We use photochemistry to make nylon and vitamins and to help babies to
survive. Sunlight interacts with the atmosphere, cells of plants and the
surface of our skin. Some of these reactions are vital, some are dangerous,
and some we would like to harness to create a cleaner, safer world
HALF of all new-born babies become jaundiced because their livers cannot
cope with a bodily waste product called bilirubin which comes from the
breakdown of red blood cells. Too much bilirubin can cause brain damage. In
the womb, bilirubin is removed across the placenta, but newly born babies have
to manufacture their own enzymes to make bilirubin soluble in waste fluids.
These enzymes are missing to begin with, so how do babies cope?
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In 1958, a nursing sister at Rochford General Hospital in Essex noticed
that babies on the sunny side of the maternity ward seldom became jaundiced.
The reason, we now realise, was that the bilirubin in their skin was being
made soluble by exposure to light. A photochemical reaction brings this about
by changing the shape of the molecule. Today every maternity hospital prevents
jaundice in new-born babies by bathing them in blue light.
Another disease that can be treated by light is cancer, and trials have now
reached an advanced stage in the technique which is called phototherapy.
Doctors inject cancer patients with a 鈥渟ensitiser鈥� molecule which is absorbed
over two or three days, mainly by cancer cells. The doctors then insert fibre-
optic probes into the patient which illuminate the cancer with red laser light
(wavelength 600-800 nanometres). Red is chosen because it is not absorbed by
healthy tissue.
The sensitiser is a so-called 鈥減orphyrin鈥� made from haemoglobin, a
component of blood. Porphyrin turns skin green, and patients injected with the
drug should avoid sunlight, but the cancer cells retain most of the drug. The
laser activates the sensitiser; which also fluoresces so we can see it
absorbing the light. The activated molecule excites oxygen in the cell to form
singlet oxygen. This is dangerously reactive and kills the cell.
A group headed by David Phillips at Imperial College, London, has discovered a
more sophisticated type of sensitiser, phthalocyanine, that is similar to a
porphyrin but very much more effective. Phthalocyanines, which turn the skin
blue, can be made soluble in water and are selectively absorbed by cancer
cells. They fluoresce and give a long-lived activated state that makes them
particularly deadly to cells that contain them. Tests by Stephen Bown at
University College Hospital, London, have shown that these compounds kill
cancers of the brain, bladder, pancreas and intestine. So far, trials have
been limited to mice and rats, but clinical trials are now being planned.
Making light work
Jumping electrons
CHEMISTRY is all about energy and molecules: energy can help to make
molecules, change them and break them. Photochemistry is all about exciting
the electrons of molecules with discreet packages of light energy called
photons, thereby making them jump from a lower to a higher energy state.
Photons of ultraviolet (UV) and visible (VIS) light can do this. Our eyes
rely on such photochemical changes in the molecules of the retina to enable us
to see.
Shorter wavelengths of light have more energy and cause bigger electron
jumps. Light may even have enough energy to strip out an electron completely
and turn a molecule into an ion. However, photochemists are more interested in
light which excites electrons rather than light which ejects them. They study
the remarkable behaviour of molecules once they have captured a photon of
light. What happens next depends on the molecule: it may break apart, possibly
to form highly reactive free radicals; or react chemically with another
molecule; or rearrange itself into another isomer; or emit light.
Photochemical reactions are very different from ones driven by heat. A photon
of light can excite a molecule in a way that heat cannot.
Nearly all molecules respond to ultraviolet and visible light. Some
molecules are particularly sensitive and have evolved in nature or been
synthesised to be so. The best-known naturally photosensitive substance is
chlorophyll which enables plants to make carbohydrates with energy from
trapped sunlight. Perhaps the best-known artificial photosensitive material is
photographic film. Other applications include thin films, liquid crystals,
polymers and microelectronic devices that we now use to detect light, to count
photons, to transport light, to record light, and even to change the
wavelength of light.
Some chemists are working to trap light and charge so-called molecular
batteries, rather like chlorophyll. In a molecular battery, a molecule excited
by light moves an electron to a more stable neighbouring site, and in so doing
creates a positive and negative pole 鈥� a miniature battery. Discharging such a
molecular battery could, in theory, be used to drive a difficult chemical
reaction, as happens in photosynthesis. Some zirconium compounds have been
shown to behave as molecular batteries, although their usual limitation is
that they stay charged only for fractions of a second.
Light which produces photochemical changes generally has wavelengths of 200
to 740 nanometres (a nanometre is a billionth of a metre and given the symbol
nm). We see the visible light in this range, but not UV light. The 鈥渃olours鈥�
within this band of the electromagnetic spectrum are shown in the adjacent
diagram. The sources of light suitable for photochemical investigations are
described in the above box.
