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Acid-Base Solubility Equilibrium Ionic Bond Acid Base theories pH values equilibrium constant common ion effect solubility cuprite
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DM Sherman, University of Bristol
Geochemistry, DM Sherman University of Bristol
The O-H bond is “polar covalent”; from the difference in electronegativity, the OH bond is 39 % ionic. This gives the water molecule a large dipole moment.
DM Sherman, University of Bristol
Water molecules are attracted to each other via the hydrogen bond (dipole-dipole interaction): Because of the hydrogen bond, water has a much higher boiling point than that for molecules of similar mass (e.g., CH 4 ).
The dipole moment of the water molecule means that water molecules will be attracted to ions in solution. This is why ionic compounds will dissolve in water.
DM Sherman, University of Bristol
The Lewis acid-base theory gives a more generalized picture: CO 2 + H 2 O = HCO 3 -^ + H+ The generalized acid-base reaction can be viewed as Acid + Base → Weaker Acid + Weaker Base Acid Base Base Acid Fe3+^ + H 2 O = Fe(OH)2+^ + H+
In aqueous systems, all acids stronger that H 2 O generate excess H+^ ions (or H 3 O+); all bases stronger than H 2 O generate excess OH-: ( or, Fe3+^ + H 2 O → Fe(OH)2+^ + H+^ ) ii) CO 3 2-^ + H 2 O → HCO 3 -^ + OH- i) Fe(H 2 O) 6 3+^ + H 2 O → Fe(H 2 O) 5 (OH)2+^ + H 3 O+ Hence, in aqueous solutions, the Lewis acid-base picture becomes the Bronsted acid-base picture.
DM Sherman, University of Bristol
Most minerals are salts of weak acids and strong bases. Such minerals react with water to raise the pH: 2KAlSi 3 O 8 + 2H 2 O = Al 2 Si 2 O 5 (OH) 4 + 2K+^ + 2OH- 2Mg 2 SiO 4 + 3H 2 O = Mg 3 Si 2 O 5 (OH) 4 + Mg2+^ + 2OH- CaCO 3 + H 2 O = HCO 3 -^ + Ca2+^ + OH- A few minerals are salts of weak bases and strong acids. Such minerals react with water to lower the pH: KFe 3 (SO 4 ) 2 (OH) 6 + 3H 2 O = K+^ + 3Fe(OH) 3 + 2SO 4 2-^ + 3 H+
Hence, regardless of the nature of the acid or base:
DM Sherman, University of Bristol
The “proton ladder” gives a useful picture of the relative strengths of acids and bases in aqueous solutions.
DM Sherman, University of Bristol
In very dilute aqueous solutions, the ions do not interact with each other. Moreover, the presence of the ions has no effect on the bulk structure of water. We will take this as the standard state for ions in aqueous solution: a 1 molal solution with the properties of infinite dilution. In dilute solutions: Activity = concentration
With increasing concentration, ions start to interact with each other. The chemical potential of an ion will depend upon the concentration of the other ions. Ion-ion interactions may be specific (e.g., ion- pairs or complexes ) or non-specific (long-range coulombic interactions).
DM Sherman, University of Bristol
CaSO 4 (s) = Ca+2^ + SO 4 - PbSO 4 (s) = Pb+2^ + SO 4 - = 6.3 x 10- = 2.0 x 10-
DM Sherman, University of Bristol
Because of charge balance , we have [Ca] + [Pb] = [SO 4 ]. But since CaSO 4 is much more soluble PbSO 4 , [Ca] >> [Pb] or [Ca] + [Pb] ≈ [Ca]. Hence [Ca] ≈ [SO 4 ] Since KCaSO4 = [Ca][SO 4 ] = 6.3 x 10- we get [SO 4 ] = 7.9 x 10-
Since KPbSO4 =[Pb][SO 4 ] = 2.0 x 10- we get [Pb] = 2.5 x 10- If CaSO 4 was absent then [Pb] = [SO 4 ] So that, [Pb] = 1.4 x 10-
DM Sherman, University of Bristol
Cl complexation changes the solubility of Cu 2 O by 8 orders of magnitude.
In addition to the simple solubility equilibrium: Fe3+^ + H 2 O = FeOH+2^ + H+^ pK = 2. Fe3+^ + 2H 2 O = Fe(OH) 2 +^ + 2H+^ pK = 5. Fe3+^ + 3H 2 O = Fe(OH) 30 + 3H+^ pK = 12. Fe3+^ + 4H 2 O = Fe(OH) 4 -^ + 4H+^ pK = 21. FeOOH + 3H+^ = Fe+3^ + 2H 2 O pK = -0. We also need to account for the hydrolysis of Fe3+:
DM Sherman, University of Bristol
By taking the “p-function” (i.e., -log) of the equilibrium expressions, we get: p[FeOH+2] = pK 2 - pH + p[Fe+3] = 1.66 + 2pH (2) p[Fe+3] = pK 1 + 3pH = -0.53 + 3pH (1) p[Fe(OH) 2 +] = pK 3 - 2pH + p[Fe+3] = 5.14 + pH (3) p[Fe(OH) 3 ] = pK 4 - 3pH + p[Fe+3] = 12.03 (4) p[Fe(OH) 4 - ] = pK 5 - 4pH + p[Fe+3] = 21.1 - pH (5) Plotting the p[Fe(OH)n] (equations 1-5) gives a map of the saturation field of FeOOH and the dominant aqueous Fe complexes at each pH.
Notice that [Fe]tot reaches a minimum at the pH of seawater. 1 2 3 4 5
