Advanced Level Organic Chemistry: 15.5 UV and Visible Spectroscopy

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Doc Brown's Advanced Chemistry

PART 15.5 UV and Visible Spectroscopy

15.5.2 Examples of uv and visible absorption and reflectance spectra

Doc Brown's Chemistry Advanced Level Pre-University Chemistry Revision Study Notes for UK IB KS5 A/AS GCE advanced A level organic chemistry students US K12 grade 11 grade 12 organic chemistry courses Spectroscopic methods of analysis and molecular structure determination

All my advanced A level organic chemistry notes


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15.5.1 The origin of colour, the wavelengths of visible light, our perception!

15.5.2 uv-visible spectroscopy theory, spectrometer, absorption & reflectance spectra explained (this page)

15.5.3 uv-visible absorption spectra - index of examples: uses, applications, more on the chemistry of colour

15.5.2 uv-visible spectroscopy theory, spectrometer, absorption & reflectance spectra explained

Sub-index for this page

(a) The molecular absorption of uv and visible light photons - formation of electronic spectra

(b) How a uv-visible light spectrometer works

(c) More on absorption spectra and their interpretation

(d) Reflectance spectra

(e) Examples and theory of coloured inorganic compounds

(f) Examples and theory of coloured organic compounds - chromophores

(g) The theory of flame emission and absorption spectroscopy (separate page)

(a) The molecular absorption of uv and visible light photons - formation of electronic spectra

This section is all about how our perception of 'colours' is related to molecules when they absorb uv or visible light photon.

When a molecule absorbs a uv-visible photon, if it is of sufficient energy, an electron can be promoted to a higher electronic quantum level e.g. electronically, the molecule's electron is raised from a ground state to an excited state (often denoted by a *) and the molecule M is then described as being 'excited'.

M  ===  h  ===>  M*

This is shown by ∆E (= h) on the diagram below, the energy change from the ground state to the excited state.

An electron is promoted to a higher energy level molecular orbital (no details required pre-university!).

For absorption spectra, it is the outer bonding and non-bonding electrons (lone pairs) that are involved in the electronic quantum level changes.

The 'inner' electrons are held too strongly by the positive nucleus to be promoted to a higher level by uv-visible photons.

diagram expaining electronic energy level changes when a molecule absorbs a uv or visible light photon complications from vibrational rotational translational kinetic energy quantum levels

A very simple diagram illustrating two possible electron level changes i.e. a lower energy excitation from a visible light photon and from a ground state to an excited state, and higher excitation energy requirement from a uv photon.

An arrow pointing up means absorption of a photon of energy.

An arrow pointing down means releasing the energy absorbed as re-emitted radiation or more complex changes resulting ultimately increasing the translation kinetic energy of the material i.e. an increase in temperature.

Unfortunately things are a bit more complicated than the 'simple' == E ==> change indicated on the diagram.

In terms of absorption of EM radiation: Euv > Evisible >> ∆Evib  >>  Eir  >>  Etrans

i.e. higher electronic energy levels are only accessed by higher energy ultraviolet photons and the substance is colourless unless E is low enough to cause absorption of visible light photons to produce a coloured material.

I've indicated the associated vibrational and rotational levels for the uv ∆E, it is the same for the smaller ∆E visible light photon absorption (but not shown).  On some absorption spectra can actually see some of the vibrational modes as finer peaks on a broad absorption band. So, this situation needs some extra explanation!

The energy of a molecule in a particular state is the sum of the electronic energy + vibrational energy + rotational energy (and strictly speaking + translational kinetic energy too).

For each electronic quantum level, there are associated vibrational quantum levels of the bonds and for each vibrational level there are associated rotational quantum levels (for each of these there are associated translation kinetic energy levels).

So for a given electronic energy level change there are actually lots of other ΔE changes possible.

The gaps between vibrational levels are around 100 times smaller than electronic energy levels and gaps between rotational levels are about 100 times smaller than vibrational level.

This is why UV-visible absorption spectra tend to show as one or more very broad absorption bands.

What does an absorption spectra look like?

The three spectra of aromatic compounds and a transition metal ion below illustrate some of the above points.

Note in all cases the broad absorption bands, you rarely see sharp peeks you see in infrared spectra or NMR spectroscopy.

Absorption spectra are usually presented on the y-axis as the intensity of absorption (in various units) versus the wavelength (x-axis), λ usually in nm, but can be frequency too.

uv-visible absorption spectrum of methylbenzene example of aromatic spectra of arenes

(1) (c) doc bMethylbenzene, a colourless liquid, whose absorption spectrum is shown above.

