# Chemistry (A-Level)/Transition elements
## Transition Metals
### Defining a transition element

A transition metal is a d-block element that can form at least one stable ion with an incomplete d sub-shell. D-block elements are so-called because their sub-shell with the highest energy level is a d sub-shell. In A2-level chemistry, you only need to know of the first row of the d-block, elements Scandium(Sc) - Zinc(Zn). Of these ten metals, all but Scandium(Sc) and Zinc(Zn) are transition elements.

### Electronic configuration

The electronic configuration for Scandium(Sc) is: 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$4s$^2$3d$^1$. As you move across the first row of the d-block, electrons are added onto the 3d sub-shells, up to Zinc(Zn): 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$4s$^2$3d$^{10}$. The 4s sub-shells are generally written before the 3d sub-shells.This is because 4s sub-shells have a lower energy level than 3d sub-shells and a therefore filled before 3d sub-shells.
There are two exceptions to this in the top row of the d-block, Copper(Cu) and Chromium(Cr). The electronic configuration for Copper(Cu) is 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$4s$^1$3d$^5$, and for Chromium(Cr) it is 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$4s$^1$3d$^{10}$. These elements act like this as it is more energetically favored for the electrons to half-fill all of the orbitals, or fill all the d-orbitals leaving the s-orbital half-filled.
When forming ions, electrons are taken from the 4s sub-shell before the 3d sub-shell. e.g. Nickel(Ni): 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$4s$^2$3d$^8$ Nickel(II)(Ni$^{2+}$): 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$3d$^8$ Nickel(III)(Ni$^{3+}$): 1s$^2$2s$^2$2p$^6$3s$^2$3p$^6$3d$^7$

### Properties

All of the transition metals have the properties of metals. This includes: high melting and boiling points, high density, shiny surface, forms giant metallic lattices and coducts electricity. As well as these there are properties that only apply to transition metals. These are:
* Variable oxidation states
* Form coloured compounds
* Catalytic properties
* From complex ions
The transition elements you need to know (Ti to Cu) can all form two or more ions with oxidation numbers +2 and above.
When light passes though a solution, some of the wavelengths are absorbed. The colour you see is a result of wavelengths that were not absorbed by the ions. e.g. A Cu$^{2+}$ solution absorbed light at red/orange wavelengths, therefore we see it as blue. Generally, if a compound does not contain a transition metal ion (one with incompletely filled d-orbitals) it will be colourless.

### Catalytic Uses

Transition metals are good catalysts. There are two main ways in which they are used.
* They can provide a surface for a reactant to adsorb onto. The reactant can then be desorbed to take part in the reaction.
* They can change their oxidation state and then bond with the reactants to give a lower activation energy.
There are some well known uses of transition metals as catalyst that you should know.
* Iron (Fe) - used in the Haber process in the production of ammonia (NH$_3$)
* Vanadium Oxide (V$_2$O$_5$) - used in the contact process in the formation of sulphur trioxide(SO$_3$) for the production of sulpuric acid(H$_2$SO$_4$)
* Nickel (Ni) - used for hydrogenation of alkenes, breaking the double bond to form alkanes
* Manganese Oxide (MnO$_2$) - in the decomposition of hydrogen peroxide(H$_2$O$_2$) which is a simple way to produce oxygen gas(O$_2$)

## Complex Ions

A complex ion contains a central transition metal ion surrounded by (usually 4 or 6) ligands. A ligand is a molecule or ion which can donate a lone pair of electrons to form a coordinate bond with a central ion in a complex ion. A complex ion is given a coordination number which indicates the number of coordinate bonds to the central ion. 6 and 4 are common coordination numbers.
Some common ligands include:
* Water(H$_2$O)
* Ammonia(NH$_3$)
* Chloride(Cl$^-$)
* Hydroxide(OH$^-$)
* Thiocyanate(SCN$^-$)
* Cyanide(CN$^-$)
An example of a complex ion would be [Cu((H$_2$O)$_6$]$^{2+}$. We can work out the charge and oxidation number by looking at the overall charge and charge of the ligands. In this example the overall charge is 2+ and the six ligands are neutral. From that we can see that the central Cu ion must have a charge of 2+ and therefore an oxidation number of +2.
[FeCl$_4$]$^{2-}$ contains four chloride(Cl$^-$) ligands that add up to a total charge of 4-. Using the same technique, we can work out that the central Fe ion has a charge of 2+. [FeCl$_4$]$^{2-}$ only has four ligands because of the size of the chloride(Cl$^-$) ions. They are too big to fit six into the complex ion.

## Redox Titrations

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