When a molecule absorbs a photon of UV-VIS radiation and is excited, how
can we study the excited state and follow what happens next? The excited state
may last a few seconds but more likely it will last for only a microsecond (a
millionth of a second), a picosecond (a trillionth of a second) or just a few
femtoseconds (a million-billionth of a second). A femtosecond is equivalent to
a minute in the lifetime of the Earth. We can carry out such studies thanks to
the pioneering work of a British chemist called George Porter, who developed a
technique known as flash photolysis. He shared the 1967 Nobel Chemistry Prize
for the study of ultra-fast chemical reactions with the German Manfred Eigen
and the Briton Ronald Norrish.
Chemists have used femtosecond flash photolysis to study photosynthesis,
the process by which plants harvest energy from sunlight. Cells called
chloroplasts in green plants have a reaction centre with two chlorophyll
molecules. The chloroplast absorbs a photon of light and in about one
picosecond this causes an electron to be transferred out of the reaction
centre leaving the system with enough energy to oxidise water in the cell to
oxygen, which is released into the air:
2H2O 鈫� O2 + 4H+ + 4e鈥�.
The hydrogen ions and the electrons released in the cell then bind to
carbon dioxide to form carbohydrate molecules such as starch or cellulose. The
process is complex but we know the oxidation takes place at a cluster of four
manganese atoms situated on one of the chlorophyll molecules. Photosynthesis
enables an energetically unfavourable reaction to take place, and this
supplies the Earth with plants and its atmosphere of oxygen:
6CO2 + 6H2O 鈫� C6H12O6 + 6O2
Shielding life on Earth
Airborne activity
IN the Cambrian period of geological time, about 600 million years ago,
only 1 per cent of the atmosphere was oxygen. Life-threatening UV-B and UV-C
radiation from the sun could reach the surface of the planet, so life only
survived where it was shielded, in stagnant water for example. But as marine
plants released more and more oxygen to the atmosphere, life on the surface
became less hazardous, thanks primarily to the formation of a protective
shield of ozone in the upper atmosphere, or stratosphere.
When molecular oxygen, O2, is exposed to UV-C, it is converted into
ozone by a photochemical reaction. Light splits the O2 into two
atoms of oxygen, each of which can then combine with another molecule of
O2 to form ozone, O3. This process occurs in the stratosphere where
there is now an estimated 3.5 billion tonnes of ozone. But the photochemistry
does not stop there, because the ozone is also capable of absorbing UV-B
radiation. When it does, it throws off one oxygen and reverts to
O2. The result is that only a little UV radiation of wavelength
shorter than 300 nm reaches the surface of the Earth. But some does, and this
is life-threatening (see diagram and box opposite).
Ozone in the stratosphere is depleted by reactions with nitrogen oxide, NO,
the hydroxyl radical (OH.) and by chlorine radicals (Cl.). NO and OH. are
considered 鈥渘atural鈥�, whereas the chlorine is a cause for concern. Chlorine is
being transported to the stratosphere as the notorious chlorofluorocarbons,
which are used in refrigerators and for more than 20 years were used in
aerosols. Because these are very stable molecules, they survive to reach the
upper atmosphere, where UV radiation eventually breaks them apart. Once
released from its CFC, a chlorine radical can turn 1000 ozone molecules into
oxygen molecules. Hence the current concern.
Smog and acid rain are other legacies of photochemical changes to airborne
pollutants. Coal and oil burnt in power stations emit copious amounts of
sulphur dioxide which dissolves in rain to form the weak acid, sulphurous
acid. However, if the sulphur dioxide is exposed to sunlight it can react with
other airborne particles to form sulphur trioxide which dissolves in rain to
form the stronger sulphuric acid.
Sunlight also changes waste gases from vehicle exhausts, paradoxically into
ozone, a gas which is damaging to lungs and plants at ground level. The
sequence of steps involves oxides of nitrogen. First there is NO and
NO2 (together called NOx) produced in car engines. The
NO2 decomposes photochemically to NO and O, and the atom of oxygen
then goes off to form O3. The ozone may then react with pollutants
and natural materials from plants to form peroxyalkyl nitrates that
characterise photochemical smog. In cities such as Los Angeles, which has the
highest density of cars per head of population, there are two daily peaks of
NOx corresponding to the morning and evening rush hours and a mid-
day peak of ozone generated by sunlight at noon (see Figure).
Light industry
Electrons at work
PHOTOCHEMISTRY has limited application in the chemicals industry, which
still drives most reactions with heat not light. The mercury arc lamps that
drive photochemical processes are not cheap, so they are only economical for
making compounds, such as drugs and cosmetics, that are highly valuable even
in very small amounts. Ideally, only one reactant absorbs light and becomes
activated. If other ingredients absorb light, they too may undergo reactions
and reduce the yield or contaminate the product.
