chromium(III) complexes are extremely numerous and varied, including many examples of isomerism. (see Appendix 2 and Appendix 3 for an introduction to complexes)

  • Ionisation isomerism in chromium(III) chloride based on Cr3+, 3Cl (and 6H2O)

    • A variety of octahedral complexes are theoretically possible and all do exist, including in some cases the crystalline isomers can be isolated.

    • [Cr(H2O)6]3+(Cl)3  (violet or grey–blue?)

    • [CrCl(H2O)5]2+(Cl)2.H2O  (pale green)

    • [CrCl2(H2O)4]+ Cl.2H2O  (dark green)

    • [CrCl3(H2O)3]0*.3H2O ?  (brown?, this I found reference to on a Russian website, doesn't seem to be in textbooks?

      • Note: * the superscript 0 is to signify an overall electrically neutral complex, it can be omitted, but clarifies the situation.

    • AND, this is not all on isomerism with this set of chromium(III) complexes.

      • The 3rd one down with two chloride ligands can exist as E/Z isomers (geometric isomerism).

      • (1) is the Z isomer (cis (1)) and (2) is the E isomer (trans), with respect to the chloride ion ligand, both illustrated below

      • (c) doc b

      • This example serves as a model for representing the other octahedral complexes which exhibit E/Z (cis/trans) isomerism.

    • With excess chloride ion you get the formation of the tetrahedral tetrachlorochromate(III) ion, colour?

      • [Cr(H2O)6]3+(aq) + 4Cl(aq) ===> [CrCl4](aq) + 6H2O(l)

      • Note the charge on the complex is (+3) + (4 x -) = -

    • You also get E/Z isomerism (cis/trans) with tetraamminedichlorochromium(III) complexes (dichlorotetraamminechromium(III) ion.

      • E/Z isomerism (cis/trans) diagrams of tetraamminedichlorochromium(III) complexes dichlorotetraamminechromium(III) ion (Z and E isomers, key as above)

  • A similar case of isomerism occurs with the chromium(III) complexes with ammonia and chloride ligands shown above.

    • All the complex ions above have a plane of symmetry and cannot exhibit R/S isomerism (optical isomerism).

    • Again, these are all octahedral complexes with a coordination number of 6.

  • [Cr(H2NCH2CH2NH2)3]3+, in this complex, H2NCH2CH2NH2, ethane–1,2–diamine (ethylenediamine), is often represented in shorthand by en,

    • and the complex simply written as [Cr(en)3]3+.

    • or more accurately as   [Cr(H2NCH2CH2NH2)3]3+

    • This is an example of a chromium(III) complex with a bidentate ligand.

    • There are three ligands but the co-ordination number is still 6 because there are 6 central metal ion-ligand bonds.

    • This complex has mirror image forms i.e. R/S isomers  - enantiomers (optical isomers).

      • This optical isomerism can be illustrated thus

      • where L–L represents H2NCH2CH2NH2  and  M = Cr3+

      • The ligand bonds via the lone pairs of electrons on the nitrogen which are donated to form the metal–ligand dative covalent bonds (co-ordinate bonds).

  • Both the hexa–aqua ions of chromium(II) and chromium(III) readily complex with EDTA

    • [Cr(H2O)6]2+(aq) + EDTA4–(aq) ===> [Cr(EDTA)]2–(aq) + 6H2O(l)

      • Kstab = [[Cr(EDTA)3]2–(aq)] / [[Cr(H2O)6]2+(aq)] [EDTA4–(aq)]

      • Kstab = 1.0 x 1013 mol–1 dm3 [lg(Kstab) = 13.0]

      • Remember [H2O] is not included in the equilibrium expression.

    • [Cr(H2O)6]3+(aq) + EDTA4–(aq) ===> [Cr(EDTA)](aq) + 6H2O(l)

      • Kstab = [[Cr(EDTA)3](aq)] / [[Cr(H2O)6]3+(aq)] [EDTA4–(aq)]

      • Kstab = 1.0 x 1024 mol–1 dm3 [lg(Kstab) = 24.0]

    • The co-ordination number is still 6 but only one EDTA ligand molecule bonds to each cobalt ion.

    • From the Kstab values, you can see that the more highly charged Cr3+(aq) ion complexes more strongly than the Cr2+(aq) ion.

