The interaction of electromagnetic radiation and The interaction of electromagnetic radiation and The interaction of electromagnetic radiation and The interaction of electromagnetic radiation and mattermattermattermatter

Name ……………………………………………………………………………………..

Archbishop McGrath Catholic High SchoolYsgol Uwchradd Gatholig Archesgob McGrath

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The interaction of light and matter

Most of our understanding of the electronic structure of atoms has come from the area of science known

as spectroscopy — the study of how light and matter interact. To understand this we must understand

the nature of light.

Chemists use two models to describe the behaviour of light - the

• wave model and the

• particle model.

Neither theory fully explains all the

properties of light. Some are best described by

the wave model: the particle model is better

for others. We choose the theory which is most

appropriate to the situation. Usually this

depends on the dimensions –

• for large scale, the particle theory is

better

• for very short distances, the wave

model is more appropriate

When we use the term 'light' we normally

mean the visible light to which our eyes

respond. But visible light is only a small

part of the electromagnetic spectrum, which

includes all the different forms of

electromagnetic radiation. There are other

regions, for example: radio waves, ultra-

violet, infra-red and γ-rays as shown on the

right.

Complete the following table, noting the wavelengths of different colours of light

Colour Wavelength (nm)

red

orange

yellow

blue

green

indigo

violet

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Different ideas about the true nature of light

The wave theory of light

Light is one form of electromagnetic radiation. Like all electromagnetic radiation, it behaves like a wave,

with a characteristic wavelength and frequency.

Amplitude, Wavelength, Frequency, and Speed

The amplitude (a) is the total distance

betweeen the crest of a wave and the centre line.

It's a distance, so can be measured in metres,

centimetres, millimetres or smaller units of length.

The wavelength (l) is the distance between

one peak of the wave and the next peak. It's a

distance so can be measured in metres,

centimetres etc. It is sometimes given the Greek

letter λλλλ (lambda). It's also the distance

between one part of the wave and the next part.

The frequency (f) is the number of complete waves passing a point each second. It's a 'number per

second' so it's measured in /s or s-1; usually called hertz (Hz) after a German physicist.

1 kilohertz = 1 kHz = 1000 Hz

1 megahertz = 1 MHz = 1,000,000 Hz

The speed of a wave (v) is just what it says - the speed at which the vibrations in the wave move

from one point to the next. Wave speed is measured in metres per second (m/s, ms-1).

Speed of light

Frequency and wavelength are very simply related. Multiplying the wavelength and frequency together

gives us a constant you get a constant. In the case of light, this constant is the speed of light.

A wave of light will travel the distance between two points in a certain time. It doesn't matter what kind of

light it is, the time is always the same.

The speed at which the wave moves, the speed of light (symbolised by c). is the same for:-

• all kinds of light, and for

• all kinds of electromagnetic radiation.

It has a value of 2.99 x 108ms-1 when the light is travelling in a vacuum.

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The speed of light is given by:-

c = f λλλλ

where:

c= speed of light

f = frequency

λ = wavelength

There is another parameter we use in infra red (ir) spectroscopy: wavenumber. This is simply the

reciprocal of the wavelength and is the number of waves per unit of distance. It is useful in ir

spectroscopy since it gives convenient numbers to handle and, in ir spectroscopy, has units cm-1.

Waves of different wavelength

Waves of different frequency

Waves of the same frequency but different

amplitude

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Theory of electromagnetic radiation

All electromagnetic radiation has fundamental

properties and behaves in predictable ways according

to the basics of wave theory.

The Victorian theoretical physicist, James Clerk

Maxwell, deduced the nature of electromagnetic

radiation to be a combination of magnetic and electric

fields

Electromagnetic radiation consists of an electrical field (E) which varies in magnitude in a direction

perpendicular to the direction in which the radiation is traveling, and a magnetic field (M) oriented at

right angles to the electrical field. Both these fields travel at the speed of light.

