Transition Metals

• transition elements - d block elements that can form one or more stable ions with with a partially filled d-subshell.

  key tricks:

  • scandium is not a transition metal because its only stable oxidation state is +3. This gives it the electronic configuration of $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^{6\ }4s^0\ 3d^0$(empty d subshell)
  • zinc is not a transition metal because its only stable oxidation state is +2. This gives it the electronic configuration of $1s^2\ 2s^2\ 2s^6\ 3s^2\ 3p^6\ 4s^0\ 3d^{10}$
• transition metals have the following properties which originate from its incomplete d-subshell of electrons:

  • complex formation - via their dative bonds to ligands
  • formation of coloured ions

  
  

  • variable oxidation states - they have electrons of similar energy in both the 3d and 4s levels. Hence these elements can form ions of roughly the same stability by losing different numbers of d subshell electrons. When forming ions, they lose 4s electrons before 3d ones.

  
  

  • catalytic activity
  • higher mps, bps, and densities than normal metals
• ligand - molecule or ion that forms a co-ordinate/dative bond with a transition metal.
• monodentate ligands donate only one pair of electrons to a transition metal. The most common examples include:

  $H_2O,\ NH_3,\ OH^-,\ and\ Cl^-$

  $H_2O,\ OH^-,\ NH_3$ are all ligands of similar size, but $Cl^-$is larger. Hence fewer $Cl^-$ ligands can fit around a transition metal ion.
• bidentate ligands can donate 2 pairs of electrons to a transition metal:

  
• multidentate ligands donate many pairs of electrons to a transition metal. The most important one is $EDTA^{4-}$ which is hexadentate:

  
• complex - a central metal ion surrounded by ligands.
• lewis acid - electron pair acceptors.

  these are the transition metals in complex ions.
• lewis base - electron pair donors.

  these are the ligands in complex ions.
• coordination number - the number of co-ordinate bonds being made to the central metal atom/ion.
• haem is an iron (II) complex with a multidentate ligand. This makes haemogloblin which transports oxygen around the body:

  • oxygen forms a coordinate bond to iron (II) in haemoglobin.
  • carbon monoxide is toxic because it binds strongly to iron (II), preventing oxygen from binding and limiting oxygen supply around the body as a result.
• cisplatin - anticancer drug. It binds to adjacent guanine bases in the same strand of DNA within cancerous cells, preventing replication as the DNA structure is disrupted.

  
  

  However, it does not differentiate between healthy and cancerous cells so has adverse health effects (hair loss and a weakened immune system).

  
  

  In spite of this, its still used in anticancer treatments because the benefits outweigh adverse health impacts that could be experienced.
• transplatin is not used as an anticancer drug as it binds to guanine on seperate strands of DNA (easier to repair).
• shapes of ligands:

  octahedral:

  • coordinate number of 6
  • occurs with small monodentate ligands.

  
• shapes of ligands:

  tetrahedral:

  • coordinate number of 4
  • occurs with large monodentate ligands (Cl−)

  
• shapes of ligands:

  square planar:

  • coordinate number of 4
  • occurs with $d^8$ complexes such as $Ni^{2+}and\ Pt^{2+}$

  
• shapes of ligands:

  linear:

  • coordinate number of 2
  • occurs with $\left[Ag\left(NH_3\right)_2\right]^+$AKA tollen's reagent

  
• cis-trans isomerism is exhibited by some octahedral complexes (reference the water groups):

  
• optical isomerism is exhibited by octahedral complexes with 3 bidentate ligands:

  
• square planar complexes with different ligands can similarly exhibit cis-trans isomerism. Cis- and trans- platin are 2 good examples:

  
• transition metals appear coloured as they absorb certain wavelengths of light.