Methylbenzene is colourless because the ∆E values are too high for visible light photon to effect excitation of an electron to a higher energy level.

However, below 380 nm (start of uv region), you do get absorption of uv photons from 275 nm downwards - several broad peaks and note the fine structure peaks from associated vibration levels in the methylbenzene molecule.

There is a λmax of 190 nm, these λmax maximum absorbance wavelengths are characteristic of each substance under investigation.

For more see The uv-visible absorption spectra of selected aromatic compounds - arenes


uv-visible absorption spectrum of 3-nitrophenol example of aromatic spectra of phenols

(2) 3-nitrophenol, pale yellow solid, same colour in a solution, whose absorption spectrum is shown above.

As well as in the uv region, 3-nitrophenol absorbs in the blue-violet light region (~380 - 500 nm) of the visible spectrum - and is yellow in colour, the complimentary of blue, see the simplified colour wheel below.

Two of bands have peaks, λmax at 275 and a λmax at 340 nm, but, although both peaks are in the 'invisible' uv region, the right-hand absorption band extends well into the violet-blue region of the visible spectrum.

Note that a different substituent group in the benzene ring alters the energy levels of the pi electrons, lowering them sufficiently to allow excitation with lower energy uv photons.

For more see The uv-visible absorption spectra of some phenols

simplified colour wheel of complimentary colours made up from primary colours secondary colours


uv-visible absorption spectrum of the hexaaquatitanium(III) transition metal complex ion

(3) A typical transition metal complex ion absorption spectrum - for the pale purple hexaaquatitanium(III) ion shown above (sometimes quoted as violet).

The λmax ~520 nm, shows strong absorption in the blue-green-yellow region, but little absorption in the violet and red regions, resulting in the ion being purple in aqueous solution from a combination of violet and red.

Note the strong absorption in the 'invisible' ultraviolet region <380 nm).

Colour theory of transition metal ions (this page section (e))

More examples transition metal ions (index of examples)


Further complications that broaden the absorption bands in uv or visible absorption spectra

In a solid, the main factors that cause broadening of the spectral line into a broad absorption band are the distributions of vibrational and rotational energies of the molecules in the sample (and also those of their excited states) which are 'super-imposed' onto the electronic level changes - so there are many possible quantum level changes.

In pure liquids or solutions, the same situation occurs, but there is now an extra broadening of the absorption band due to collisions between the molecules - which is significant enough to cause even more blurring together of the energy differences between the different rotational and vibrational states, such that the spectrum consists of broad absorption bands instead of discrete lines.

This means when we refer to a colour, we are dealing with a band of frequencies, but wavelengths are usually quoted on the y axis of uv-visible absorption spectra.


What happens to the energy of absorbed uv-visible photons?

Molecules in an excited state are inherently in an unstable condition and the electron returns to a lower level, and, in doing so, the absorbed radiation is re-emitted, but not at once - the radiation can be re-emitted in stages via intermediate energy levels.

The relatively smaller amounts of higher vibrational level energies are readily lost by molecular collisions, so the molecule can drop to lower vibrational levels for the same electronic level., this increases the translational kinetic of the molecules, slightly increasing the temperature of the material.

(b) How a uv-visible spectrometer works

how an uv-visible light absorption spectrometer works doc brown's advanced level chemistry notes

A uv-visible absorption spectrometer works in a similar way to an infrared spectrometer.

The substance under investigation is usually dissolved in a suitable solvent.

After passing through a diffraction grating that scans through the uv-visible wavelength range, the beam of uv-visible light is split into two identical beams, one passes through a reference cell of the pure solvent and the other through the solution of the material under investigation (the reference beam and sample beam).

There must be no absorption by the solvent in the wavelength region being investigated for a particular material.

This is not usually a problem with a coloured substance, because there are many colourless solvents that, in terms of electronic quantum level changes, only begin to absorb in the ultraviolet region of the spectrum <380 nm).

The solvent and solution are contained in cuvettes, usually of cross-section 1cm2, and made of high quality 'optical' standard pure silica glass to minimise interfering absorbances.

The detectors are synchronised with the diffraction grating to scan through the uv-visible bands and measure the absorbance versus the wavelength of the light beam.

A = log10(Io / I0 ) = ɛcL

A = absorbance, Io = Intensity of reference beam, I = intensity of sample beam, c = concentration, l = path length

ɛ = extinction coefficient, the absorption coefficient for a particular compound at a particular wavelength

This is mathematical expression of the Beer-Lambert Law which infers that the absorbance is proportional to concentration - often valid for relatively low concentrations.

