Transition Metals Pt2 Flashcards
What happens when transition metals form ions?
They lose the 4s electrons before the 3d
This process influences their oxidation states and reactivity.
Why can transition metals easily oxidize and reduce?
Ions contain partially filled sub-shells of d electrons that can easily lose or gain electrons
This characteristic allows for a variety of oxidation states.
What is the trend in the relative stability of oxidation states across the period?
Relative stability of +2 state with respect to +3 state increases across the period
This indicates that the +2 oxidation state becomes more favorable as you move across the transition metals.
What do compounds with high oxidation states tend to be?
Oxidising agents
MnO is a compound with a high oxidation state.
Examples include V2+ and Fe2+ for reducing agents
How many main oxidation states does vanadium have?
Four main oxidation states
These include +5, +4, +3, and +2.
What is the color of the VO3- ion (oxidation state +5)?
Yellow solution
This ion is often found as ammonium vanadate (V).
What is the color of the VO2+ ion (oxidation state +4)?
Blue solution
This is one of the oxidation states of vanadium.
What color corresponds to the V3+ ion (oxidation state +3)?
Green solution
This indicates the presence of vanadium in the +3 state.
What is the color of the V2+ ion (oxidation state +2)?
Violet solution
This is the lowest oxidation state for vanadium.
What happens when acid is added to NH4VO3?
Turns into a yellow solution containing the VO2+ ion
This demonstrates the behavior of vanadium in acidic conditions.
What occurs when zinc is added to vanadium (V) in acidic solution?
Reduces the vanadium down through each successive oxidation state
The color changes from yellow (+5) to blue (+4) to green (+3) to violet.(+2)
What does the E value being more negative tell us?
Better reducing agent
Other one is better oxidising agent
Reaction when zinc in HCl added to a dichromate (VI) ion - Cr2O7 2-? What happens if Fe2+ is used instead of zinc?
Reduced Cr2O7 2- (orange) into Cr3+ (green) and Cr2+ (blue)
Fe2+ weaker reducing agent so only reduce dichromate to Cr3+
Conditions needed when zinc and dichromate react?
Keep zinc/dichromate under hydrogen atmosphere so can be reduced into Cr2+
As O2 in air will oxidise Cr2+ to Cr3+
- ACIDIC CONDITIONS
Cr2O7 2- + 14H+ + 3Zn —> 2Cr3+ + 7H2O + 3Zn2+
Cr2O7 2- + 14H+ + 4Zn —> 2Cr2+ + 7H2O + 4Zn2+
How can Cr2+ be stabilised?
Form a stable complex ion with ligand, ethanoate ion
- bubble Cr2+ ions through SODIUM ETHANOATE —> stable red precipitate of chromium(II) ethanoate forms
What is the reaction between Fe2+ and Cr2O7 2- an example of?
Quantitative redox titration - doesn’t need indicator
Cr2O7 2- + 14H+ + 6Fe2+ —> 2Cr3+ + 7H2O + 6Fe3+
ORANGE GREEN
In terms of electrode potentials, why does zinc reduce dichromate down to Cr2+ and iron only to Cr3+?
Electrode potential of iron(II) is between the 2 chromium half equations - so reduce down to Cr3+
Zinc has more negative electrode potential than all chromium half equations so zinc will reduce chromium down to Cr2+
How is chromium oxidised?
Acidified (not easy to oxidise) so add excess NaOH to form alkaline (easier to oxidise)
[Cr(H2O)6]3+ —-> [Cr(OH)6]3-
Oxidised by using ox agent - hydrogen peroxide
2 [Cr(OH)6]3- + 3H2O2 —> 2CrO4 2- + 2OH- + 8H2O
GREEN SOLUTION. YELLOW SOLUTION
How can chromate (CrO4 2-) and dichromate ions (Cr2O7 2-) be converted into each other ?
Equilibrium reaction
2CrO4 2- + 2H+ —> Cr2O7 2- + H2O
Yellow solution. Orange solution
NOT REDOX as both have ox number +6 so ACID BASE reaction
- adding acid/alkali can shift position of equilibrium left and right respectively
Iron (II) Metal aqua ions with limited OH- and NH3
[Fe(H2O)6]2+ (aq) + 2OH- (aq) →[Fe(H2O)4(OH)2](s) + 2H2O (l)
GREEN solution—> GREEN PRECIPITATE darkens on standing
[Fe(H2O)6]2+ (aq) + 2NH3 (aq) → [Fe(H2O)4(OH)2](s) + 2NH4+ (aq)
GREEN SOLUTION —> GREEN PRECIPITATE
Fe3+ Metal aqua ions with OH- /NH3?
[Fe(H2O)6]3+ (aq) + 3OH- (aq) → [Fe(H2O)3(OH)3] (s) + 3H2O (l)
YELLOW SOLUTION —> BROWN PRECIPITATE
[Fe(H2O)6]3+ (aq) + 3NH3 (aq) → [Fe(H2O)3(OH)3] (s) +3NH4+ (aq)
YELLOW solution —> BROWN PRECIPITATE
Copper (II) metal aqua ions with OH- and NH3?
