1.1. Reactions of Inorganic Compounds in Aqueous Solutions 1.2. Difference between Fe2+ and Fe3+ acidity
Dissolving Iron Nitrate in water Fe3+ is more acidic than Fe2+
Water molecules cluster around the iron so in aqueous Because: Fe3+ has a higher charge, so there is a higher
solutions it actually exists as [Fe(H20)6]2+ charge density than Fe2+, making it more strongly
polarising
The water molecules act as ligands forming 6 co-ordinate This means it draws electrons from the oxygen
bonds = octahedral which is co-ordinately bonded to it
This weakens the OH bonds
The same happens with an Iron (III) salt, except the complex is [Fe(H20)6]3+ Causing H+ ions to be more readily lost in the
solution
These are called aqua ions Decreasing the pH, making Fe3+ more acidic
**Rule – M3+ are generally more acidic than M2+
1.6. Distinguishing Iron Ions
1.3. Theories of Acidity Hydrolysis Reaction – using water molecules to break compounds
down (the water molecule is likely to break too)
Fe2+ and Fe3+ exist in aqueous solutions as octa
Lavoisier (1777) – said all acids contained oxygen, HCl disproved theory [Fe(H2O)6]2+ is pale green
Davy (1816) – said all acids contained hydrogen, doesn’t explain why [Fe(H20)6]3+ is pale brown when dilute,
hydrogen is important
Liebig (1838) – defined acids as substances containing hydrogen which Test – add dilute alkali, precipitates hydroxides w
could be replaced by a metal Inorganic Chemistry:
Fe3+ + 3OH- get → Fe(H2O)3(O
Arrhenius (1887) – acids produced H+ ions Reactions of Aqueous Ions
Fe2+ + 2OH- get → Fe(H2O)4(O
Bronsted Lowry – defined as substance that can donate a proton, base can 1.5. Reactions with CO3
accept a proton, difficulty with acids that do not contain
hydrogen (AlCl3) Iron (III) Carbonate doesn’t exist but Iron (II) Carbonate does
Lewis Theory – acids are electron pair acceptors and bases as electron pair Due to greater acidity of Fe3+
donors Carbonate can remove protons from [Fe(H2O)6]3+
but can’t remove them from the Fe2+ complex
1.4. Acid-Base Reactions of M2+ and M3+ ions
2[Fe(H2O)6]3+ +3CO32- → 2[Fe(H2O)3(OH)3] + 3CO2 + 3H2O 1.7. Amphoteric Hydroxides
[Fe(H2O)6]2+ + CO32- → FeCO3 + 6H2O Amphoteric – showing both acidic and basic p
**Rule – Carbonates of M2+ exist, while M3+ carbonates do not Aluminium Hydroxide is example, reacts with
Al(H2O)3(OH)3 + 3HCl → A
Reacts with acid, therefore
The neutral metal hydroxide formed is essentially M(OH)3, it’s
Ammonia has the same effect as hydroxide ions in Al(H2O)3(OH)3 + OH- → [Al
uncharged, insoluble and forms a ppt
removing protons – both basic
The neutral metal hydroxide formed is essentially M(OH)2, it’s
uncharged, insoluble and forms a ppt
Dissolving Iron Nitrate in water Fe3+ is more acidic than Fe2+
Water molecules cluster around the iron so in aqueous Because: Fe3+ has a higher charge, so there is a higher
solutions it actually exists as [Fe(H20)6]2+ charge density than Fe2+, making it more strongly
polarising
The water molecules act as ligands forming 6 co-ordinate This means it draws electrons from the oxygen
bonds = octahedral which is co-ordinately bonded to it
This weakens the OH bonds
The same happens with an Iron (III) salt, except the complex is [Fe(H20)6]3+ Causing H+ ions to be more readily lost in the
solution
These are called aqua ions Decreasing the pH, making Fe3+ more acidic
**Rule – M3+ are generally more acidic than M2+
1.6. Distinguishing Iron Ions
1.3. Theories of Acidity Hydrolysis Reaction – using water molecules to break compounds
down (the water molecule is likely to break too)
Fe2+ and Fe3+ exist in aqueous solutions as octa
Lavoisier (1777) – said all acids contained oxygen, HCl disproved theory [Fe(H2O)6]2+ is pale green
Davy (1816) – said all acids contained hydrogen, doesn’t explain why [Fe(H20)6]3+ is pale brown when dilute,
hydrogen is important
Liebig (1838) – defined acids as substances containing hydrogen which Test – add dilute alkali, precipitates hydroxides w
could be replaced by a metal Inorganic Chemistry:
Fe3+ + 3OH- get → Fe(H2O)3(O
Arrhenius (1887) – acids produced H+ ions Reactions of Aqueous Ions
Fe2+ + 2OH- get → Fe(H2O)4(O
Bronsted Lowry – defined as substance that can donate a proton, base can 1.5. Reactions with CO3
accept a proton, difficulty with acids that do not contain
hydrogen (AlCl3) Iron (III) Carbonate doesn’t exist but Iron (II) Carbonate does
Lewis Theory – acids are electron pair acceptors and bases as electron pair Due to greater acidity of Fe3+
donors Carbonate can remove protons from [Fe(H2O)6]3+
but can’t remove them from the Fe2+ complex
1.4. Acid-Base Reactions of M2+ and M3+ ions
2[Fe(H2O)6]3+ +3CO32- → 2[Fe(H2O)3(OH)3] + 3CO2 + 3H2O 1.7. Amphoteric Hydroxides
[Fe(H2O)6]2+ + CO32- → FeCO3 + 6H2O Amphoteric – showing both acidic and basic p
**Rule – Carbonates of M2+ exist, while M3+ carbonates do not Aluminium Hydroxide is example, reacts with
Al(H2O)3(OH)3 + 3HCl → A
Reacts with acid, therefore
The neutral metal hydroxide formed is essentially M(OH)3, it’s
Ammonia has the same effect as hydroxide ions in Al(H2O)3(OH)3 + OH- → [Al
uncharged, insoluble and forms a ppt
removing protons – both basic
The neutral metal hydroxide formed is essentially M(OH)2, it’s
uncharged, insoluble and forms a ppt