  • See also oxidation of a chromium(III) ion to a chromium(VI) ion

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

    • It shows the weak oxidising or weak reducing power of chromium(III), Cr3+(aq), the strong reducing power of chromium(II), Cr2+(aq) and the strong oxidising power of chromium(VI), Cr2O72–

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

CHROMIUM(VI) oxidation state chemistry

  • The 'simple' hexaaquachromium(VI) cation, [Cr(H2O)6]6+, cannot exist in aqueous media.

    • In fact, I doubt if the 'simple' Cr6+ ion can exist in any ionic compound, the high charge would 'theoretically draw 'over' the electron cloud of the negative anion to form a covalent bond and it would be a very acidic complex.

    • Note that chromium(VI) oxide, CrO3 and chromium(VI) fluoride, CrF6, (both compounds have Cr in +6 oxidation state), are covalent compounds, despite the relatively large electronegativity difference between the metal and non–metal.

      • BUT, being very electronegative, they can stabilise chromium compounds-ions in the higher +6 oxidation state.

    • If the oxidation state of the central metal ion is over +3, it appears that deprotonation via proton transfer to water is so facilitated that in most cases (there may be exceptions?) all protons are 'theoretically' lost to give the oxyanion.

      • ie the theoretical [Cr(H2O)6]6+ ends up in reality as Cr2O72– or CrO42– depending on the pH of the aqueous solution.

      • The reason for this situation is that the high charge density of the 'theoretical' central metal ion, gives it a high polarising power.

        • You can theoretically conceive the situation of imagining the central metal ion pulling on the electrons of the M–OH2 and M–OH ligand bonds in two stages to leave a M=O bond in the oxyanion

          • ie considering one co–ordinated water molecule (and ignoring the charge on intermediate complexes) ...

            • M–OH2  ==proton loss==>  M–OH  ==proton loss==>  M=O

    • There is, theoretically, always an equilibrium between the chromate(VI) ion and the dichromate(VI) ion which can be expressed in two ways.

      • (i) 2CrO42–(aq) +  2H+(aq) Cr2O72–(aq) + H2O(l)      (adding acid to a chromate(VI) salt solution)

      • (ii) Cr2O72–(aq) +  2OH-(aq) 2CrO42–(aq)H2O(l)  (adding alkali to a dichromate(VI) salt solution)

      • Therefore from Le Chatelier's principle, high pH (alkaline) favours the formation of the chromate(VI) ion and low pH (acid) favours dichromate(VI) ion formation.

  • Oxidation of a chromium(III) ion to a chromium(VI) ion:

  • When hydrogen peroxide is added to an alkaline chromium(III) solution, oxidation occurs to give the yellow chromate(VI) ion CrO42–

    • 2Cr3+(aq) + 3H2O2(aq) + 10OH(aq) ===> 2CrO42–(aq) + 8H2O(l)

    • Redox changes: oxidation 2Cr(+3) ==> 2Cr(+6), and for the corresponding ....

      • reduction 6 O(–1) in 3H2O2 ==> 6(–2) in 6 of the 8H2O

      • a total of 6 'units' oxidation state change, which I sometimes unofficially call 6 'electrons worth' of change!

    • Both Cr(VI) compounds and hydrogen peroxide and are oxidising agents e.g.

      • EØ = +1.33V for the half-cell reaction: Cr2O72–(aq) + 14H+(aq) + 6e rev 2Cr3+(aq) + 7H2O(l)

      • Hydrogen peroxide is a stronger oxidising agent, and for the half–cell reaction:

      • EØ = +1.77 V  for H2O2(aq) + 2H+(aq) + 2e rev 2H2O(l)

      • BUT, both of the above are for acid–neutral conditions, so different half-cell potential data must be used for this Cr(III) to Cr(VI) conversion.

        • What we are looking at here is a case of a pH change producing different half-cell potentials and different species involved in a redox reaction.

      • For strongly alkaline conditions for the conversion of chromium(III) ion to the chromate(VI) ion the following half–cell potential data should be used involving the perhydroxyl ion (HO2 or HOO):

        • (a) EØ = +0.88 V  for the half–cell reaction: HO2(aq) + H2O(aq) + 2e rev 3OH(aq)

        • So alkaline hydrogen peroxide solution is still quite a strong oxidising agent.

        • Note the oxidant species is considered to be the perhydroxyl ion.

      • However, unlike the orange dichromate(VI) ion (Cr2O72–), the yellow chromate(VI) ion is a very weak oxidising agent in alkaline solution (in acid solution it reverts to the dichromate(VI) ion).

