Appendix 1. Hydrated salts, metal–aqua complex ions and their relative acidity, salt hydrolysis

(1) Introduction

Examples of hydrated salts are described and water of crystallisation explained.

The reason why many salts of metals do not dissolve to give neutral solutions i.e. many are acidic because hexa–aqua ions of transition metals and other metals like aluminium behave as Bronsted Lowry acids in a process sometimes described as salt hydrolysis.

The effect of ion charge and the ionic radius on the acidity of the hydrated ions is described and explained by considering the polarising power of the central metal ion.

Examples are described of both 3d-block transition hydrated metal ions and non-transition hydrated metal complex ions

(not necessarily just transition metal ions)

(2) Hydrated salts and water of crystallization

• All metal ions in solution are 'associated' with water. The water molecules can also be weakly bonded or more strongly as a ligand to form a complex ion, and these can also present in solid 'hydrated' salts on crystallisation e.g.

  • FeSO₄.7H₂O_(s), CoCl₂.6H₂O_(s), CuSO₄.5H₂O_(s) etc.

    • Iron(II) sulfate heptahydrate, cobalt(II) chloride hexahydrate and copper(II) sulfate pentahydrate.

    • The above crystals contain 7, 6 and 5 molecules of water of crystallisation respectively.

    • A hexa–aqua ion is present in the first two, [M(H₂O)₆]²⁺ (M = Fe, Co)

  • What is the difference between water of crystallization and the co–ordinated water molecules bonded to the central metal ion via the dative covalent bonds?

    • There may or not be a difference!

    • BUT, what ever, the water molecules are chemically bonded in some way to the positive metal ion and the negative anion.

    • The water of crystallisation is the total number of water molecules incorporated into the crystal structure irrespective of the nature of the chemical bonding involved OR any intermolecular associations.

    • The case of copper(II) sulfate pentahydrate is considered below, where 4/5 water molecules are ligands and the 5th water molecule is held in place by hydrogen bonding.

    • However in the case of magnesium chloride (magnesium isn't a transition metal), the crystal lattice consists of hexaaqua magnesium ions and chloride ions.

    • Cl^– <=> [Mg(H₂O)₆]²⁺ <=> Cl^– are all bonded together by electrostatic attraction in the crystal lattice.

    • No hydrogen bonding is involved, and the number of molecules of water of crystallisation is equal to the ligand coordination number of the central metal ion.

    • Just in passing, a complex ion is a complex ion, it doesn't matter if the central metal ion isn't a 3d block transition metal ion etc. Its still the same sort of structure with the same sort of bonding and shape.

  • In the case of copper(II) sulfate, 4 water molecules are covalently bonded to form effectively a square planar complex ion,

    • [Cu(H₂O)₄]²⁺ and the 5th water molecule H₂O is hydrogen bonded to this ion and hydrogen bonded to a neighbouring sulfate ion

    • SO₄^2– thus helping to hold the crystal lattice together, though the main force of attraction is the electrostatic attraction between copper complex ion and the sulfate ion.

    • Therefore, the water of crystallisation number doesn't equal the co–ordination number of the central metal ion.

    • ... H–O–H ... are the three components of the crystal structure of copper(II) sulfate pentahydrate, and all three are linked by hydrogen bonds. The full structure is a bit complicated to draw but the 5th and 6th octahedral positions of the Cu²⁺ ion are occupied by oxygen atoms of the sulfate ion and the 5th water molecule is held in position by hydrogen bonding.

    • However, this blue crystal lattice is readily broken down on heating, a classical demonstration of a reversible reaction, since the white anhydrous solid turns blue on adding water (a simple test for water.

    • CuSO₄.5H₂O_(s) CuSO_4(s) + 5H₂O_(g/l)

  • So three words–phrases to know ...

    • Water of crystallisation – the molecules of water incorporated into the crystal structure either by acting as a ligand to a metal ion (e.g. four in hydrated CuSO₄) or just hydrogen bonded into the lattice (one in hydrated CuSO₄), so a total of five molecules of water of crystallisation as discussed above.

      • The term originates from analysis of salts crystallised from water.

    • Hydrated salt – lattice contains molecules of associated water e.g. water of crystallisation in the case of salts.

    • Anhydrous salt – devoid of water molecules in the crystal lattice e.g. dehydrated salts or those that cannot crystallise with water of crystallisation from a concentrated aqueous solution.

    • It should be pointed out that the term anhydrous is used quite generally to mean a substance has had all water–moisture removed from it.

  • See also

  • Calculations of water of crystallization – % composition & simple experimental determination including practice questions with worked out answers

(3) Acid-base theories

• Lewis acid–base theory reminders:

  • A base is an electron pair donor and an acid is an electron pair acceptor.

  • Ligands like water, can donate a pair of non–bonding electrons (lone pair) into a vacant orbital of a central metal ion and so dative covalent (co–ordinate) bonds hold a complex together.

