# Ligands and complex ions

### Ligands and complex ions

#### Lessons

In this lesson, we will learn:

• The definition of complex ion, ligand and coordination number.
• How to name simple ligands and how they bond to metals to create complex ions.
• The conditions that create coloured transition metal complexes.
• How ligand size and denticity affects the coordination number of a complex ion.
• Some important examples of complex ions and mutidentate ligands.
Notes:

• In the previous lesson, Introduction to transition metals and the d-block, we saw that one of the properties of transition metals is their ability to form coloured compounds and have variable oxidation states. These properties don’t come from the metal ions alone; they occur when metal ions form compounds with other molecules bonding to them, forming what we call coordination complexes or complex ions.

• Complex ions are made of a central metal ion(s) surrounded by ligands bonding to the metal via dative bonding. In dative (coordinate) bonding both electrons are ‘donated’ by one species, which here is the ligand.
Since it involves two electrons coming from one species, many ligands in transition metal compounds are molecules with lone pairs that datively bond to positively charged metal centres. See the image below for an example of dative bonding:

• This is sometimes called dative covalent bonding, because the electrons are shared (covalent) despite both coming from one source.

• A ligand is any molecule that datively bonds to a metal centre using a lone pair. There are many different ligands – anything that can donate a lone pair could be a ligand:
• Neutral molecules with lone pairs, such as water, are very common ligands. When water acts as a ligand it is called ‘aqua’ in the compound’s name. Ammonia (NH3) is another common ligand, called ‘ammine’ in complex ions.
For example, the compound below is called iron (II) hexaaqua, for the six molecules datively bonding to the one metal centre.

• The charge on the metal and ligand will combine, just like in any other bonding. Since water is neutral, there isn’t a change in charge from what the iron was before: Fe3+.
If it’s more than one atom, the ligand is written in regular brackets with the entire compound in square brackets, the charge outside. This is shown in the diagram as [Fe(H2O)6]3+.

• Anions can also act as ligands by donating a lone pair. An example of this is the cobalt (II) chloride complex. See the image below:

• As chloride is a negatively charged species, their dative bonding to the Co2+ ion leads to an overall 4- charge for the complex, the result of Co2+ and 4Cl- bonding together. This is written [CoCl4]2-.

• Ligands bond to metal centres using EMPTY orbitals – they do not interact with the partially filled 3d orbitals already containing electrons!
This is because the ligands are contributing two electrons to make the bond. They cannot fit into an orbital that already contains an electron.
This is explained more below.

• When ligands bond to metal ions, they cause splitting of the metal’s d-orbitals which causes the colour of the compound to appear.
Before the ligands bond, the five d-orbitals are all equal energy (known as degenerate orbitals). But when the ligands interact with the empty higher energy orbitals, it causes a splitting in the partially filled d-subshell.
• For example, Fe2+ as an ion has the electron configuration [Ar] 4s0 3d6, with one of the orbitals holding a pair of electrons and the other four each with one. When the ion interacts with water and becomes [Fe(H2O)6]2+, the dative bonds from the water ligands cause the d-orbitals to split – now three are lower energy than before, and two are higher.
This gap between the lower and higher energy d-orbitals can be transitioned between, if the electrons absorb a frequency of light in the visible region.
If electrons absorb some visible light, the rest is reflected as colour.

• Different ligands have different strength in their interactions, so they produce different d-orbital splitting and therefore a different $\triangle E$ (gap in energy) between the lower and higher d-orbitals. The energy an electron must absorb to transition from the lower to the higher level is a particular frequency and wavelength.
Whatever light is not absorbed in the $\triangle E$ transition is reflected as the colour we see in the solution. In short, changing the ligand can change the colour of the metal ion.
See below for a comparison between [Fe(H2O)6]2+ and [Fe(CN)6]4-.

• This is why metals that are not transition metals (partially filled d-orbitals) have colourless salts and solutions.
• If the d-orbitals are completely empty (like in Sc), there are no electrons to make d-d transitions, and the ligands interact with those empty orbitals instead.
• If the d-orbitals are completely full (like in Zn), there is no space for d-d electron transitions to happen!

• The oxidation number and charge of the metal is also responsible for different coloured complex ions. The differing oxidation number will affect the strength of ligand interactions and lead to a different $\triangle E$.

• The number of ligands around a metal ion can affect the colour produced too. The number of ligands around a metal centre is called the coordination number.
• How many ligands can fit around a metal centre depends on their size – chloride ligands for example are larger than water ligands. As a result, a metal centre can normally only accommodate four chloride ligands instead of six water ligands. This then affects the overall geometry of the complex:
• Complex ions with a coordination number of 6 usually have an octahedral geometry. This is the geometry taken with hexaaqua and hexaammine complex ions like [Co(H2O)6]2+.
• Complex ions with a coordination number of 4 usually have a tetrahedral geometry. This is the geometry taken with four chloride ligands, for example [CoCl4]2-.
• Some complex ions, due to the metal ion, can form square planar complexes with a coordination number of four. This is the geometry of cis-platin, a platinum-based anticancer drug.
Importantly, the anticancer properties are only found in the cis-isomer, not the trans isomer where the two NH3 ligands are opposite each other.
See the diagram below:

• Another factor that affects coordination number is denticity – this is the number of dative bonds that the ligand makes to the metal centre. Most ligands are monodentate ligands which means only one dative bond is made, but there are many multidentate ligands. See the table below for some common multidentate ligands.

• As well as cis-platin, the haemoglobin molecule (the oxygen-carrying agent in our blood) is a coordination complex. It is an iron (II) based complex with a multidentate ligand that carries oxygen because O2 binds as a ligand to the Fe2+ ion. It is a relatively weak interaction. That’s good; it means the O2 ligand can leave and oxygen can get into our cells!
Unfortunately, ligands binding to haemoglobin is exactly why carbon monoxide (CO) is highly toxic to humans. CO is far stronger than O2 at binding to haemoglobin, so CO molecules will replace O2 ligands on the haemoglobin complex and prevent O2 from being transported in our body.
This type of reaction is called a ligand exchange which we’ll look at next lesson!
• Introduction
Transition metal complexes and ligands.
a)
Introduction to ligands and complex ions.

b)
How ligands bond to metals.

c)
How does colour form in transition metal compounds?

d)
Coordination number and ligand denticity.

e)
Important coordination complexes.

• 1.
Explain the difference between colour appearances in d-block compounds.
a)
Explain why [Cu(H2O)6]2+ and [Co(H2O)6]2+ are both coloured compounds, but Zn2+ does not form any coloured compounds.

• 2.
Understand How ligands affect coordination number due to denticity and size.
a)
State the coordination number of the compounds:
1. [Cu(H2O)6]2+
2. [Cu(en)3]2+, where 'en' is ethylenediamine NH2CH2CH2NH2
3. [CuCl4]2-

b)
Explain why [Cu(en)3]2+ has a lower coordination number than [Cu(H2O)6]2+

c)
Explain why [CuCl4]2- has a lower coordination number than [Cu(H2O)6]2+.