Metal Aqua Ions

Created by

khadija

Cards (48)

in aqueous solutions, without the presence of other ions, metal ions exist as metal aqua ions
these have a central metal ion with six water ligands
octahedral shape
water acts as a lewis base - lone pair donor
metal ions act as lewis acids - lone pair acceptor
when salts are crystallised from a solution, the metal aqua ions are present in the crystals
complex colours:
Cu2+ - blue
Co2+ - pink
Fe2+ - green
V2+ - green
Fe3+ - pale violet but appears brown/orange irl
Cr3+ - violet
Al3+ - colourless
Fe3+ should be pale violet but appears brown/orange due to the small amount of Fe2+ formed by hydrolysis
reactions of metal aqua ions:
hydrolysis - loss of H+ from H2O ligands (O-H bond in H2O ligand breaks)
substitution - replacement of H2O by other ligands (metal-ligand bond breaks)
redox - metal changes oxidation state (gain or loss of electrons)
hydrolysis:
in solution metal aqua ions lose H+ ions from one or more H2O ligands
acidity of solutions of aqueous metal ions:
in solution metal aqua ions lose H+ ions from one or more H2O ligands
in water the main species is still [M(H2O)6]n+ but some hydrolysis happens making the solution acidic
the higher the charge on the metal, the more acidic the solution
higher charge on metal = ion is more attracted to water = more acidic solution
to lose H+ the O-H bond has to break
to break the bond the electron pair is pulled towards the O, making the bond more polar and making the H more acidic
the higher the charge on the metal ion and/or the smaller the metal ion, the stronger the pull on the electrons towards the O
this breaks the O-H bond
pH of aqua ion solutions:
1+ has pH 7
2+ has pH 6
3+ has pH 3
4+ has pH 0
reaction with bases:
if a base is added to metal aqua ions hydrolysis may take place - the base removes the H+ and shifts the equilibria to the right
common bases - OH-, NH3, CO3 2-
other reactions e.g. ligand substitution could also take place
if hydrolysis happens the insoluble neutral complex is formed as a precipitate
if an excess of the base is added the insoluble complex may react further
reaction with bases:
metal (II) aqua ions react with NaOH to form precipitates of metal(II) hydroxides by hydrolysis
metal(III) aqua ions react with NaOH to form precipitates of metal (III) hydroxides by hydrolysis
Al(OH)3 reacts further with excess NaOH to form [Al(H2O)2(OH)4]-
reaction with bases:
metal (II) aqua ions react with NH3 to form precipitates of metal (II) hydroxides by hydrolysis
metal (III) aqua ions reacts with NH3 to form precipitates of metal (III) hydroxides by hydrolysis
Cu2+ reacts further with excess NH3 to form [Cu(H2O)2(NH3)4]2+ by ligand substitution
reaction with bases:
metal(II) aqua ions react with CO3 2- ions to form precipitates of metal (II) carbonates by precipitation
metal (III) aqua ions react with CO3 2- ions to form bubbles of CO2 and precipitates of metal (III) hydroxides by hydrolysis
3+ aqua ions are acidic enough to react in an acid-base reaction with carbonate ions by hydrolysis but 2+ aqua ions aren't acidic enough
reactions of aqua ions in bases:
hydrolysis with OH-
ligand substitution with NH3
precipitation with CO3 2-
6 water molecules form co-ordinate bonds with the metal ion
a lone pair on the oxygen allows this bond to form
(1) hydrolysis of metal aqua ions:
when 3+ metal aqua ions like Al3+ or Fe3+ are hydrolysed an equilibrium is established
if OH- ions are added the equilibrium shifts right as H+ ions are removed
(2) hydrolysis of aqua ions:
further hydrolysis takes place in the new equilibrium
OH- ions remove H+ ions
equilibrium shifts to the right
(3) hydrolysis of aqua ions:
further hydrolysis leads to the formation of a neutral complex that is solid
this forms a precipitate in solution
further hydrolysis can be seen with 2+ aqua ions also, such as Cu2+ and Fe2+