  For example, if a complex absorbs green light and transmits red and blue, the complex will appear purple (a combination of the transmitted colours):

  
• when a complex ion is formed, the d orbitals split in energy.

  wavelengths of light are absorbed, to excite an electron to the higher energy d orbital (from ground to excited state).

  the energy of the light absorbed corresponds to the energy gap between the ground and excited states — the larger the splitting gap, the greater the energy gap and hence the frequency of light absorbed:

  you see the complementary colour to what's absorbed

  ΔE = hv = hc/wavelength

  • ΔE - energy gap between orbitlas in J.
  • h - Planck's constant ($6.63\times10^{-34}$ Js)
  • v - frequency (hertz or $s^{-1}$)

  
• naming complex ions:

  $\left[Pt\left(NH_3\right)_4\right]^{2+}$= tetraammineplatinum (II)

  $\left[CrCl4^-\right]$ = tetrachlorochromate(III)

  
• factors that affect the colour of complexes:

  • the identity of the transition metal
  • the ligand identity and shape of the complex.
  • the oxidation number of the transition metal.

  the above all change the energy gap between orbitals, within the 3d subshell when split.
• Vanadium exists as 4 main oxidation states (+5, +4, +3, and +2)

  Addition of zinc in acidic solution will reduced Vanadium (V) through each succesive oxidation state — changing through each colour:

  $Zn(s)+$$2\ VO_2^+$$(aq)+$$4H^+$$(aq)\ \rightarrow\ 2\ VO^{2+}(aq)\ +$$2H_2O(l)\ +$$Zn^{2+}(aq)$

  $Zn(s)\ +$$2\ VO^{2+}(aq)\ +$$4\ H^+$$(aq)\ \rightarrow\ 2\ V^{3+}(aq)\ +$$2H_2O(l)\ +$$\ Zn^{2+}(aq)$

  $Zn(s)\ +$$2\ V^{3+}(aq)\ \rightarrow\ 2\ V^{2+}(aq)\ +$$Zn^{2+}(aq)$
• Vanadium colour wheel:

  
  

  You'd Better Get Vanadium

  
• Chromium exits as three main oxidation states: +6, +3, and +2.

  Addition of zinc in acidic solution will reduce dichromate (VI) to chromium (II):

  $3\ Zn(s)\ +$$\ Cr_2O_7^{2-}(aq)+$$14\ H^+$$(aq)\ \rightarrow\ 2\ Cr^{3+}(aq)+$$7\ H_2O(l)\ +$$\ 3\ Zn^{2+}(aq)$

  $Zn\ (s)\ +$$\ 2\ Cr^{3+}(aq)\ \rightarrow\ 2\ Cr^{2+}(aq)\ +$$\ Zn^{2+}(aq)$

  
• In addition (Or You'll go blind)

  chromium (III) can be oxidised to chromate (VI) using alkaline hydrogen peroxide. This is because in alkaline conditions chromium (III) exists as $\left[Cr\left(OH\right)_6\right]^{3-}$

  $\left[Cr\left(OH\right)_6\right]^{3-}(aq)\ +$$1\frac{1}{2}H_2O_2\ (aq)\ \rightarrow\ CrO_4^{2-}(aq)+$$\ 4\ H_2O(l)\ +$$\ OH^-(aq)$

  chromate (VI) can be converted to dichromate (VI) by adding an acid:

  $2\ CrO_4^{2-}(aq)\ +$$\ 2\ H^+$$(aq)\ \rightarrow\ Cr_2O_7^{2-}(aq)\ +$$\ H_2O(l)$

  
• $\left[Sc\left(H_2O\right)_6\right]_{(aq)}^{3+}$ is colourless because $Sc^{3+}$ has an empty d sub-shell, meaning that no d electron can be excited to a higher energy by absorbing wavelengths of visible light.
• $\left[Zn\left(H_2O\right)_6\right]_{(aq)}^{2+}$is colourless because $Zn\ ^{2+}$has an full d sub-shell, meaning that no d electron can be excited to a higher energy by absorbing wavelengths of visible light.
• ligand substitution reactions can occur with no change in co-ordination number, when similar sized ligands are exchanged.