        • EØ = –0.11 V for the half–cell reaction (b): CrO42–(aq) + 4H2O(l) + 3e rev Cr(OH)3(s) + 5OH(aq)

      • From the standard electrode potentials (+0.88 V > –0.12 V) you can clearly see that the hydrogen peroxide can oxidise the chromium(III) ion/hydroxide to the chromate (VI) ion in alkaline conditions.

        • EØreaction = EØred EØox = EØHOO–/OH–    EØCrO42–/Cr(OH)3 = +0.88 0.12 = + 1.00 V, a very feasible reaction!

        • Note that the chromium species used in the EØ argument involves chromium(III) hydroxide.

      • It is quite legitimate to do so, because when you add sodium hydroxide to a chromium(III) salt solution you get a green precipitate of chromium(III) hydroxide.

        • simple equation: Cr3+(aq) + 3OH(aq) ===> Cr(OH)3(s)

      • When the hydrogen peroxide is added, this green precipitate is oxidised and dissolves to give a yellow solution of sodium chromate(III).

      • So the reaction can better written as (though derivation not shown, see equation (d) at the end):

      • (c) 2Cr(OH)3(s)  +  3H2O2(aq)  +  4OH(aq)  ===>  2CrO42–(aq)  +  8H2O(l)

      • The oxidation state change numbers and balancing are as above.

      • Some of these redox equations are quite tricky to work out – do a triple check ...

        • .... especially when combining two half–cell equations, which I've illustrated for this reaction ...

        • (i) balance the number of species in the equation with the oxidation state changes

          • 2Cr(+3) ==> 2Cr(+6) and 6 O(–1) in 3H2O2 ==> 6 O(–2) in 6 of the oxygen's of the H2O's, 6e change

          • The total oxidation state increases must equal the total decrease in oxidation states.

        • (ii) check the ionic charge balance

          • on the left 4– is balanced by 2 x 2– = 4–, if it involves both + and – ions, just obey the usual mathematical rules

        • (iii) double check the atom count

          • 2Cr + 16O + 16H on both sides of the equation finally!

        • Do all three and you shouldn't go wrong!

      • However on reflection, you do not get equation (c) by combing half–cell equations (a) and (b) even though it is a legitimate equation,

        • BUT, using the perhydroxyl half–cell equation I've worked it through in one of my more 'nerdy' moments on my website to obtain equation (d), set out below.

        • I've kindly left out the state symbols until the end for logic clarity!

        • Check out the 'triple check' at each stage in deriving equation (d) for practice, that is well worth doing!

        • I think equation (d) is the most accurate depiction of the redox reaction that actually happens.

      • You may have to know how to carry out the reaction for your examination, but not all the details, BUT, you should understand all the principles used in this explanation and appreciate that changing the pH of an oxidant can change both its oxidising power and species involved.

3 x (a) for 6e change

3HO2 + 3H2O + 6e



2 x (b) reversed for 6e change

2Cr(OH)3 + 10OH


2CrO42– + 8H2O + 6e

= initial total, 6e cancel out

3HO2 + 3H2O + 2Cr(OH)3 + 10OH

===> 2CrO42– + 8H2O + 9OH

3H2O and 9OH cancel out to give

3HO2 + 2Cr(OH)3 + OH

===> 2CrO42– + 5H2O

leaving the final equation (d)!

2Cr(OH)3(s)  +  3HO2(aq)  +  OH(aq) ===> 2CrO42–(aq)  +  5H2O(l)

  • As mentioned already, when the resulting yellow solution from above is acidified with dilute sulfuric acid, the orange dichromate(VI) ion Cr2O72–  is formed.

  • The equilibrium is pH dependent. From 'Le Chatelier's Principle':

    • in more acidic solution, more H+, decrease pH ==> more orange (net change L to R) or in

    • more alkaline, less H+ (removed by OH), increase pH <= more yellow (net change R to L).

    • 2CrO42–(aq) + 2H+(aq) rev Cr2O72–(aq) +  H2O(l) (no change in ox. state)

  • The dichromate(VI) ion is reduced in two stages by a zinc and dilute hydrochloric/sulfuric acid mixture.

