H2O, H2O2
H2O
Water can function both as an acid or as a base. Such compounds known as ampholytes, can give themselves a proton, in a process called autoprotolysis. The equilibrium constant of waters autoprotolysis is called the ionic product of water (pKw = 13,997). The formulas for the solubility product of metal oxides have the same structure as similar metal hydroxides.
H2O2
Hydrogen peroxide is a stronger acid than water (pKa1 = 11,65). The second acidity constant has not been measured accurately, but it is extremely small (16 < pKa < 18) and so the peroxide ion (O22−) in practice is not present in solutions.
Ag
Silver(I) ion appears in solutions as a colourless tetraaquasilver(I) ion ([Ag(H2O)4]+).
Silver(I) nitrate forms a mildly acidic solution with water (0,1 M solution pH = 6,48).
There is no known solid silver(I) hydroxide, instead the hydroxide ion precipitates brown silver(I)oxide (pKs = 7,71, precipitation window = 7,29-12,3). The oxide does not dissolve in excess hydroxide into hydroxide complexes (lgβ = 2,0 and 3,99).
Pb
In an aqueous solution lead(II) ions appear as colourless octaaqualead(II) ion ([Pb(H2O)8]2+). Lead(II) nitrates aqueous solution is acidic (0,1 M solution pH = 4,09).
The hydroxide ion precipitates white lead(II) oxide-hydroxide (pKs = 14,9, precipitation window = 7,05-9,55), which turns spontaniously into lead(II)oxide (15,1(yellow) ja 15,3(red)). The precipitate dissolves in excess reagents into hydroxide complexes (lgβ = 6,6, 11,4, 14,8, 17,1 and 7,6(lgβ21)).
Hg
In an aqueous solution mercury(I) ions apper as diaquamercury(I) ion ([Hg2(H2O)2]2+) and mercury(II) ions as hexaaquamercury(II) ion ([Hg(H2O)6]2+). The aqueous solution of Mercury(I) nitrates is strongly acidic (0,1 M Hg-solution pH = 2,56). Dissolving mercury(II) nitrate into water always results in the precipitation of mercury(II) oxide (dissolving 0,1 mol mercury(II) nitrate into a litre of water creates 0,0065 mol of mercury(II) oxide). However, mercury(II) nitrate dissolves completely into a 0,1 M solution in nitric acid as mild as 0,015 M. Mercury(I) ions most notable hydroxide complex is [(Hg2)2OH]3+, which has a stability constant of lgβ = 11,52. The values regarding dissolution are given without consideration of the hyrdoxide nitrates Hg2(OH)(NO3) and Hg(OH)(NO3), which can also be created, but whose stability is not known. Mercury(II) ions hydroxide complexes stability constants are: lgβ = 10,6, 21,83 ja 20,9.
The hydroxide ion precipitates mercury(I) ions as black mercury(I) oxide (pKs = 23,74, precipitation window = 3,13-8,13), which does not dissolve in reagent excess, but will turn gray when heated, as mercury(I) ions disproportionate into yellow mercury(II) oxide and free mercury. The hydroxide ion precipitates mercury(II) ions as mercury(II) oxide (pKs = 25,44, precipitation window = 1,78-4,28).
Bi
Bismuth is stable in an aqueous solutiuon only at oxidation state III. The normal salts of strong acids (except HI) with bismuth(III) are technically water soluble. When dissolving them into water however, practically quantitatively equal amounts of dihydroxide salts (Bi(OH)2NO3 (pKs = 30,55), Bi(OH)2Cl (pKs = 35,8),Bi(OH)2Br (pKs = 35,45) ja Bi2(OH)4SO4) are precipitated. At the same time the solution becomes extremely acidic as the bismuth(III) salt gives away two equivalents of the following strong acid:
|
|
Bi(NO3)3↓ + 2 H2O → Bi(OH)2NO3↓ + 2 H+ + 2 NO3– |
|
Therefore in order to dissolve bismuth(III) said acid needs to be added, in which case the formation of complexes affects the dissolution. E.g. 0,1 mol Bi(NO3)3·5H2O dissolves into a litre of 0,4 M nitric acid as [Bi(NO3)2]+ ions become the dominant specimen. As sodium hydroxide is added Bi(OH)2NO3 changes at pH 5,45 into Bi(OH)3, which dissolves poorly into hydroxide complexes in reagent excess (lgβ = 12,9, 23,5, 33,0 and 34,8, dissolving into a 0,001 M solution requires that [OH–] = 4 M).
Cu
Copper(II) ion appears in aqueous solutions as blue hexaaquacopper(II) ions. Copper(II) nitrates aqueous solution is acidic (0,1 M solution pH = 3,75).