Light which is radiating in all directions is termed non-polarised light.

It is possible to effectively filter the light and permit through only light radiating in a particular plane.

This is called plane-polarised light and is extremely important in Chemistry since the only difference in

the properties of optical isomers is the ability of optical isomers (d- and l- or R- and S-) to rotate plane

polarised light in opposite directions by the same angle of rotation. It is also the reason why PolaroidTM

sunglasses work – by effectively restricting the total amount (intensity) of light passing through the lens.

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The particle theory of light

In some situations, the behaviour of light is easier to explain by thinking of it not as waves but as

particles. This idea, originally suggested by Isaac Newton, (who referred to ‘corpuscles’ of light) was

used by Albert Einstein in 1905.

As the name suggests, the particle theory of light regards light as a stream of tiny 'packets' of

energy called photons.

The energy of the photons is related to the position of the light in the electromagnetic spectrum.

For example, photons with energy 3 x 10-19 J would correspond to red light.

Combining the wave and particle theories of light

The two theories of light - the wave and photon models - are linked by a relationship which was

proposed by Max Planck around the turn of the century:

E = hf

where

• E = the energy of a photon

• f = frequency of the light on the wave model multiplied by a constant

• h is a constant (the Planck constant and has a value of 6.63 x 10-31 J Hz-1.

Example

For a photon of red light with an energy of 3 x 10-19 J, the frequency would be given by

3 x 10-19 = 6.63 x 10-31 f

and so

f = 3 x 10-19 / 6.63 x 10-31 = 4.5 x 1011 Hz

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Emission spectra

It has been known for hundreds of years

that white light contains light of all the

colours of the rainbow.

It has also been known for a long time that

white light can be split into its constituents

colours using a prism.

It has also been known for a long time that different elements may have different characteristic colours

associated with them. We can use this property in the flame tests where we note any characteristic

colour. Equally, we could pass the emitted light through a prism and notice the different individual

colours.

Atoms can become excited by absorbing energy, for example, from:-

• flames

• an electric discharge or from

• radiation in the stratosphere or in outer space.

When atoms lose energy, it is often emitted as electromagnetic radiation. This radiation is

usually in the infra-red visible or ultra-violet region. The emitted light can be split up into an

emission spectrum by passing it through a prism or a diffraction grating (see below).

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The spectrum of white light contains a continuous series of lines and so a continuous spectrum

contains all the colours of the rainbow.

Different elements may have different colours and it has long been known that some elements can

be identified in this way.

Splitting the light emitted by different elements into the spectrum demonstrates that all elements

have different, characteristic spectra as shown by the absorption spectra of H, Hg and Ne below.

• If a cold gas is given energy, then

the atoms can absorb light of

particular frequency – this as an

absorption spectrum.

• If a gas is hot then it can emit light

of particular frequency. This is an

emission spectrum.

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• In absorption spectra, the lines appear black since the light is absorbed and then re-emitted in

an infinite number of directions. Effectively, almost no light of these particular wavelengths

then get to the detector and so these wavelengths appear to be black. The black lines in the

absorption spectrum are at exactly the same wavelength as the coloured lines in the emission

spectrum.

The sequence of lines is characteristic of each element and, like human fingerprints, can be used to

identify the element.

The composition of stars, for example, can be determined from vast distances away by using this

technique. This is also how helium was first detected and the element was named after the Greek

word for the sun, ‘helios’.

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The spectrum of hydrogen atom

The electromagnetic spectrum contains many wavelengths not visible to the human eye. This does

not mean there are no absorptions or emission outside of the visible region. There are. The

spectra are often named after the physicist who discovered them:-

• Lyman

• Balmer

• Paschen

• Brackett

• Pfund

The Balmer series is shown schematically below.

It is very important to notice that:-

• There are discrete lines

• As the wavelength decrease, the lines get closer and closer together

The full hydrogen spectrum contains the other series of lines, one in the visible and several in the

infra-red region. The spectrum was interpreted in 1913 by the Danish scientist, Niels Bohr.