happens in 2 steps instead of 3 because only two water ligands need to be deprotonated to form a neutral complex
further hydrolysis of 2+ aqua ions:
step 2
forms an insoluble metal hydroxide
A) M(H2O)5(OH)
B) +
C) M(H2O)4(OH)2
amphoteric metal hydroxides:
can act as an acid or a base
dissolve in both excess acid AND base
aluminium hydroxide is amphoteric - when a base is added it acts as a bronsted-lowry acid and donates H+ ions to react with OH- and dissolves
when an acid is added it acts as a Bronsted-Lowry base and accepts H+ ions to form H3O+ in solution and dissolves
aluminium hydroxide reacting with a base:
A) OH-
B) [Al(H2O)2(OH)4]
C) -
D) H2O
aluminium hydroxide reacting with acid:
A) [Al(H2O)6]
B) 3+
C) 3H2O
in solution ammonia exists in the equilibrium:
NH3 + H2O <--> NH4+ + OH-
this reaction produces OH- ions so adding small quantities of NH3 to a metal aqua ion produces the same metal hydroxides as adding NaOH
test tube reactions:
add a sample of the unknown metal ion solution into 3 test tubes
drop by drop add NaOH to tube 1 and observe any changes, add more NaOH to see if an excess produces a change, record observations
repeat step 2 with a second test tube, using ammonia solution instead of NaOH, record observations
in test tube 3, add sodium carbonate drop by drop and record observations
test tube reactions - safety:
wear gloves and goggles as some chemicals used can cause irritation
ammonia gives off pungent fumes - conduct in a fume cupboard
Al3+ with NaOH:
Al3+ ions in solution in a test tube
add NaOH
precipitate forms
add an excess of NaOH
precipitate dissolves
ammonium hydroxide is the only one that dissolves in excess NaOH because it is amphoteric
Cu2+ with ammonia:
Cu2+ ions in solution in a test tube
add NH3
precipitate forms
add excess NH3
precipitate dissolves and a dark blue colour is produced
copper hydroxide is the only precipitate that dissolves in excess NH3 because there is a ligand substitution reaction
Fe2+ and Cu2+ react with sodium carbonate to form precipitates only
Al3+ and Fe3+ react with carbonates to form precipitates and CO2, so bubbles are also observed
metal ions dissolved in water only have the formula [M(H2O)6]n+ which can be simplified to M n+
Fe2+ in NaOH:
green solution to green precipitate
[Fe(H2O)6]2+ + 2OH- --> [Fe(H2O)4(OH)2] + 2H2O
in excess OH- no visible further reaction
Cu2+ with NaOH:
blue solution to blue precipitate
[Cu(H2O)6]2+ + 2OH- --> [Cu(H2O)4(OH)2] + 2H2O
excess NaOH - no further reaction
Fe3+ with NaOH:
pale violet (orange) solution to red/brown precipitate
[Fe(H2O)6]3+ + 3OH- --> [Fe(H2O)3(OH)3] + 3H2O
excess NaOH - no visible further reaction
Al3+ with NaOH:
colourless solution to white precipitate
[Al(H2O)6]3+ + 3OH- --> [Al(H2O)3(OH)3] + 3H2O
excess NaOH - precipitate dissolves, colourless solution again
[Al(H2O)3(OH)3] + OH- --> [Al(H2O)2(OH)4]- + H2O
Fe2+ with NH3:
green solution to green precipitate
hydrolysis
[Fe(H2O)6]2+ + 2NH3 --> [Fe(H2O)4(OH)2] + 2NH4+
excess NH3 - no visible further reaction
Cu2+ with NH3:
blue solution to blue precipitate
hydrolysis
[Cu(H2O)6]2+ + 2NH3 --> [Cu(H2O)4(OH)2] + 2NH4+
excess NH3 - precipitate dissolves, dark blue solution appears
substitution
[Cu(H2O)4(OH)2] + 4NH3 --> [Cu(H2O)2(NH3)4]2+ + 2H2O + 2OH-
Fe3+ with NH3:
pale violet (orange) solution to red/brown precipitate
hydrolysis
[Fe(H2O)6]3+ + 3NH3 --> [Fe(H2O)3(OH)3] + 3NH4+
excess NH3 - no visible further reaction
Al3+ with NH3:
colourless solution to white precipitate
hydrolysis
[Al(H2O)6]3+ + 3NH3 --> [Al(H2O)3(OH)3] + 3NH4+
excess NH3 - no visible further reaction
Fe2+ with Na2CO3:
green solution to green precipitate
precipitation
[Fe(H2O)6]2+ + CO3 2- --> FeCO3 + 6H2O
Cu2+ with Na2CO3:
blue solution to blue precipitate
precipitation
[Cu(H2O)6]2+ + CO3 2- --> CuCO3 + 6H2O