23.4 Variable oxidation states of transition elements
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Why can transition metals have a greater variety of oxidation states in different compounds
. A typical transition metal can use its 3d electrons as well as its 4s electrons in bonding. This can be compared to eg group 2 which can only lose its outer 2 electrons to become a 2+ charge
Give the pattern for the oxidation states of transition metals
. All transition elements form +1 and +2 oxidation states Only the lower states of transition metals actually exist as simple ions Eg Mn2+ ions exist but not Mn7+ which is found in the (MnO4)- ion
Give the electron configuration of nickel and Ni2+ and explain reasoning for this
Ni: 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Ni2+ : 1s2, 2s2, 2p6, 3s2, 3p6, 3d8 So when the 2 electrons are lost, the 4s sub shell always empties before the 3d because once full it has more energy than the 3d shell so places itself above the 3d shell.
Give electron configuration of iron and Fe2+
Fe: 1s2 2s2 2p6 3s2 3p6 3d6 4s2 Fe2+: 1s2 2s2 2p6 3s2 3p6 3d6
Lots of the reactions in transition metal chemistry are redox reactions where the metals are either oxidised or reduced Give the example of irons two stable oxidation states: Which one is more stable Give reaction of Fe2+ with chlorine and its products
Fe2+ and Fe3+ . Fe3+ is more stable, eg Fe2+ is oxidised by oxygen in the air . It can also happen by chlorine: Fe2+(aq) + Cl2(g) —> 2Fe3+ + 2Cl- Here, chlorine is the oxidising agent as its oxidation number drops from 0 to -1 as is gains an electron, whilst Fe2+ loses an electron .
How does potassium manganate (VII) (KMnO4) act in acidic solution
It acts as an oxidising agent in solution that containins H+ ions . This is because it is oxidation state +7 so has lost all its valence electrons so will try and attract them back from other species, to reduce itself. So as a result will oxidise Fe2+ to Fe3+ During this reaction, the oxidation state of the manganese drops from +7 to +2
How would you write the equation for kMnO4 and Fe2+
.First do half equation for MnO4- to Mn2+ MnO4-(aq) + 8H+ +5e- –>Mn2+ + 4H2O .Then do half equation for Fe2+ to Fe3+ Fe2+ —> Fe3+ + e- 5Fe2+ –> 5Fe3+ + 5e- Then multiply the Fe equation by 5 to balance the number of electrons so they can cancel out Combine the equations by putting everything on the left side of arrow of both equations on the left and doing the same for the right. The charged electrons will cancel out. MnO4- + 8H+ + 5fe2+ –> Mn2+ + 4H2O + 5Fe3+
What are redox titrations
If you wish to measure the concentration of an oxidation or reducing agent , it is similar to an acid base titration
What colour is Fe2+ ion What colour is Fe3+ ion
Fe2+ is pale green Fe3+ is pale violet
What colour is MnO4- ion What colour is Mn2+ ion
Intense purple Pale pink
How would you do the redox titration between FeSO4 and KMnO4 What are the colour changes
Using a burette, add the KMnO4 (aq) to a solution containing Fe2+ ions which is already acidified with sulphuric acid The purple colour disappears as the MnO4- ions are converted to Mn2+ which is pale pink. Once enough MnO4- ions have been added to react with Fe2+ ions, one more drop of KMnO4- will turn the solution purple This is the end point of the titration
Why can’t you use HCl instead of sulphuric acid in the KMnO4 and FeSO4 reaction
. HCl contains Cl- ions which are oxidised by MnO4- ions This would affect the titration because the manganate ions should only be used to oxidise Fe2+ ions.
How can you use potassium dichromate (VI) for titrations K2(Cr2O7) Do the equations for it with iron
. It can also titrate against Fe2+ ions as it contains (Cr2O7)2- ions All aqueous (Cr2O7)2- + 14H+ 6e- —> 2Cr3+ + 7H2O Fe2+ —> Fe3+ + e- multiply by 6 6fe2+ —> 6Fe3+ + 6e- Combine: (Cr2O7)2- + 14H+ + 6Fe2+ —> 2Cr3+ + 7H2O + 6Fe3+
What ions do potassium dichromate turn into when reduced by the iron ions
Cr3+ ions
Why is sulphuric acid added to these titration reactions
To provide the H+ ions required
What indicator is used for the potassium dichromate and iron titration
sodium diphenylaminesulfonate Turns from colourless to purple at end point
How does the PH of a solution affect which oxidation state of a metal is reduced
A metal with a high oxidation state is reduced in acidic solution. So gains electrons A metal with a low oxidation state of a transition metal occurs in more alkaline solutions
Why do transition metals with low oxidation states get oxidised in alkaline solution more Why do transition metals with high oxidation states get reduced in acidic solution
. There is a tendency to form negative ions in alkaline solution Because oxidation is the loss of electrons, so this is easier from negatively charged species than for positively charged or neutral ones . Because they are electron deficient so it is more easy for them to gain electrons
What would the typical transition metal species be in acid solution M represents transition metals
M(H2O)6 2+ as it is positively charged
What would the typical transition metal species be in neutral solution M represents transition metals
M(H2O)4(OH)2 as it is neutral
What would the typical transition metal species be in alkaline solution M represents transition metals
M(H2O)2(OH)4 2- as it is negative
How would you oxidise a transition metal to a high oxidation state
. Add an alkali Then add an oxidising agent