  • The central metal ion with vacant bonding orbitals can act as a Lewis acid.

  • Ligands act as Lewis bases by electron pair donation to form the metal–ligand bond.

• Bronsted–Lowry acid–base theory reminders

  • (essentially a sub–set of Lewis Theory)

  • For more details see Equilibria Part 5

  • A base is a proton acceptor.

    • This is via an electron lone pair on the base (a Lewis base is a lone pair donor).

    • e.g. NH₃, HCO₃^–, OH^– etc.

  • An acid is proton donor.

    • This involves a heterolytic breakage of an X–H bond (a Lewis acid is an electron pair acceptor).

    • e.g. HCl, HCO₃^–, H₂SO₄, CH₃COOH etc.

(4) Salt hydrolysis and the acidity of hexaaqua ions

• Many hexa–aqua complex ions can undergo acid–base reactions with water to produce solutions of pH less than 7.

• Usually group 2, 3 and transition metal ions.

• The positive central metal ion polarises a water molecule ligand, releasing a proton, H⁺.

• In the deprotonation reaction the proton transfers to water and the overall charge on the complex falls by 1 unit since the H₂O – H⁺ = OH^–, i.e. one of the ligands is now a hydroxide ion instead of the original water molecule.

• In these reactions the hydrated ions act as Bronsted Lowry acids and water acts as a Bronsted–Lowry base.

• These reactions are examples of what is termed 'salt hydrolysis' because the metal ion (of usually a salt) reacts with water to give, in this case, two products.

• You can get salt hydrolysis with e.g. carbonate salts via the carbonate ion acting as a base (e.g. aqueous sodium carbonate is alkaline), but they are discussed elsewhere (see Equilibria section 6.1.3).

• These are acid–base reactions NOT redox reactions, even if they involve transition metal ions – there is NO change in oxidation state of the metal!

• e.g. for hexaaquametal(II) ions ...

• [M(H₂O)₆]²⁺_(aq) +  H₂O_(l) [M(H₂O)₅(OH)]⁺_(aq) + H₃O⁺_(aq) 

  • e.g. when M = Mn, Fe, Co, Ni, Cu, Mg etc. gives a very weak acid solutions with pH's just less than 7.

  • The hydrated M²⁺ ions are not as acidic as the hydrated M³⁺ ions - which have a higher charge and usually a smaller ionic radius.

  • Ti(II), V(II) and Cr(II) M²⁺ ions are redox unstable in the presence of air, but theoretically their salts give very weakly acid solutions.

  • They are usually prepared by zinc–acid reduction from higher oxidation states.

  • Note that M⁺ hexaaqua ions show little acidity due to having the smallest central metal ion charge i.e. salts of group 1 alkali metals with strong mineral acids e.g. in the chloride, nitrate and sulfate salt solutions, the [M(H₂O)₆]⁺_(aq) where M = Li, Na and K etc. give ~pH 7 neutral solutions.

• Salt hydrolysis and acidity of hexaaquametal(III) ions

• The hexaaqua ions of Al³⁺ and Fe³⁺ can donate 1 to 3 protons to water molecules giving acidic solutions e.g. the first two, and consecutive proton donations can be written as ...

• 1. [M(H₂O)₆]³⁺_(aq) + H₂O_(l) [M(H₂O)₅(OH)]²⁺_(aq) + H₃O⁺_(aq) 

• 2. [M(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) [M(H₂O)₄(OH)₂]²⁺_(aq) + H₃O⁺_(aq) 

  • e.g. M = Ti, V, Cr, Fe, Al etc. give very weak acids solutions (but generally stronger than for M²⁺) of pH's in the 3–5 region.

  • In the presence of alkali, OH^–, in removing H₃O⁺ ions, the  equilibrium moves more to the right and more protons are lost from the complex in stages until the hydroxide precipitate is formed e.g. for iron(III), chromium(III) or aluminium.

  • [M(H₂O)₆]³⁺_(aq) + 3OH^–_(aq) [M(H₂O)₃(OH)₃]⁰_(s) + 3H₂O_(l) 

  • Some of the M³⁺ hydroxides are amphoteric and dissolve in excess strong alkali (1.) or strong acid (2.) e.g. to eventually form for chromium(III) or aluminium, 1. the soluble hexa–hydroxo complex anion or 2. the original hexa–aqua cation.

  1. [M(H₂O)₃(OH)₃]⁰_(aq) + 3OH^–_(aq)  [M(OH)₆]^3–_(aq)+ 3H₂O_(l) 

  2. [M(H₂O)₃(OH)₃]⁰_(s) + 3H₃O⁺_(aq)  [M(H₂O)₆]³⁺_(aq) + 3H₂O_(l)

  3. Reactions 1. and 2. apply to ions such as Al³⁺ or Cr³⁺ whose insoluble hydroxides are amphoteric - they dissolve in both strong acids or strong alkalis.

