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Doc Brown's Chemistry  Advanced Level Inorganic Chemistry Periodic Table Revision Notes – Transition Metals

Part 10. Transition Metals 3d–block:  10.6 Chromium Chemistry

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.

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

(c) doc b GCSE/IGCSE Periodic Table Revision Notes * (c) doc b GCSE/IGCSE Transition Metals Revision Notes

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


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. 3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
Electrode potential M(s)/M2+(aq) na –1.63V –1.18V –0.90V –1.18V –0.44V –0.28V –0.26V +0.34V –0.76V
Electrode potential M(s)/M3+(aq) –2.03V –1.21V –0.85V –0.74V –0.28V –0.04V +0.40 na na na
Electrode potential M2+(aq)/M3+(aq) na –0.37V –0.26V –0.42V +1.52V +0.77V +1.87V na na na

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 [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

Advanced Inorganic Chemistry Page Index and Links

  • 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. This 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.

Advanced Inorganic Chemistry Page Index and Links


The Chemistry of CHROMIUM

CHROMIUM(III) oxidation state chemistry

  • Chromium forms the stable green (greyish dark green almost violet sometimes?) chromium(III) ion, [Cr(H2O)6]3+(aq).

  • 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) 

      • The hydroxide readily dissolves in acids to form salts,

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

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

          • or more simply Cr(OH)3(s) + 3H+(aq) rev Cr3+(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+ => [Cr(OH)(H2O)5]2+ => [Cr(OH)2(H2O)4]+ => [Cr(OH)3(H2O)3](s) precipitate

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

    • VIEW ppts. with OH, NH3 and CO32–, & 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 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) is formed

    • No Cr2(CO3)3 is formed because of the acid–base reaction above, due to the acidity of the chromium(III) ion. Note the similarly highly charged small ions Al3+ and Fe3+ behave in the same way.

    • * 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)   (from original hexa–aqua ion)

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

        • 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 all octahedral complexes with a co–ordination number of 6.

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

  • Advanced Inorganic Chemistry Page Index and LinksIonisation isomerism in chromium(III) chloride based on Cr3+, 3Cl and 6H2

    • [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? *the 0 to signify an overall electrically neutral complex can be omitted)

    • and this is not all, the 3rd one down with two chloride ligands can exist as cis (1) or trans (2) geometric isomers (Z/E isomers) illustrated below, and also serve as models for representing the other octahedral complexes which exhibit cis/trans or Z/E isomerism.

    • (c) doc b

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

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

    • You also get geometrical cis/trans isomers (Z/E) with tetraamminedichlorochromium(III) complexes.

    • (c) doc b

  • 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 optical isomerism.

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

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

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

      • This complex has mirror image forms i.e. enantiomers of optical isomers.

        • This optical isomerism can be illustrated thus

        • where L–L represents H2NCH2CH2NH2

        • The ligand bonds via the lone pairs of electrons on the nitrogen which are donated to form the metal–ligand dative covalent 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]

    • [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]

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

Advanced Inorganic Chemistry Page Index and Links


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.

    • 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.

    • 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 pH.

      • 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.

      • 2CrO42–(aq) + 2H+(aq) rev Cr2O72–(aq) + H2O(l)

      • 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.

  • 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), reduction 6 O(–1) in 3H2O2 ==> 6(–2) in 6 of the 8H2O, total of 6 'units' oxidation state change.

    • Both H2O2 and Cr(VI) compounds are oxidising agents but in alkaline solution H2O2 is the stronger oxidising agent.

      • EØ = +?V details to add???

    • When the resulting solution from above is acidified with dilute sulphuric 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/dilute sulphuric 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.

  • 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 sulphuric acid, using a redox indicator like barium diphenylamine sulphonate 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!

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

  • The dichromate(VI) ion 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 thiosulphate solution using starch indicator.

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

  • 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)

    • A few drops of silver chromate is used as an indicator when titrating chloride solutions with silver nitrate solution in neutral solution. The solubility product for the white ppt. of silver chloride

    • Ksp = [Ag+(aq)][Cl(aq)] = 2 x 10–10 mole2dm–6

    • is exceeded before the solubility product of silver chromate(VI)

    • Ksp = [Ag+(aq)]2[CrO42–(aq)] = 3 x 10–12 mole3dm–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 ppt

  • CHROMIUM(II) oxidation state chemistry:

    • 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 green chromium(III) ions by dissolved oxygen unless protected by an inert atmosphere.V3+/V2+ EØ = –0.26V, O2+H+/H2O EØ = +1.23V , see Redox Electrode Potential Chart


 

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: 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


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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

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