A colorimeter is simple type of of visible light spectrometer (spectrophotometer) and measures intensity of absorption over a narrow range of wavelengths governed by a filter - but it is a simple and accurate method to measure concentration.

See Colorimetric analysis and determining a transition metal complex ion formula

AND note how a colorimeter works and the use of a linear calibration graph in colorimetric analysis.

(c) More on absorption spectra and their interpretation

First study the Three spectra were looked at in section (c) (opened in new window for convenience)

Interpreting a uv-visible spectrum (with particular reference to colour chemistry)

There are three main features a spectrum

(i) The wavelengths of the uv-visible wavelengths absorbed (or reflected), λ usually shown in nanometres (nm).

The wavelength of absorption depends on the values of the electronic quantum levels.

(ii) The intensity of the absorption, particularly the characteristic λmax peaks in the uv-visible spectrum.

The  λmax peaks are important for quantitative analysis, giving the maximum sensitivity possible.

(iii) The shape of absorption bands across the region where absorption takes place, see point (iii) below.

All three characteristics are dependent on the structure of the organic/inorganic molecule or ion.


More on aspects of molecular structure and absorption wavelength and intensity

(i) For organic molecules with delocalised electron systems, the longer the conjugated carbon chain, the more intense is the absorption and the longer the correspond λmax.

For example see the difference between the absorption spectra of buta-1,3-diene (colourless) and carotene (orange), the deriving from multiple conjugated C=C double bonds.

(ii) The intensity of absorption also depends on concentration of the solution of the compound under investigation or analysis.

(iii) Colour chemists are very interested in not just the absorption spectra and resulting colour, but how intense is the colour for commercial dyes for fabrics - the latter affects the quantity of the organic molecule needed.

The shape and width of the absorption band/bands controls the shade and purity-quality of the colour observed - you can even do reflectance spectra of a dye absorbed on a fabric to compare it with its solution..

(d) Reflectance spectra

Absorption spectra of coloured compounds are obtained using a solution of the in a colourless-transparent solvent.

However, if the material is an insoluble or is being investigated as part of a surface e.g. a forensic examination of a valuable painting, then a different technique is used to obtain a reflectance spectrum.

In this technique white light is shone onto the surface of the solid (pigment powder or painting surface) and the reflected light analysed in a reflectance spectrometer.

Here you are analysing the light wavelengths reflected, NOT the light wavelengths absorbed.

A reminder diagram to show the origin of the blue colour of a molecule in terms of transmission light through its blue solution and reflectance from the surface of the solid.

diagram of a pigment explaining transmission and absorption of visible light to explain its blue colour in solution and the solid

If the molecule absorbs in the yellow (or green-red) it appears violet-blue (the complimentary colour).

See methylene blue on the page of aromatic uv-visible spectra


We can now compare the difference between the absorption spectrum and the reflectance spectrum for the same blue pigment (or dye) described above.

diagram comparing explaining the difference between an absorption spectrum and reflectance spectrum for a blue pigment dye spectra

 The bulk absorption, mainly in the red-orange visible light region, and not in the blue-green region, produces a blue solution absorption spectrum.

 The reflected wavelengths, mainly in the visible light blue region, and surface absorption of the red-orange region, produces a blue solid reflectance spectrum.

(e) Examples and theory of some coloured inorganic compounds

(i) Transition metal complexes (confined to the 3d block)

Many compounds of transition metals e.g. complex ions, are coloured inferring the ∆E for electronic changes involving the central metal ion can be caused by photons in the visible region of light.

The ligands in transition metal complexes cause a splitting of the d orbitals in the d sub-shell (see diagram below).

The observed colour resulting from the ∆Eelec changes due to the 3d (or any d) orbital splitting depend on:

(i) the d electron configuration of the central metal ion and its oxidation state - the electronic state of the d orbital sub-shell,

(ii) the nature of the ligand and strength of its bond with the central metal ion - different ligands have different effects on the relative energies of the d orbitals of a particular ion.

(iii) the number and spatial arrangement of the ligands - this affects the splitting of the d sub-shell.

This can be illustrated by a simple reduction experiment using a soluble vanadium(V) and reducing it with zinc and dilute sulfuric acid. A series of vanadium ions of ever decreasing oxidation state are formed and all have a different colour.

Vanadium V(+5, yellow)  ==> V(+4, blue)  ==>  V(+3, green)  ==> V(+2, violet)

For full details of these chemical reactions see the chemistry of vanadium

The colours originate from the splitting of the 3d orbitals under the influence of the ligands - illustrated above.

