Transition metals form ions with an incomplete d subshell. They become complex ions when they form ionic compounds with other atoms or molecules. (Find out more in Transition Metals and Substitution Reaction.) One of the most fascinating characteristics of complex ions is that they can show up in a rainbow of beautiful colours!
For example, you may be familiar with the orange-brown colour of iron oxide, otherwise known as rust. Or, perhaps, you have seen the shiny bronze colour of copper roofs slowly turn blue-green as time passes. Colour is not only a cool characteristic of complex ions, it is also helpful when we want to identify a transition metal.
What gives each complex ion its unique colour, and why do some ions change colour? Let’s find out!
- We will discuss the origin of the colours of complex ions.
- We will discuss the factors that affect the colours of complex ions.
- We'll learn about the various colours of different complex ions, and why ions change colour.
- Finally, we will discuss visible light spectroscopy, and how to determine the concentration of coloured ions using colorimetry.
Did you know the Statue of Liberty in the USA is made of copper? It was once a dull bronze colour! After 1900 it slowly began to change to the pale blue-green colour we see today. Copper naturally changes colour when exposed to the air for a long period of time. Through a process of oxidation, the outermost layer of the copper turns into what is known as basic copper carbonate. When close to the sea it becomes basic copper chloride.
Some artists and designers use artificially aged copper in their work because they find the blue-green colour of oxidised copper pleasing. We describe naturally and artificially aged copper as a patina or verdigris.
Fig. 1 - Oxidation has changed the colour of the copper surface of the Statue of Liberty to blue-green. Pexels.com
Factors affecting the colour of complex ions
To explain why transition metal complexes have different colours, you'll need to be familiar with the concept of the electromagnetic spectrum. Perhaps you’ve heard of the idea that visible light, or ‘white light’, is made up of a mixture of colours. You can see this if you shine white light through a prism - it splits into all of the colours of the rainbow. The light we can see is just a small part of the electromagnetic spectrum, which includes other kinds of light we cannot see like infra-red, and ultraviolet.
Fig. 2 - The visible spectrum
Each of the light colours in the visible spectrum has a certain wavelength. At one end we have red light, which has a wavelength of 700 nanometers (nm). At the other end we have violet light, with a wavelength of 400 nm. All the other colour waves exist between red and violet and have specific wavelength values between 700 and 400 nm, e.g. 438.4, or 605.0. As you can see, there's an infinite range of colours in the visible spectrum!
When white light passes through a transition metal complex in an aqueous solution, it absorbs some of the wavelengths. The remaining light waves are reflected or transmitted. These remaining light waves give the solution its characteristic colour. For example, copper(II) sulphate solution appears cyan (a pale blue) because, as light passes through it, it absorbs the light in the red region of the visible spectrum. In other words, the light that comes out the other side of the solution has all the other light wavelengths except for red. Mixing these wavelengths together forms the colour cyan.
Fig. 3 - Complex ions in solution absorb certain wavelengths of light and transmit others
You may be wondering what happened to the other colours not absorbed by the copper(II) sulphate solution. After all, an infinite number of colours exist in the visible spectrum. If the ion complex only absorbed red light, why do we see it as cyan blue, and not yellow or pink? We need to use a colour wheel like the one in the image below to find the answer.
Fig. 4 - A complex ion in solution transmits the complementary colour of the lightwave it absorbs
We call colours that are opposite each other on a colour wheel complementary colours. What colour can you see opposite red on the colour wheel? If a solution of a transition metal complex absorbs a colour, the light that passes through it will contain proportionately more of its complementary colour. When copper(II) sulphate absorbs red light, the light that gets transmitted will have more of its complementary colour, i.e., blue. That is why copper(II) sulphate solution appears blue.
Of course, this is a rather simple explanation for the colours of complex ions. We can explain the colours of aqueous solutions of complex ions even further by looking at the arrangement of the electrons in their 3d energy level.
Degenerate and non-degenerate d-orbitals
In the sub-section below, we shall be exploring electron promotion. In order to do this, however, we need to explore what degenerate and non-degenerate d-orbitals are.
Degenerate: d-orbitals that have the same energy levels
Non-degenerate: d-orbitals that do not have the same energy levels
In a transition element metal or ion there are 5 d-orbitals. Without any ligands bonded, all 5 are at the same level, therefore, they can be described as degenerate. However, after ligands bind to the transition metal ion, it is no longer by itself. When ligands bind, they form dative covalent bonds and this splits the 5 d-orbitals into two sets. These two sets do not have equal energy and are therefore described as non-degenerate. Now let's explore how the degenerate 5 d-orbitals split in an octahedral complex and a tetrahedral complex.