  e.g. $\left[M\left(H_2O\right)_6\right]_{(aq)}^{2+}+$$6\ NH_3(aq)\ \rightarrow\ \left[M\left(NH_3\right)_6\right]_{(aq)}^{2+}+$$6\ H_2O(l)$

  They can also occur with a change in co-ordination number, when a larger ligand (like $Cl^-$) substitutes a smaller ligand. This also involves a change in the shape of the complex.

  e.g. $\left[M\left(H_2O\right)_6\right]_{(aq)}^{2+}+$$4\ Cl^-(aq)\ \rightarrow\ \left[MCl_4\right]_{(aq)}^{2-}+$$6\ H_2O(l)$
• Chelate Effect- the substitution of monodentate ligands with bidentate or multidentate ligands leads to a more stable complex:

  $\left[M\left(H_2O\right)_6\right]_{(aq)}^{2+}+$$EDTA_{(aq)}^{4-}\ \rightarrow\ \left[M\left(EDTA\right)\right]_{(aq)}^{2-}+$$6H_2O(l)$
• Chelate Effect explanation:

  • there is a positive value for the entropy change of the system, as there are more products than reactants.
  • there is a small, if not 0, value for the enthalpy change, because there are similar/the same number of bonds in the product and reactant complexes.

  this means that these substitution reactions are more likely to be spontaneous, as $ΔG\le0$

  
• when concentrated hydrochloric acid is added to $\left[Cu\left(H_2O\right)_6\right]_{(aq)}^{2+}$ and $\left[Co\left(H_2O\right)_6\right]_{(aq)}^{2+}$ a ligand substitution reaction takes place and tetrahedral complexes form:

  $\left[M\left(H_2O\right)_6\right]_{(aq)}^{2+}+$$4\ Cl^-(aq)\ \rightarrow\ \left[MCl_4\right]_{(aq)}^{2-}+$$6\ H_2O(l)$

  
• tollens reagent reduced silver (I) to silver metal when it oxidises aldehydes.
• Redox titration between $Fe^{2+}and\ MnO_4^-$ (part 1)

  Its self indicating as $MnO_4^-$ is purple and $Mn^{2+}$is colourless.

  $MnO_4^-(aq)\ +$$\ 8H^+$$(aq)+$$5\ Fe^{2+}(aq)\ \rightarrow\ Mn^{2+}(aq)\ +$$\ 4H_2O(l)+$$5\ Fe^{3+}(aq)$
• Redox titration between $Fe^{2+}and\ MnO_4^-$ (part 2)

  a sufficient amount of sulphuric acid (only option) must be used to provide enough $H^+$. If not a brown precipitate of $MnO_2$ would form:

  $MnO_4^-(aq)+$$4\ H^+$$(aq)\ +$$3e^-\rightarrow\ MnO_2(s)+$$2\ H_2O(l)$
• Redox titration between $Fe\ ^{2+}$ and $MnO_4^-$ (part 3)

  • $HCl$ can't be used as it would produce lots of poisonous $Cl_2$.
  • $HNO_3$ can't be used either as $NO_3^-$ ions are oxidising agents, and oxidise $Fe^{2+}$ to $Fe^{3+}$ leading to the production of less $Mn^{2+}$
• heterogenous catalyst examples

  • $Fe$in the Haber process — $N_2(g)+$$3\ H_2\ \rightarrow\ 2\ NH_3(g)$
  • $Pt\$in catalytic converters — remove$NO_x\ and\ CO$
• Heterogenous catalyst example:

  • $V_2O_5$ in the Contact process — conversion of $SO_2\ to\ SO_3$ to help make sulphuric acid:

  overall equation: $SO_2(g)+$$1\frac{1}{2}O_2(g)\ \rightarrow\ SO_3(g)$

  $SO_2(g)+$$V_2O_5(s)\ \rightarrow\ SO_3(g)+$$V_2O_4(s)$

  $V_2O_4(s)+$$\frac{1}{2}O_2(g)\ \rightarrow\ V_2O_5(s)$