    • Cr(VI, +6) ==> Cr(III, +3): Cr2O72–(aq) + 14H+(aq) + 6e rev 2Cr3+(aq) + 7H2O(l)

      • orange (+6) ==> green (+3),  EØ = +1.33V 

    • Cr(III, +3) ==> Cr(II, +2): Cr3+(aq) + e rev Cr2+(aq)

      • green (+3) ==> blue (+2), EØ = –0.41V, so Cr(II) is readily oxidised by dissolved oxygen and can only be retained in an inert atmosphere.

    • Note the  EØZn(s)/Zn2+(aq) is –0.76V, so the reducing power of zinc is sufficient to effect either of the two chromium oxidation state reduction changes.

      • The full redox equations for the reactions which happen on the surface of the zinc are:

      • Cr2O72–(aq) + 3Zn(s) + 14H+(aq) rev 2Cr3+(aq) + 3Zn2+(aq) + 7H2O(l)

      • 2Cr3+(aq) +  Zn(s) rev 2Cr2+(aq) + Zn2+(aq)

      • You will see hydrogen formed as a by–product of the zinc–acid reaction but the reductions take place on the surface of the zinc.

      • The reductions occur by electron transfer on the surface of the zinc.

  • Potassium dichromate(VI), K2Cr2O7,  can be crystallised to high purity standard without water of crystallisation, and is a valuable 'standard' redox volumetric reagent.

    • e.g. It can used to titrate iron(II) ions in solution acidified with dilute sulfuric acid, using a redox indicator like barium diphenylamine sulfonate (sulfonate) which is less readily oxidised than iron(II) ions, but once all the iron(II) ions are oxidised the indicator is oxidised to a blue colour.

    • The iron(III) ions formed affect the indicator to give an inaccurate end point so phosphoric(V) acid is also added at the start to complex the Fe3+ ions as they form.

    • Cr2O72–(aq) + 14H+(aq) + 6Fe2+(aq) ===> 2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l)

    • Theoretically, there are actually two simultaneous colour changes, both masked by the redox indicator change.

      • The orange dichromate(VI) ion changes on reduction to the green chromium(III) ion,

      • and the pale green iron(II) ion changes on oxidation to the orange iron(III) ion,

      • so without the indicator I'm not sure exactly how the colour change you would really observe would pan out, but the problem is got round by using a special redox indicator!

      • One of the most common indicator is sodium diphenylaminesulfonate, which turns from colourless to purple with first tiny excess of dichromate at the end-point.

      • You must have excess dil. sulfuric acid in the titration flask mixture - you can use dil. hydrochloric acid, because the dichromate(VI) ion is not a powerful enough oxidising agent to oxidise chloride ion to chlorine.

    • See also fully worked examples of redox volumetric titration calculation questions, and ...

    • Constructing full inorganic redox equations from half–equations and redox titrations

    • and Balancing redox equations involving transition metal ions

  • The dichromate(VI) ion in acid solution is a strong oxidising agent – examples of oxidising action:

  • See above for oxidation of iron(II) ions.

  • It oxidises iodide ions to iodine.

  • Cr2O72–(aq) + 14H+(aq) + 6I(aq) ===> 2Cr3+(aq) + 3I2(aq) + 7H2O(l)

    • The released iodine can be titrated with standard sodium thiosulfate solution using starch indicator.

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

    • This reaction between the released iodine and sodium thiosulfate can be used to estimate oxidising agents like dichromate(VI) 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.

  • Soluble chromate(VI) salts give yellow solutions, but lead(II) ions give a yellow ppt. of lead(II) chromate(VI) and silver ions a dark red ppt. of silver chromate(VI).

  • Pb2+(aq) + CrO42–(aq) ===> PbCrO4(s)   and 

    • 2Ag+(aq) +  CrO42–(aq) ===> Ag2CrO4(s)

    • Note that a few drops of potassium chromate(VI) is used as an indicator when titrating chloride solutions with silver nitrate solution in neutral solution.

      • The solubility product for the white precipitate. of silver chloride is

      • Ksp = [Ag+(aq)][Cl(aq)] = 2 x 10–10 mol2dm–6

    • and is exceeded before the solubility product of silver chromate(VI) because of the relatively high concentration of chloride ions prior to reaching the end-point.

      • Ksp = [Ag+(aq)]2[CrO42–(aq)] = 3 x 10–12 mol3dm–9

      • until all the chloride is precipitated.

      • The next drop of silver nitrate causes the precipitation of brownish–red silver chromate, so the end point is the formation of the dark red precipitate.