Splitting in an octahedral complex:
So firstly what is an octahedral complex? This is a complex that has 6 ligands covalently bonded to the central transition metal ion. We will be using the diagram on the right to explore this splitting. [Insert image 'Splitting of 3d orbitals in octahedral complex, StudySmart Original' here]We can first identify that the degenerate 5 d-orbitals have split into two levels with dx2-y2 and dz2 at a higher level and dyz, dxz, dxy at a lower level.Now you may be thinking, why does this happen? This is because the lone pairs of electrons in the dx2-y2 and dz2 orbitals repel more than in the dyz, dxz and dxy. So now you may be thinking, but why do they repel more? This is because when the ligands bind to the central metal ion it does this along the x, z and y axis and dx2-y2 and dz2 have lobes on these axes, So the electrons are closer together, leading to more repulsion which leads it to have a higher energy level than dyz, dxz and dxy. The energy difference between the two is ΔE.
Splitting in a tetrahedral complex:
A tetrahedral complex has 4 ligands covalently bonded to the central transition metal ion. We will be using the diagram on the right to explore this splitting. [Insert image 'Splitting of 3d orbitals a tetrahedral complex, StudySmart Original' here]
In a tetrahedral complex the ligands line up with dyz, dxz and dxy, meaning that there is more repulsion than at dx2-y2 and dz2. Therefore, dyz, dxz and dxy will be at a higher energy level, while dx2-y2 and dz2, will be at a lower and more stable energy level compared to the former. As in the splitting in octahedral complexes, the energy difference between the two is ΔE.
Electron promotion
A transition metal is one that forms stable ions with a partially filled d subshell. How do we know the colours of complex ions in solution are connected to their partially filled d subshell? Well, when we observe the colour of non-transition metal ions in solution, they are colourless. For example, scandium and zinc, although in the d-block, do not form ions with a partially filled d orbital. Their ions appear clear in solution. On the other hand, the transition metal ions with an incomplete 3d energy level form coloured solutions.
Just how does an incomplete 3d energy level give complex ions in solution colour? Let's take copper(II) sulphate and zinc sulphate as an example. As we've discussed, copper(II) sulphate solution is cyan blue, while zinc sulphate solution is colourless. They both contain sulphate, so any difference in colour must be caused by the copper and zinc ions. Cu2+ has the electron configuration 1s2, 2s2, 2p6, 3s2, 3p6, 3d9 whilst Zn2+ has the electron configuration 1s2, 2s2, 2p6, 3s2, 3p6, 3d10. You'll notice that Cu2+ has a partially-filled 3d subshell whereas Zn2+ has a complete 3d subshell.
In an aqueous solution, water ligands attach to the copper ion and split its 3d orbitals into two sets of non-degenerate orbitals. This means the two sets have slightly different energies, as explained above. The lower energy level has six electrons while the higher one has three. We call this splitting. Splitting only occurs if the 3d subshell is partially filled.
Fig. 5 - Ligands split the 3d orbitals in Cu2+ into two sets of non-degenerate orbitals
When the copper(II) sulphate solution absorbs light from the visible spectrum, an electron from the lower energy level gets promoted or excited by moving to the higher energy level. We say it moves from the ground state to an excited state. The energy absorbed by the excited electron depends on the difference in energy between the two levels.
Fig. 6 - An excited electron gets promoted when it absorbs a specific light frequency
A small energy difference between the split 3d energy levels of the Cu2+ ion means that the excited electron absorbs a small light frequency. Light frequencies are inversely proportional to their wavelengths, so large light frequencies have small wavelengths, and vice versa. In other words, red light has a large wavelength and a small frequency but violet light has a small wavelength and a large frequency. In our example, the promoted electron in the Cu2+ ion absorbs a small light frequency. This means that it absorbs light with a large wavelength, namely light from the red end of the visible spectrum. That is why we see copper(II) sulphate solution as cyan blue, the complementary colour of red.
On the other hand, the zinc ion has a completely filled 3d subshell so no splitting takes place. This means the electrons don't get excited. There aren't transitions that absorb light energy, so zinc complexes are colourless.
Fig. 7 - The Zn2+ ion has a completely filled 3d sub-shell, so no splitting takes place between the 3d orbitals
How to find ∆E of an excited electron
As we've discovered, when an electron moves from the ground state to an excited state it absorbs energy in the form of light. Light travels in waves and particles we call photons. Electrons absorb light energy in the form of photons. Albert Einstein showed us that the amount of energy in a photon equals the frequency of a lightwave, v, multiplied by Planck’s constant (6.626 x 10-34 Js), h. In other words, Ephoton = hv.