• As a general rule the greater the polarising power of the central metal ion, the lower the pH of the resulting aqueous solution, i.e. the acid–base equilibrium is shifted more to the right causing an increase in acidity of the solution.

  • This effect and process facilitated by the central metal ion on one water ligand molecule can be envisaged for one of the water molecule ligands as ...

    • [M–O–H₂]ⁿ⁺ ==> [M–O–H]^(n–1)+ + H⁺

    • (conceptually think of a proton transferred to a water molecule)

    • One of the O–H bond pairs is 'attracted' onto the oxygen atom by the electric field effect of the central metal ion of charge n+, allowing proton transfer to the base water.

  • Polarising power is a function of ionic charge (n+)/ionic radius (r) ratio

  • i.e. polarising power of the central metal ion is a function of n+/r

  • Therefore ...

  • the greater the charge on the central metal ion (n+), the more acidic the hexaaqua ion, hence a lower pH solution,

  • and the smaller the ionic radius of the central metal ion of the complex, the more acidic the hexaaqua ion, hence a lower pH solution,

  • and so these factors increase the electric field effect of the central metal ion on the surrounding ligand bonded water molecules ...

  • both increasing charge, or decreasing the central cation radius intensify the electric field polarising effect on a water ligand which facilitates proton donation from the complex ion to a free water molecule.

• The acidity of the hexaaqua ions M³⁺_(aq) due to the polarising influence of the central highly charged M³⁺ ion accounts for the lack of stability-existence of ...

  • e.g. aluminium carbonate, iron(III) carbonate or chromium(III) carbonate, which don't exist as far as I know?,

  • whereas MgCO₃ , ZnCO₃ and FeCO₃ etc. with the less polarising M²⁺ ion exist and although insoluble, they are relatively stable in the presence of water

  • It also accounts for why you see bubbles of carbon dioxide when (i) hydrogencarbonates or (ii) carbonates are mixed with aluminium chloride, iron(III) chloride or chromium(III) chloride solutions.

  • You can write a variety of Bronsted-Lowry acid-base equations to illustrate this e.g.

  • (i) donating one proton to a hydrogencarbonate ion releasing carbon dioxide

  • [Al(H₂O)₆]³⁺_(aq) + HCO₃^–_(aq) [Al(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)

  • [Fe(H₂O)₆]³⁺_(aq) + HCO₃^–_(aq) [Fe(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)

  • [Cr(H₂O)₆]³⁺_(aq) + HCO₃^–_(aq) [Cr(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)

  • (ii) donating two protons to a carbonate ion releasing carbon dioxide

  • 2[Al(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) 2[Al(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)   

  • 2[Fe(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) 2[Fe(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)

  • 2[Cr(H₂O)₆]³⁺_(aq) + CO₃^2–_(aq) 2[Cr(H₂O)₅(OH)]²⁺_(aq) + H₂O_(l) + CO_2(g)   

  • There several other legitimate permutations based on these equations.

(5) Amphoteric nature of some complex ions and hydroxide precipitate formation and dissolution

• Amphoteric nature of ions some hydrated Al³⁺ or Cr³⁺ ions and their hydroxide precipitates.

  • In the above chemistry the acidic nature of the hexa–aqua ions was emphasised, BUT as soon as one proton has been lost the resulting complex ion can then act as a base.

  • e.g. in solutions of the weakly acidic Cr³⁺ or Al³⁺ species with excess strong acid the hexaaqua metal ion would predominate (M = Cr or Al).

  • [M(H₂O)₄(OH)₂]⁺_(aq) + 2H⁺_(aq) [M(H₂O)₆]³⁺_(aq)

  • or more correctly written as the full Bronsted=Lowry acid-base equation

  • [M(H₂O)₄(OH)₂]⁺_(aq) + 2H₃O⁺_(aq) [M(H₂O)₆]³⁺_(aq)  +  2H₂O(l)

• The effect of step-wise adding alkali (e.g. from NaOH(aq)) to an initial aqueous solution of the of the hexaaqua ion - two sequences illustrated below.

1 [Cr(H2O)6]3+(aq)	2 [Cr(H2O)5(OH)]2+(aq)	3 [Cr(H2O)4(OH)2]+(aq)
4 [Cr(H2O)3(OH)3]0(s)	5 [Cr(H2O)2(OH)4]-(aq)	6 [Cr(H2O)(OH)5]2-(aq)
7 [Cr(OH)6]3-(aq)	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 with an overall charge changing from 3+ to 3–. Only the stage 4. complex is insoluble [Cr(OH)3(H2O)3]. Equations for 1 to 7 illustrated below. From 1 to 7 happen as you add more alkali, increasing pH and the OH– concentration, removing protons from the chromium(III) complex. From 7 back to 1 represents what happens when you add acid, decreasing pH, increasing H+/H3O+ concentration and protonating the complex.