The hydroxide ion precipitates copper(II) ions as blue copper(II) hydroxide (pKs = 18,7, precipitation window = 5,15-7,65), which dissolves into a strongly basic solution (dissolving into a 0,1 M solution requires that [OH−] = 4,46) as hydroxide complexes (6,5 (lgβ1 ) and 8,4 (lgβ21)). Copper(II) hydroxide turns spontaneously into dark brown copper(II) oxide (pKs = 19,5).
Cd
Cadmium appears in aqueous solutions as colourless hexaaquacadmium(II) ions ([Cd(H2O)6]2+). Cadmium nitrates aqueous solution is acidc (0,1 M solution pH = 5,46).
The hydroxide ion precipitates white cadmium hydroxide (pKs = 14,10(gamma), 14,35(beta), precipitation window = 7,45-9,95), which dissolves faintly in reagent excess (dissolving into a 0,001 M solution requires that [OH]– = 5,53) into hyrdoxide complexes (lgβ = 3,7, 7,7, 10,3, 8,7 and 4,6 (lgβ21)).
Sb
Antimony appears most commonly as antimony(II) oxide (Sb2O3) which dissolves only in strong acids as [Sb(OH)2]+ and in bases as [Sb(OH)4]– (pKs(Sb2O3)= 50,2 (Ks = [Sb3+][OH–]3) and hydroxide complexes lgβ = -, 33,1, 45,7 and 47,9, at 10 M H+ [Sb(OH)2+] = 0,008 M and at 10 M OH– [Sb(OH)4]– = 0,05 M). Antimony(III) compounds are better dissolved only with complex formation. For the oxidizing effects of atmospheric oxygen on cadmium(III) ions in basic conditions, see the redox section.
Co
In its aqueous solutions cobolt(II) ions appear as pink hexaaquacobolt(II) ions ([Co(H2O)6]2+). Cobolt(II) nitrates aqueous solution is acidic (0,1 M solution pH = 5,35).
The hydroxide ion precipitates cobolt(II) ions as bluish or reddish cobolt(II) hydroxide (pKs = 14,9(amorphous, blue) and 15,7(crystalline, red), precipitation window = 7,05-9,55). Cobolt(II) hydroxide dissolves faintly in reagent excess (dissolving into a 0,001 M solution requires that [OH–] = 4,09) into hydroxide complexes (lgβ = 4,3, 9,2, 10,5, 10,2).
Ni
In its aqueous solutions nickel(II) appears as green hexaaquanickel(II) ions ([Ni(H2O)6]2+). Nickel(II) nitrates aqueous solution is acidic (0,1 M solution pH = 5,45).
The hydroxide ion precipitates green nickel(II) hydroxide (pKs = 15,1(amorphous), 17,2(crystalline), precipitation window = 6,95-9,45). The precipitate dissolves slightly in reagent excess (dissolving into a 0,01 M solution requires that [OH–] = 3,08) into hydroxide complexes (lgβ = 4,1, 9,0, 12, 12).
Mn
Manganese(II) ion appears in an aqueous solution as hexaaquamanganese(II) ions ([Mn(H2O)6]2+), which is faintly pink in colour. Manganese(II) nitrates aqueous solution is acidic (0,1 M solution pH = 5,78).
The hydroxide ion precipitates white manganese(II) hydroxide (pKs =12,8, precipitation window = 8,1-10,6). The precipitate is practically insoluble in reagent excess (dissolving into a 0,001 M solution requires that [OH–] = 11,1) into hydroxide complexes (lgβ = 3,4, 5,8, 7,2, 7,7 and 3,4 (lgβ21)). Ammonium partly, or completely prevents the precipitation of hydroxides.
Fe
Iron(II) ion appears in an aqueous solution as very faintly light green hexaaquairon(II) ions ([Fe(H2O)6]2+). Iron(III) on the other hand foms colourless hexaaquairon(III) ions ([Fe(H2O)6]3+), that in a pure iron(III) nitrate solution hydrolyzes almost entirely into orange iron(III) oxide-hydroxide. An aron(II) nitrate solution is acidic (0,1 M solution pH = 1,52). Iron(III) nitrates solution is the most acidic after the chrome(III) nitrate solution (0,1 M solution pH = 1,52). Because iron(III) hydroxides solubility reduces with time, to the originally clear iron(III) nitrate solution, iron(III) hydroxide is precipitated over time. In order to prevent this completely a 0,1 M iron(III) nitrate solution must contain 0,8 mol/l of nitrcic acid.