We express the change in energy when a electron goes from the ground state to an excited state as ∆E. You can find the wavelength of light absorbed by the excited electron by using the equation:
Where:
∆E is the exact amount of energy absorbed in joules, J
h is Planck’s constant, 6.626 x 10-34 Js
v is the frequency of light in hertz, Hz, or s-1
c is speed of light, 3.0 x 108 ms-1
𝝀 is the wavelength of light in meters, m
Why do these ions change colour?
You have seen that the amount of light energy absorbed (∆E) by electrons in the split d orbital gives ion complexes their unique colours. Any factor that changes the level of splitting in the d orbitals changes ∆E and so changes the colour of the ion complex. These include:
- Oxidation state
- Ligands
- Coordination number
Let's discuss how these factors change the ion colours.
Oxidation state
Oxidation states are numbers we give to ions that show how many electrons those ions has lost or gained, compared to the element in its uncombined state.
The oxidation state or oxidation number of an ion is the same as its ionic charge. An ion with a 2+ charge has an oxidation state of +2 and an ion with a 3- charge has an oxidation state of -3. Transition metals have a unique quality that allows them to have varying oxidation states. As oxidation state increases, so does splitting in the d orbitals and vice versa. Increased splitting simply means there is a larger energy gap between the split orbitals. This means that when there is a change in oxidation state, the colour of the ion complex also changes.
Ligands
Atoms or molecules bonded to a complex ion are called ligands. The identity of the ligands bonded to a complex ion can also change its colour. How so? Well, there are five 3d orbitals that all have the same energy. You can say they are degenerate. When ligands attach to a transition metal ion by dative covalent bonds, the 3d orbitals split into two sets of non-degenerate orbitals. That is to say, they no longer have the same energy - some have a higher energy than the others.
Electrical fields surrounding the ligands influence the energy gap between the split 3d orbitals. The stronger the electrical field, the larger the energy gap. The difference in energy between the split 3d orbitals determines the size of the light wavelength an ion complex can absorb. For example, ammonia causes more splitting than water in copper(II) sulphate. Water ligands give copper(II) sulphate a light blue colour, while ammonia ligands give it a deep navy colour.
Fig. 8 - Copper(II) sulphate in water and in ammonia
Coordination number
The coordination number is the number of ligands attached to the central ion. It is usually a number between two and nine.
A change in coordination number changes the colour of an ion by changing the amount of splitting in the 3d orbitals. For example, an octahedral ion has more splitting in the 3d orbitals than a tetrahedral ion. This change in coordination number changes the colour of the ion.
You can't just look at the change in coordination number when you want to find out why the colour of an ion complex has changed. This is because changing the coordination number generally involves changing the ligands attached to a complex ion, which changes the colour anyway!
To illustrate this, have a look at the reaction between copper(II) sulphate solution and concentrated hydrochloric acid shown below. When we add the acid slowly and continuously to the ion complex, the colour changes gradually from blue to green to yellow.
From the equation, can you see if the ligands have changed? What about copper's coordination number?
The 6 water ligands have changed to 4 chloride ion ligands.
The coordination number has changed from 6 to 4.
Copper's oxidation state is still the same even though the overall charge on the ionic molecule has changed from 2+ to 2-. This is because the copper ion itself still has a charge of 2+.
Remember that these factors all affect the splitting of the 3d orbitals, which changes the colour of the complex ion.
Visible light spectroscopy
Spectroscopy is the name of a branch of chemistry that studies the absorption and emission of light and other kinds of radiation. Different substances absorb different wavelengths of light. You have seen this in the case of copper(II) sulphate, which absorbs red light from the visible spectrum and transmits blue light. So copper(II) sulphate appears blue. Even colourless substances absorb specific wavelengths, but in the ultraviolet region, so our eyes don't notice. In visible light spectroscopy, we can use this knowledge to identify substances by means of a colorimeter.
To clarify, when we speak of spectroscopy, we mean the absorption and emission of all kinds of radiation and light from the electromagnetic spectrum. This includes wavelengths of light we cannot see with our eyes. Colorimetry is spectroscopy that involves light from the visible spectrum only.
A simple colorimeter measures light absorbance, or the amount of visible light that a substance absorbs. The amount of light absorption depends on the concentration of the substance in solution. This means a colorimeter can help to determine the concentration of coloured ions in solution.
In visible light spectroscopy, you pass different frequencies of light through a coloured filter and a sample of complex ions in solution. A detector and recorder is placed at one end to record the frequency of the transmitted light. The filter you use must match the part of the spectrum that the coloured complex absorbs the most. For example, a solution that absorbs red light appears blue, so you must use a red filter. This way, only red light passes through the solution, and maximum absorption can take place.
Fig. 9 - A colorimetry setup
You will need to produce a calibration graph or calibration curve to figure out the concentration of the sample solution. Let us use the example below to find the concentration of a sample of [Cu(H2O)6]2+ ions in solution.
