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Aqueous ion solution

AQA A-level inorganic chemistry transition metals year 13

TermDefinition
How most inorganic compounds of transition metals react when aqueous Depending on the strength of the metal-ligand bond, most anions will dissociate & be replaced with water ligands
Ligand exchange Ligands may be swapped around for one another, usually resulting in a colour change, this is due to thermodynamics & kinetics
Neutrally charged monodentate ligand exchange Water & ammonia are both neutrally charged monodentate ligands which means that they can be substituted for each other in solution; neither the charge of the complex nor the coordination number will change. This is reversible
Chelate effect When multiple monodentate ligands are substituted for bi or multidentate ligands (such as en of EDTA) the new complex is more thermodynamically stable which pushes equilibrium very far to the right making it difficult to reverse
EDTA & en meaning EDTA = ethylenediametetraacetic acid en = ethylenediamine
How chelation is thermodynamically favourable Negative ΔG due to entropy increasing from releasing many water ligands from coordination in exchange for a few ligands & breaking stronger chelated bonds isn’t energetically favourable compared to monodentate ligands
Biological example of chelation The molecule porphyrin is a polydentate ligand which forms 4 coordinate bonds with Fe2+, 1 bond with the globin subunit, & 1 with water which is substituted for oxygen in the lungs
Hydrolysis H+ ion dissociates from H2O ligand leaving an OH- ligand in its place
Why hydrolysis occurs and what affects it Ligand is polarised by metal ion through its coordinate bond causing electron density to shift towards it which strains the O-H bond causing deprotonation. Smaller atomic radius & higher charge causes greater polarisation leading to lower pH
How equilibrium affects aqueous ion hydrolysis One equilibrium has moved completely to the right for the first ligand hydrolysis, a new equilibrium is established for the second & so on
Lewis acids & bases rules Defines acids as electron pair acceptors & bases as electron pair donors, this means there can be acids without protons (such as AlCl3). Also, if any Lewis base other than water coordinates to the metal ion it does reduce the acidity
Why complexes form precipitates when neutralised Once there is an equal number of hydroxide groups to the oxidation state of the ion, it is no longer polar which makes it water insoluble. However, it can be redissolved in dilute acid which can regenerate the water ligands & polarise it
Amphoteric hydroxides Hydroxides with both acidic & basic properties (e.g., Al(H2O)3(OH)3) since they can react with protons to regenerate water ligands & form a positively charged complex or be hydrolysed further into a negatively charged complex. These are both water soluble
Reactions of aqueous ions: Hydroxides Removes a proton from aqueous ions by hydrolysis causing the charge to decrease, potentially enough to form a precipitate
Reactions of aqueous ions: Ammonia For most complexes just causes hydrolysis by acting as a weak base but may sometimes be substituted with copper if in excess. This is because while water & ammonia are similar sizes, ammonia is a slightly better ligand
Reactions of aqueous ions: Carbonates 1+ & 2+ ions replace all of their water ligands with the carbonate ion forming a metal carbonate. Since 3+ complexes are more acidic, carbonate ions just cause hydrolysis into hydroxides with effervescence from carbon dioxide gas
Redox/redox potential When metal cations change oxidation state through oxidation or reduction. Redox potential is the tendency of metal cations to change oxidation state
Created by: Study_B
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