Bohr's theory explained why the hydrogen atom only emits a limited number of specific

frequencies. The frequencies - predicted by Bohr's theory matched extremely Well with the

observed lines.

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How Bohr's theory explained emission spectra

The basic idea behind Neils Bohr-s theory is this - atomic emission spectra are caused by electrons in atoms

moving between different energy levels (also known as shells and sub-shells).

When an atom is excited, electrons jump into higher energy levels. Later, they drop back into lower levels

and emit the extra energy as electromagnetic radiation, which gives an emission spectrum.

Bohr-s theory not only explained how the cause of absorption and emission spectra, it also gave a new model for

the electronic structure of the atom.

However, Bohr's theory was controversial –

• It made use of Planck’s new idea of quantisation of energy which was at odds with much of what was

then thought about energy.

• It only explained the spectra of hydrogen. Bohr himself stated that he could not use it to explain the

spectra of helium or larger elements however since he could calculate the spectrum of hydrogen.

Since the theory fitted completely with experimental observation, there was clearly something

significant in it even if the theory was not complete.

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Summary of Bohr’s theory of the hydrogen atom

• The electron in the H atom is only allowed to exist in certain definite energy levels

• A photon of light is emitted or absorbed when the electron changes from one energy level to another

• The energy of this photon is equal to the difference between the two energy levels (∆E)

• The frequency of the emitted or absorbed light is related to ∆E by:

∆E = h f

where

• h is Planck’s constant

• f is the frequency of the absorbed (or emitted) light.

There are refinements to be made to Bohr’s model since the lines in emission and absorption spectra get

closer together with decreasing wavelength and eventually converge, producing a continuum.

This suggests that the shells are not equally spaced:-

and that, eventually, the electron has enough energy to escape the atomic environment entirely. The

frequency where the continuum begins represents the 1st ionization energy of the atom.

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Ionisation energy

Notice that as energy increases the levels become more closely spaced until they converge. After this

point, which corresponds to the electron breaking away from the atom, the electron is free to move

around with any energy. The H atom has lost its electron and become an H+ ion. This is ionisation

and the energy difference between this point and the ground state is called the ionisation energy.

Ionisation can be represented by the equation

X ( g ) X+( g ) + e -

where X stands for an atom of any element and e- for an electron.

Notice that X is shown as X(g), indicating that the atoms are separated from one another, in the

gaseous state.

In the case of hydrogen, we can work out the ionisation energy.from

the point where the lines of the Lyman series converge together.

The entire spectrum of hydrogen (Lyman, Balmer, Paschen etc;) can be explained by this theory

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Summary

• In the Bohr model, the rings represent the energy levels of the electron in the hydrogen atom.

• The further away from the nucleus, the higher the energy.

• Levels are labelled with numbers starting at 1 for the lowest level — the ground state.

• The lines of the Lyman series correspond to changes in electronic energy from various upper levels

to one common lower level, level 1. Each line corresponds to a particular energy level change,

such as level 4 to level 1.

• The series of lines which lies in the visible region, the Balmer series, arises from changes to

level 2 from levels 3, 4, 5 .... etc.

Example questions

The diagram below shows a part of the atomic emission spectrum of hydrogen

Explain why it consists of a series of sharp lines and is not a continuous spectrum [2]

QWC [1]

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What happens when radiation interacts with matter?

Energy interacts with matter

Electromagnetic radiation can interact with matter, transferring energy to the chemicals involved.

The chemicals absorb energy, and the absorbed energy can make changes happen in the chemicals.

Just what changes occur depend on

• the chemical involved

• the amount of energy involved_

We must also remember that molecules are doing energetic things all the time. They:-

• move around

• rotate and

• the bonds in the molecule vibrate.