First, you need to set up your colorimeter with a yellow filter, since [Cu(H2O)6]2+ ions absorb yellow light.
You will need to use different filters with different complex ions since they absorb different wavelengths of light.
Second, you measure the absorbance of sample solutions of [Cu(H2O)6]2+ that you already know the concentrations of. Here's an example of results you might get:
| Concentration (mol dm-3) | Relative Absorbance |
| 0.2 | 0.050 |
| 0.3 | 0.055 |
| 0.4 | 0.100 |
| 0.6 | 0.150 |
| 1.0 | 0.250 |
| 1.6 | 0.400 |
Plot these values on a calibration graph and draw a line of best fit, as shown below. The concentration must go on the x-axis and relative absorbance must go on the y-axis.
Fig. 10 - A calibration graph
Measure the absorbance of the solution you don't know the concentration of, then read its concentration from the graph you drew.
In this example, the absorbance measured from the sample of unknown concentration was 0.225. So the concentration of the sample of [Cu(H2O)6]2+ ions is around 0.9 mol dm-3.
- This graph is a straight line because at low concentrations, relative absorbance is directly proportional to the amount of coloured ions in solution.
- Spectrometry is handy because you can quickly measure the concentration of lots of samples of coloured ions, even very low concentrations. Also, none of the samples get used up or interfere with other reactions.
- The more concentrated a solution is, the more light it will absorb.
A colorimeter may not be sensitive enough to accurately record absorption in solutions that are very dark or very pale. In the case of pale solutions, a ligand exchange may help change the colour of the solution to a stronger one.
You may be curious about the colours of common ion complexes that you might encounter in everyday life. Earlier you learned about the colour of oxidised copper on the roofs of buildings. To round this lesson off, let us consider the colour of the ferric ion, manganese ion, and chromium ion.
The colour of the ferric ion
Ferric is the word chemists use to describe iron in the +3 oxidation state, Fe(III). In an aqueous solution, the ferric ion is yellow. On the other hand, the ferrous ion, which is the word we use to describe the Fe(II) ion, is green in solution. Many living organisms use iron's ability to switch between these two oxidation states in order to carry out vital processes. For example, iron in haemoglobin changes oxidation states when it binds to oxygen in the blood.
Another example is ferritin. If we ever have too much iron in the blood, we store it as the ferric ion in a protein called ferritin. When our cells need extra ions, ferritin releases ferric ions in a controlled way. These ferric ions are then reduced to ferrous ions before entering our cells. That's a lot of ferr- words!
The colour of the permanganate ion
Permanganate is the name we give compounds that have manganese in the +7 oxidation state. The permanganate ion (MnO4-) is a vibrant purple colour. Be careful not to confuse permanganate with manganate, which is the word we use for manganese in the +6 oxidation state. Manganate appears green in solution.
The colour of chromium ions
The chromium(III) ion (also known as hexaaquachromium(III)) is green in solution. The name chromium comes from the Greek word chroma which means 'colour'. This is because complexes that contain the chromium ion at higher oxidation states show us a wide range of colours. For example, in the past we used chromium(VI) in pigments called chrome yellow and chrome red. However, we have since discovered that chromium(VI) compounds are toxic and may cause cancer. Chromium also causes the colour in precious stones like the red in rubies, the pink in sapphires and the green in emeralds. So you see, the fascinating colours of complex ions make our world a more colourful and vibrant place!
Fig. 11 - Colours of common complex ions
Ion Colours - Key takeaways
- Complex ions in solution absorb specific wavelengths of light from the visible spectrum and transmit other wavelengths.
- Transition metals complexes reflect or transmit the complementary colour of the wavelength they absorb.
- Complex ions can absorb light energy from the visible spectrum because of splitting in the incomplete 3d subshell.
- The 3d orbitals in complex ions split into two non-degenerate levels.
- When an electron in the lower non-degenerate level absorbs light energy, it becomes excited and gets promoted to the higher non-degenerate level.
- The light wave absorbed by an electron in the split 3d subshell must be equal in energy to the energy difference between the two non-degenerate orbitals.
- You can find the wavelength of light absorbed by the excited electron by using the equation
- Factors that affect the colour of complex ions include oxidation state, coordination number, and ligands present.
- We can use visible light spectroscopy or colorimetry to find the concentration of coloured ions in solution.
- A calibration graph or calibration curve will help to determine the concentration of ions in solution.
- At low concentrations, relative absorbance is directly proportional to the concentration of coloured ions.
- Spectrometry is handy because it allows you to quickly measure the concentration of lots of samples of coloured ions, even those with very low concentrations.
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