CHROMIUM(II) oxidation state chemistry:

  • diagram of the octahedral shape of the aqueous blue hexaaquachromium(II) ion Cr2+(aq) [Cr(H2O)6]2+The blue hexaaquachromium(II) ion, [Cr(H2O)6]2+(aq), can be formed by reducing chromium(III) salt solutions with zinc and hydrochloric acid but it is rapidly oxidised back to violet-green chromium(III) ions by dissolved oxygen unless protected by an inert atmosphere.

  • See Redox Electrode Potential Chart and given the half-cell potentials ...

  • (i) O2(g)  +  4H+(aq)  +  4e-  ===>  2H2O(l)   (Eø = +1.23V, will act as the oxidising agent, and reduced)

  • (ii) Cr3+(aq) +  e-  ===>  Cr2+(aq)   (Eø = -0.42V, will act as the reducing agent, and oxidised)

  • Eøreaction = Eøreduction  -  Eøoxidation

  • Eøreaction = (+1.23)  -  (-0.42)  =  +1.65V  (very positive, so very feasible!)

  • for the reaction:

    • O2(g)  +  4H+(aq)  +  4Cr2+(aq)  ===>  4Cr3+(aq)  +  2H2O(l

    • Note the balancing from the 4 electron change equation.

    • One of (i) balances four of (ii) reversed.

    • Redox equation triple check!

      • Check on oxidation number changes,  the increase must numerically balance the decrease in oxidation states,

      • Check the total ion charge is same on both sides of equation,

      • Finally, do the usual check on the atom sums of each element are the same on each side of the equation.