The hydroxide ion precipitates from iron(II) solutions pale iron(II) hydroxide (pKs = 14,5(amorphous) and 15,1(crystalline), precipitation window = 7,25-9,75), and from iron(III) solutions red-brown iron(III) hydroxide (pKs = 38,6, 39,3 (stagnated), 41,5 (FeOOH, α), 42,7 (Fe2O3, α), precipitation window = 1,52-3,13). Iron(II) compounds oxidize at all pH values readily into iron(III) compounds (see redox section). Iron(II) ions form medium strength hydroxide complexes (lgβ = 4,6, 7,4, 11,0 and 9,6) and iron(III) ions form strong hydroxide complexes (lgβ = 11,81, 22,4, 30,2, 34,4 and 49,7 (lgβ34)).
Zn
Zinc appears in aqueous solutions as colourless hexaaquazinc(II) ions ([Zn(H2O)6]2+). Zinc(II) nitrates aqueous solution is acidic (0,1 M solution pH = 4,98).
The hydroxide ion precipitates white zinc(II) hydroxide (pKs = 15,52(amorphous), 16,24(β1), 16,20(β2), 16,26(γ), 16,15(δ), 16,46(ε), precipitation window = 6,74-9,24), which dissolves in an excess of hydroxide ions (dissolving into a 0,1 M solution requires that [OH–] = 0,694) into hydroxide complexes (lgβ = 5,0, 11,1, 13,6, 14,8 and 5,0(lgβ21)).
In aqueous solutions aluminium appears as colourless hexaaquaaluminium(III) ions ([Al(H2O)6]3+). Aluminium nitrates aqueous solution is strongly acidic (0,1 M solution pH = 2,93).
The hydroxide ion precipitates white gelatinous aluminium hydroxide (pKs = 33,7, precipitation window = 3,10-4,77), which dissolves in an excess of hydroxide ions (dissolving into a 0,1 M solution requires that [OH–] = 0,025) into hydroxide complexes (lgβ = 9,00, 17,7, 25,3, 33,3 and 42,1 (lgβ34)).
Cr
In aqueous solutions chromium(III) appears as grey-silver hexaaquachromium(III) ions ([Cr(H2O)6]3+). Chromium(III) ion functions in an aqueous solution as an almost strong acid (0,1 M Cr(NO3)3 solution pH = 1,15).
The hydroxide ion precipitates grey-green gelatinous chromium(III) hydroxide (pKs =30,2, precipitation window = 4,27-5,93), which dissolves in reagent excess (dissolving into a 0,1 M solution requires that [OH–] = 3,98) into hydroxide complexes (lgβ =10,3, 18,3, 24,0, 28,6 and 37,0 (lgβ34)) which turn the solution green.
Mg
In an aqueous solution magnesium appears as colourless hexaaquamagnesium(II) ions ([Mg(H2O)6]2+. The magnesium ion hydralizes more than other alkaline earth metals and because of this, magnesium nitrates aqueous solution is weakly acidic (0,1 M solution pH = 6,20).
The hydroxide ion precipitates white magnesium hydroxide (pKs = 9,2(active), 11,1(brucite), precipitation window = 9,9-12,4), which does not dissolve in reagent excess into hydroxide complexes (lgβ1 = 2,58).
Ca
Calcium appears in aqueous solutions as colourless hexaaquacalcium(II) ions ([Ca(H2O)6]2+. Calcium hydrolyzes only slightly and therefore calcium nitrates aqueous solution is almost neutral (0,1 M solution pH = 6,76).
The hydroxide ion does not precipitate calcium hydroxide (pKs = 5,29), but atmospheric carbon dioxide causes calcium to precipitate from a basic solution as calcium carbonate. Calcium forms a weak hydroxide complex (lgβ1 = 1,30).
Sr
Strontium appears in aqueous solutions as colourless octaaquastrontium(II) ions ([Sr(H2O)8]2+. Strontium hydrolyzes only slightly and therefore strontium nitrates aqueous solution is almost neutral (0,1 M solution pH = 6,89).
The hydroxide ion does not precipitate strontium hydroxide (solubility = 2,25 g/100 g H2O), but atmospheric carbon dioxide causes strontium to precipitate from a basic solution as strontium carbonate. Strontium forms a weak hydroxide complex (lgβ1 = 0,82).
Ba
Barium appears in aqueous solutions as colourless octaaquabaarium(II) ions ([Ba(H2O)8]2+. Barium hydrolyzes only slightly and therefore barium nitrates aqueous solution is almost neutral (0,1 M solution pH = 6,92).
The hydroxide ion does not precipitate barium hydroxide (pKs = 3,6 (octahydrate)), but atmospheric carbon dioxide causes barium to precipitate from a basic solution as barium carbonate. Barium forms a weak hydroxide complex (lgβ1 = 0,64).