For octahedral complexes the split is 3 lower and 2 higher 3d orbitals and all the diagrams refer to [M(H2O)6]n+.

As a result, an electron can be excited-promoted from a lower 3d orbital to a higher.

In the left diagrams above (1) and (4) represent situations where an electronic excitation to produce colour cannot happen because either, there is no electron available for excitation (1), or there is no available orbital to accept an excited electron (4).

In the case of (2) and (3) it is possible to excite an electron from a lower 3d level (top-left diagram) to a higher 3d level (bottom-left diagram).

The right-hand diagram shows the colours and electron configurations of various hexaaqua 3d block transition metal ions.

For more details see Electron configuration of transition metal ions and colour theory

and for examples of transition ion absorption spectra see

08 The uv-visible absorption spectra of some copper complex ions

01 The uv-visible absorption spectra of some cobalt complex ions

05 The uv-visible absorption spectra of selected titanium complex ions

12 The uv-visible absorption spectra of selected nickel complex ions

13 The uv-visible absorption spectra of selected manganese complex ions

06 The uv-visible absorption spectra of some chromium ions


(ii) Absorption spectra of the halogen molecules (separate page)

(f) Examples and theory of some coloured organic compounds - chromophores

Many coloured organic compounds contain unsaturated groups:

 e.g. C=C (alkene), (aromatic), C=O (carbonyl), -N=N- (azo linkage)

When these groups form part of an extended delocalised electron system (conjugated) the electronic energy levels can be low enough for photons of visible light to be absorbed and electrons promoted to a higher level - giving the molecule its colour.

Electrons in double bond systems (π orbitals) are more spread out than single bonds (σ bonds), and, dispersion of charge lowers the potential energy, so these 'pi' electrons require less energy to be excited to a higher level i.e. visible light photons.

However, if lone pairs of electrons can interact with the pi bond systems, they are effectively become part of the chromophore and modify its colour effect, particularly if attached to a benzene ring.

e.g. the functional groups -OH, -NH2, -NR2 (R = alkyl), where the N or O atom has 1 or 2 lone pairs of electrons, all of which are known interact with pi electron clouds if they are directly attached to a benzene ring.

Also groups such as -NO2 or -SO3- extend the delocalised system of the benzene ring and hence change the absorption spectra and the colour of the compound.

Small changes in the energy of the delocalised pi electrons affects the Evis needed for excitation and so affects absorption and the resulting colour of the compound.

The part of the molecule, including the extended delocalised electron system, that allows visible light absorption, hence colour, is called the chromophore.

BUT, note that some small parts of the molecule might be described as the chromophore e.g. the -N=N- azo group, so take care when having to describe the chromophore in a given molecule.


(i) Example of an alkene chromophore

skeletal formula isoprene beta-carotene 2-methylbuta-1,3-diene molecular structure advanced organic chemistry

The beta-carotene molecule is an orange organic pigment found in carrots

Note that the 'monomer' Z-2-methylbuta-1,3-diene (isoprene, shown on the right), with a much shorter delocalised electron system, from which beta-carotene is biosynthesised, is colourless.

The extended conjugated alternate C-C single and C=C double bond system becomes the chromophore, sufficiently low in energy to allow pi electron excitation by visible light photons to give an orange colour.

For more details see The absorption spectra of alkenes


(ii) Examples of dyes e.g the azo dye chromophore

The diagram below illustrates absorption spectra of three dyes with different visible light absorption bands, giving rise to three different colours.

simplified diagram of the visible light absorption spectra of a yellow, red and blue dye molecule azo dyestuff

Dye A absorbs in the blue and will look yellow, λmax 410 nm.

Dye B absorbs in the blue-green-yellow and will look red-purple?, λmax 510 nm

Dye C absorbs in the yellow-orange, will look blue-cyan, λmax 590 nm

A series of dyes often contain the same chromophore e.g. C6H5-N=N-C6H4

(c) doc b is orange(c) doc b is yellow

For more details of absorption spectra of dyes and natural pigments see:

02 The uv-visible absorption spectra of alkenes

03 The uv-visible absorption spectrum of chlorophyll, photosynthesis (porphyrin pigment)

04 The uv-visible absorption spectra of the photopigments in the human eye

11 The uv-visible absorption spectra of some azo dyes

11 The uv-visible absorption spectrum of methylene blue

16 The uv-visible absorption spectrum of haemoglobin (porphyrin pigment)

17 The uv-visible absorption spectrum of melanin pigments


All Advanced Organic Chemistry Notes


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