[Cu(H2O)6]2+ (aq) + 2OH- (aq) →** [Cu(H2O)4(OH)2]**(s) + 2H2O (l)
[Cu(H2O)6]2+ (aq) + 2NH3 (aq)—> [Cu(OH)2(H2O)4](s) + 2NH4 + (aq)
BLUE SOLUTION —> pale BLUE PRECIPITATE
Excess AMMONIA, copper hydroxide:
Cu(OH)2(H2O)4 + 4NH3 (aq) —> [Cu(NH3)4(H2O)2]2+ (aq) + 2OH- (aq) + 2H2O (l)
** BLUE PRECIPITATE —> DEEP BLUE SOLUTION**
Ligand exchange
Cobalt (II) with OH- and NH3?
[Co(H2O)6]2+ (aq) + 2OH- (aq) →** [Co(H2O)4(OH)2]**(s) + 2H2O (l)
[Co(H2O)6]2+ (aq) + 2NH3 (aq) →[Co(H2O)4(OH)2](s) + 2NH4 + (l)
Pale pink solution —> blue precipitate (turns brown on standing)
Excess NH3 with cobalt hydroxide :
Co(H2O)4(OH)2 + 6NH3(aq) —> [Co(NH3)6]2+ + 2OH- (aq) + 4H2O (l)
Ligand exchange reaction : BLUE PRECIPITATE dissolved —> YELLOW SOLUTION
Chromium (III) with OH- and NH3?
[Cr(H2O)6]3+ (aq) + 3OH- (aq) —> Cr(OH)3(H2O)3 + 3H2O (l)
[Cr(H2O)6]3+ (aq) + 3NH3 (aq) —> Cr(OH)3(H2O)3 + 3NH4 + (aq)
Green solution —> grey/green precipitate
[Cr(H2O)3(OH)3] (s) + 3OH- (aq) → [Cr(OH)6]3- (aq) + 3H2O (l)
GREEN SOLUTION
When Excess NaOH added - further deprotonation
Cr hydroxide dissolves
What is ampoteric behavior? An example of this?
A metal hydroxide that can act as a base and acid is an amphoteric hydroxide
Cr(III) hydroxide acts as an acid, as it reacts with a base (excess NaOH)
Cr(III) hydroxide can also act as a base as it reacts with acids:
[Cr(H2O)3(OH)3] (s) + 3H+ (aq) → [Cr(H2O)6]3+ (aq)
Wat type of reaction is it when meal aqua ions react with limited OH- and NH3?
Deprotonation acid-base reaction
- removal of proton (2 (or3) hydroxide ions remove hydrogen ions from 2 (or 3) water ligands converting them to water molecules)
When excess NH3 added, which metals react and what type of reaction occurs?
Cu, Co,Cr
Ligand substitution - NH3 is Lewis base (donate electron pair)
Ligands NH3 and H2O imilar size/uncharged
- ligand exchange occurs without change in cooodination no. For Co and Cr
Cr reacting with excess NH3?
Cr(OH)3(H2O)3 (s) + 6NH3 (aq) —> [Cr(NH3)6]3+ (aq) + 3H2O (l) + 3OH- (aq)
Ligand exchange. PURPLE SOLUTION
What does the chlorine ligand react with and what changes does it cause?
Ligand substitution reaction with Cu, Co, Fe(III)
Cl- ligand larger than H2O and is charged so change in:
Coord no. From 6 to 4
Octahedral —> tetrahedral
Overall charge of complex
Cl- acts as Lewis base - donates electrons
Reaction of Cu, Co, Fe(III) with chloride ligand?
[Cu(H2O)6]2+ (aq) + 4Cl– (aq) → [CuCl4]2– (aq) + 6H2O (l)
BLUE —> YELLOW/GREEN SOL
ALL LIGAND EXCHANGE
[Co(H2O)6]2+ (aq) + 4Cl- (aq) —> [CoCl4]2– (aq) + 6H2O (l)
PINK —> BLUE
[Fe(H2O)6]3+ (aq) + 4Cl- (aq) —> [FeCl4]– (aq) + 6H2O (l)
YELLOW —-> ORANGE
What is chelate effect?
Replacement of monodentate ligands with bidentate/multidentate ligands in complex ions —> MORE STABLE COMPLEX FORMED
ΔS is positive because increase in no. Particles + reactions between aqueous ions have small enthalpy change , THEREFORE ΔG IS NEGATIVE
E.g EDTA chelates with aqueous cobalt (II)
[Co(H2O)6 ]2+ (aq) + EDTA4- (aq) → [CoEDTA]2- (aq) + 6H2O (l)
( 2 reactants become 7 products)
What happens if solid copper chloride (or nay other metal) is dissolved in water?
Aqueous [Cu(H2O)6]2+ complex forms NOT [CuCl4]2-
Why does the green precipitate from IRON (II) reacting with NaOH or NH3 turn brown on top after a few minutes ?
Fe2+ is oxidised by oxygen in air