Na
Sodium appears in aqueous solutions as tetraaquasodium ions ([Na(H2O)4]+). Sodiums aqueous solutions are neutral (0,1 M sodium nitrate solution pH = 6,97).
Sodium hydroxide is water soluble (solubility = 100 g/100 g H2O). Sodium also forms a hydroxide complex (lgβ1 = 0,1).
K
Potassium appears in aqueous solutions as tetraaquapotassium ions ([K(H2O)4]+). Potassiums aqueous solutions are neutral (0,1 M potassium nitrate solution pH = 6,98).
Potassium hydroxide is water soluble (solubility = 121 g/100 g H2O). Potassium also forms a hydroxide complex (lgβ1 = 0,0).
NH4+
In aqueous solutions ammonium functions as an acid (pKa = 9,244) and ammonia as a base. When ammonium salts are heated with hydroxide ions, ammonia is released, which will turn wet pH paper held above the test tube blue.
Redox
H2O
As water is oxidized, oxygen gas is realeased, and when reduced, hydrogen gas is released. The reduction-oxidation reaction is seen below.
|
|
O2 + 4H+ + 4e− → 2H2O |
E° = +1,229 V, pH = 0 |
|
|
O2 + 2H2O + 4e− → 4OH− |
E° = +0,401 V, pH = 14 |
|
|
2 H+ + 2e− → H2↑ |
E° = +0,00 V, pH = 0 |
|
|
2H2O + 2e− → H2 +2OH− |
E° = −0,828 V, pH = 14 |
As we can see from the equation, pH has a significant impact on waters redox potential. Water oxidizes most readily in basic solutions and reduces in acidic ones. Pourbaix diagrams (E-pH- diagrams) map the stability of water with dashed lines.
In the standard state many oxidizing agents like potassium permanganate can quickly oxidize water. Permanganate solutions (and other aqueous solutions of oxidizing agents) are however always milder than the standard stat (less than 1 M) and therefore stable. With the Nernst equation it is possible to calculate the reduction potential at a non-standard state. Water can be reduced with, for example, alkali- and alkaline earth metals. Of these especially sodium is used to remove water for organic solvents.
H2O2
Hydrogen peroxide can function as both an oxidizing or reduction agent depending on the conditions. Below is shown the reduction–oxidation reactions of hydrogen peroxide in both acidic and basic solutions.
|
|
O2↑ + 2H+ + 2e− → H2O2 |
E° = 0,695 V, acidic solution |
|
|
O2↑ + H2O +2e− → OH− + HO2− |
E° = −0,0649 V, basic solution |
|
|
H2O2 + 2H+ + 2e− → 2H2O |
E° = +1,763 V, acidic solution |
|
|
HO2− + H2O + 2e− → 3OH− |
E° = +0,867 V, basic solution |
From the potentials we can see, that hydrogen peroxide has a very positive reduction potental in acidic solutions, which means it is a very strong oxidizer in acidic solutions. Even in basic solutions hydrogen peroxide is very good oxidizing agent.
Strong oxidizing agents can oxidize hydrogen peroxide into oxygen gas, in which case hydrogen peroxide will function as a reducing agent. Often in these scenarios the oxidizing agents reduced form oxidizes back because of the strong oxidizing ability of hydrogen peroxide. E.g. below.
|
|
Br2 + H2O2 → 2H+ + 2Br− + O2↑ |
E° = +0,392 V |
|
|
2Br− + H2O2 + 2H+ → Br2 + 2H2O |
E° = +0,676 V |
|
|
2H2O2 → O2 + 2H2O |
E° = +1,068 V, acidic solution |
The last reaction is the sum of the first two, the disproportionation reaction of hydrogen peroxide, and it can also occur without bromine. Bromine functions in this scenario as a catalyst. Disproportionation also happens in basic solutions, but then the voltage is slightly less positive (E° = +0,9319 V). Hydrogen peroxide is then very unstable regardless of pH, but without a catalyst it decomposes very slowly.
In practice hydrogen peroxide functions in acidic solutions often as a reducing agent and in basic solutions as an oxidizing agent (e.g. below). In neutral solutions a large group of elements and compounds instead function as catalysts for the decomposition of hydrogen peroxide. This behaviour is the consequence of a multitude of thermodynamic and kinetic factors.
|
|
2MnO4− + 5H2O2 + 6H+ → 2Mn2+ + 5O2↑ + 8H2O |
E° = 1,51 V – 0,695 V = 0,815 V, acidic solution |
|
|
HO2− + Mn(OH)2↓ → MnO2↓ + OH− + H2O |
E° = 0,867 V + 0,05 V = 0,917 V, basic solution |
Other reactions