• The electrons in the molecule have energy too and they can move bebween the different

electronic-energy levels.

A molecule has energy associated with several different aspects of its behaviour, including:

• energy associated with translation (the molecule moving around as a whole)

• energy associated with rotation (of the molecule as a whole)

• energy associated with vibration of the bonds

• energy associated with electrons.

These different kinds of energetic activities involve different amounts of energy-- for example, making

the bonds in a molecule vibrate generally involves more energy than making the molecule as a

whole rotate. The energy needed increases in the general order shown below.

electronic energy

vibrational energy

rotational energy

translational energy

Increasing energy

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Quantisation of energy

Max Planck originally suggested that energy could exist as definite, fixed packets which were termed

quanta. At the time this was a revolutionary idea since classical physics assumed that bodies could

have energies of any amount i.e. there was a continuum of energy.

Niels Bohr suggested that electrons in the hydrogen atom could only have fixed amounts of energy

and that the proposed transitions between energy levels would only happen if the electron was

given precisely the correct amount of energy to make the transition – too much or too little and the

transition would simply not occur. Likewise, when the electrons moved from a higher energy level

to a lower level, a fixed amount of energy was released.

We can extend these ideas to other types of energy – generally all other types of energy

(translational. rotational and vibrational) are quantised as well.

Change occurring. Size energy change (J) Type of radiation absorbed

Change of

− rotational energy level

1 x 10-22 to 1 x 10-20

microwave

− vibrational energy level 1 x 10-20 to 1 x 10-15 infra red

− electronic energy level 1 x 10-15 to 1 x 10-16 ultra violet and visible (uv/vis)

The greater the transition, the larger the energy change and the higher frequency (lower

wavelength) of light absorbed or emitted.

Excitation energy

Excited state

Ground state

Small jump – corresponds to

visible light

Large jump – corresponds to

uv light

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Colours of materials or why are plants green?

We learn lower down the school that objects are a particular colour because of absorption of the

other colours. For example, a red object is said to be red because it does not absorb red but

simply reflects that colour.

Examining the amount of uv/visible light absorbed by a chemical compound is termed uv/visible

spectroscopy. The basic principle is straightforward. A sample tube is placed in a beam of light.

The amount of light absorbed at each wavelength examined is measured and a graph is drawn.

This is the uv/vis spectrum of chlorophyll.

The higher the peak at any wavelength, the more

light of that wavelength is absorbed.

It should be clear that chlorophyll absorbs blue

light and red light. Since chlorophyll a does not

absorb the green light, chlorophyll a and hence

plants are green.

light source

sample

detector

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Interpreting a spectrum

uv/vis spectra can be deceptively simple since we are examining electronic transitions. However,

bearing in mind that the molecule can also be moving and rotating, the broad peaks show electronic

transitions with all these other transitions superimposed. When recording the spectrum, chemists often

give the wavelength of the maximum absorption (λλλλmax).

We are usually only really concerned with three features of a uv/vis spectrum:

• the wavelength of the radiation absorbed (remember, for a compound to he coloured, at least

part of the absorption band must he in the visible region)

• the intensity of the absorption

• the shape of the absorption band.

For example, carotone has the following spectrum:

λλλλmax = 453nm (blue region) and hence carrots are orange.

The intensity of the absorption depends on the concentration of the solution and on the distance the

light travels through the solution. Standard molar values are quoted so that values for different

compounds can be compared.

It is enough to know that the higher the peak, the more intense the absorption. With substances used

as dyes, the intensity of the absorption is important commercially since because it determines the

amount of pigment or dye needed to produce a good colour. The shape and width of the absorption

band is important because it governs the shade and purity of the colour seen.

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This is the spectrum of oxygenated and de-oxygenated haemoglobin.

Why is blood red in healthy people? …………………………………………………………………………………………………………

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Why do people suffering from lack of oxygen appear to have a blue colouring? ……………………………………

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Mark scheme for specimen question