keywords redox reactions ligand substitution displacement balanced equations formula complex ions complexes ligand exchange reactions redox reactions ligands colours oxidation states: chromium ions Cr(0) Cr(+2) Cr2+ Cr(II) Cr3+ Cr(+3) Cr(III) Cr(+6) Cr(VI) [Cr(H2O)6]3+ CrO42– Cr3+ + 3OH– ==> Cr(OH)3 [Cr(H2O)6]3+ + 3OH– ==> [Cr (OH)3(H2O)3] + 3H2O Cr(OH)3 + 3H+ ==> Cr3+ + 3H2O [Cr(OH)3(H2O)3] + 3H3O+ [Cr(H2O)6]3+ + 3H2O Cr(OH)3 + 3H+ ==> Cr3+ + 3H2O [Cr(H2O)6]3+ => [Cr(OH)(H2O)5]2+ => [Cr(OH)2(H2O)4]+ => [Cr(OH)3(H2O)3] precipitate dissolving => [Cr(OH)4(H2O)3]– => [Cr(OH)5(H2O)]2– => [Cr(OH)6]3– 2[Cr(H2O)6]3+ + CO32– ==> 2[Cr(H2O)5(OH)]2+ + H2O + CO2   [Cr(OH)3(H2O)3] [Cr(H2O)6]3+ + H2O [Cr(H2O)5(OH)]2+ + H3O+ [Cr(H2O)6]3+ + 6OH– ==> [Cr(OH)6]3– + 6H2O (from original hexa–aqua ion) or [Cr(OH)3(H2O)3] + 3OH– ==> [Cr(OH)6]3– + 3H2O (from hydroxide ppt.) or more simply Cr(OH)3 + 3OH– ==> [Cr (OH)6]3– [Cr(H2O)6]3+ + 6NH3 ==> [Cr(NH3)6]3+ + 6H2O (from original hexa–aqua ion) or [Cr(OH)3(H2O)3] + 6NH3 ==> [Cr (NH3)6]3+ + 3OH– + 3H2O (from hydroxide ppt.) or more simply Cr(OH)3 + 6NH3 ==> [Cr(NH3)6]3+ + 3OH– [Cr(H2O)6]3+(Cl–)3  (violet or grey–blue?) [CrCl(H2O)5]2+(Cl–)2.H2O  (pale green) [CrCl2(H2O)4]+ Cl–.2H2O  (dark green) [CrCl3(H2O)3]0.3H2O [Cr(H2O)6]3+ + 4Cl– ==> [CrCl4]– + 6H2O [Cr(H2NCH2CH2NH2)3]3+, H2NCH2CH2NH2 [Cr(en)3]3+ [Cr(H2O)6]2+ + EDTA4– ===> [Cr(EDTA)]2– + 6H2O Kstab = [[Cr(EDTA)3]2–] / [[Cr(H2O)6]2+] [EDTA4–] [Cr(H2O)6]3+ + EDTA4– ===> [Cr(EDTA)]– + 6H2O Kstab = [[Cr(EDTA)3]–] / [[Cr(H2O)6]3+] [EDTA4–] CrO42– 2Cr3+ + 3H2O2 + 10OH– ==> 2CrO42– + 8H2O 2CrO42– + 2H+ Cr2O72– + H2O Cr(VI, +6) ==> Cr(III, +3): Cr2O72– + 14H+ + 6e– 2Cr3+ + 7H2O Cr(III, +3) ==> Cr(II, +2): Cr3+ + e– Cr2+ Cr2O72– + 3Zn + 14H+ 2Cr3+ + 3Zn2+ + 7H2O 2Cr3+ + Zn 2Cr2+ + Zn2+ Cr2O72– + 14H+ + 6Fe2+ ==> 2Cr3+ + 6Fe3+ + 7H2O Cr2O72– + 14H+ + 6I– ==> 2Cr3+ + 3I2 + 7H2O Pb2+ + CrO42– ==> PbCrO4 and 2Ag+ + CrO42– ==> Ag2CrO4 Ksp = [Ag+][Cl–]  oxidation states of chromium, redox reactions of chromium, ligand substitution displacement reactions of chromium, balanced equations of chromium chemistry, formula of chromium complex ions, shapes colours of chromium complexes  Na2CO3 NaOH NH3 2CrO42– + 2H+ <=> Cr2O72– + H2O transition metal chemistry of chromium for AQA AS chemistry, transition metal chemistry of chromium for Edexcel A level AS chemistry, transition metal chemistry of chromium for A level OCR AS chemistry A, transition metal chemistry of chromium for OCR Salters AS chemistry B, transition metal chemistry of chromium for AQA A level chemistry, transition metal chemistry of chromium for A level Edexcel A level chemistry, transition metal chemistry of chromium for OCR A level chemistry A, transition metal chemistry of chromium for A level OCR Salters A level chemistry B transition metal chemistry of chromium for US Honours grade 11 grade 12 transition metal chemistry of chromium for pre–university chemistry courses pre–university A level revision notes for transition metal chemistry of chromium  A level guide notes on transition metal chemistry of chromium for schools colleges academies science course tutors images pictures diagrams for transition metal chemistry of chromium A level chemistry revision notes on transition metal chemistry of chromium for revising module topics notes to help on understanding of transition metal chemistry of chromium 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 chromium Edexcel A level chemistry notes on transition metal chemistry of chromium for OCR A level chemistry notes WJEC A level chemistry notes on transition metal chemistry of chromium CCEA/CEA A level chemistry notes on transition metal chemistry of chromium for university entrance examinations physical and chemical properties of the 3d block transition metal chromium, oxidation and reduction reactions of chromium ions, outer electronic configurations of chromium, principal oxidation states of chromium, shapes of chromium's complexes, octahedral complexes of chromium, tetrahedral complexes of chromium, square planar complexes of chromium, stability data for chromium's complexes, aqueous chemistry of chromium ions, redox reactions of chromium ions, physical properties of chromium, melting point of chromium, boiling point of chromium, electronegativity of chromium, density of chromium, atomic radius of chromium, ion radius of chromium, ionic radii of chromium's ions, common oxidation states of chromium, standard electrode potential data for chromium, ionisation energies of chromium, polarising power of chromium ions, industrial applications of chromium compounds, chemical properties of chromium compounds, why are chromium complexes coloured?, isomerism in the complexes of chromium, formulae of chromium compounds, tests for chromium ions how is chromium extracted? colour and chemistry of the chromium(III) ion, structure and chemistry of the octahedral complexes of chromium(III), effect of ammonia or sodium hydroxide on the hexaaquachromium(III) ion, EDTA complex with chromium(III) ion, standard electrode potential of Cr3+, oxidation of chromium(III) to chromium(VI) with alkaline hydrogen peroxide, chemistry of dichromate(VI) ion Cr2O7 2-, chemistry of the chromate(VI) ion CrO4 2-, formation of the chromium(II) ion Cr2+, isomerism in the octahedral complexes of chromium(III)


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 (= group 8). To account for the d block elements and their 'vertical' similarities, in the modern periodic table, groups 3 to group 0/8 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 0/8 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.

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