Cholesterol, a fatty component of blood, can be converted to vitamin D, and
one step in the process is best done photochemically. The form of the vitamin
this produces, called D3, is vital to poultry and it is added to their feed.
Vitamin D can also be made by irradiating certain yeasts which are rich in
ergosterol. The vitamin D content of milk can be raised by irradiating it, but
this also changes the taste so it is usual to add the vitamin directly.
The German chemical giant, BASF, has a process for making vitamin A. Light
is used to twist double bonds through 180掳 from a so-called cis to a trans
arrangement. This is done with visible light in the presence of a
鈥渟ensitiser鈥�, in this case chlorophyll. Sensitisers, and dyes such as
methylene blue, are also used in photo-oxidation reactions in which oxygen is
excited to its singlet state. This super-active form of the gas is used to
make artificial fragrances such as rose oxide, starting from more available
perfumes such as citronellol which is derived from lemon grass.
Few companies make bulk chemicals photochemically, but one, Toray Industries
of Japan, devised a economical method of making caprolactam, the material from
which nylon is made. Toray鈥檚 plant can produce up to 135 000 tonnes of
caprolactam a year from cyclohexane and nitrosyl chloride (NOCl). The company
bathes the reactants in light from a high-pressure mercury lamp which can
operate efficiently for more than a year.
The light splits the NOCl into chlorine radicals, Cl., and nitrosyl
radicals, NO. The chlorines snip a hydrogen atom from the cyclohexane, and
then the nitrosyl radicals move in to form caprolactam from which nylon is
made. The printing industry also nowadays dries inks photochemically in a
fraction of a second.
Sometimes, we may want things to be destroyed by light, and degradable
plastics come into this group (see Inside Science No 50). There are also
instances when we want to protect plastics and polymers from the effects of
light and then we add absorbers to prevent UV light being absorbed by
the plastic, or antioxidants to mop up any free radicals that are produced.
Exciting fuel cells
Return of water power
HYDROGEN gas poses attractions as a fuel 鈥� all it produces on burning is
water. However, there are no natural sources of hydrogen gas, and industry
harvests it from methane gas, CH4, by partial oxidation which
liberates hydrogen, H2, and forms carbon monoxide, CO. Cheap
electricity would allow chemists to harvest hydrogen from water by
electrolysis, but another method using water as the source of hydrogen has
even more appeal; photochemical decomposition by sunlight:
2H2O + UV 鈫� 2H2 + O2
water + sunlight 鈫� hydrogen + oxygen
In the 1970s it was first suggested that this process was possible using a
catalyst such as titanium dioxide. If particles of this were fine enough and
they were activated with other metals such as platinum, UV light from the Sun
could be trapped and excite a platinum atom which would then have the energy
to reduce water to its elements. It worked! However it is very inefficient and
yields of hydrogen gas were low. Recently, Japan chemists Hironori Arakawa and
Kazuhiro Sayama of the National Chemical Laboratory for Industry at Ibaraki
discovered that by adding sodium carbonate (washing-soda crystals) to the
water they can boost the amount of gas given off. Their best results came from
a catalyst of TiO2 doped with 0.3 per cent platinum, and 0.3 grams
of this catalyst can produce 2 millilitres of hydrogen gas per hour. A tonne
of catalyst could yield about 75 000 litres of hydrogen gas (or 100 litres of
liquid hydrogen) per sunny day 鈥� enough to run a family car.
In 1990 the Musashi-8 car was demonstrated by Nissan and it travelled 300
kilometres on 100 litres of liquid hydrogen. A European project involving
Germany鈥檚 Daimler-Benz, the Paul Scherrer Institute in Switzerland, and a
British catalyst firm OUP is also looking at ways of using hydrogen as a fuel
for vehicles. They point out that with a fuel cell operating an electric
motor, hydrogen has an efficiency of 60 per cent compared with the 35 per cent
of an internal combustion engine.
Sources of light used by photochemists
FOR a photochemical experiment chemists ideally need monochromatic light,
light of one colour and preferably of just one wavelength. Most sources of
light are white (polychromatic) and have to be 鈥渇iltered鈥� to remove the
unwanted colours, or passed through a prism or grating to get the required
colour. Incandescent tungsten-filament lamps are good sources of visible light
but the glass bulb itself cuts out UV below 350 nm. By adding a little iodine
and placing the lamp inside a quartz bulb, chemists can generate UV light down
to 200 nm. Oxygen in the atmosphere absorbs UV with shorter wavelengths than
this, so chemists can only work at these shorter wavelengths with the help of
vacuum apparatus and special materials for lamps, windows, lenses and reaction
cells.
Discharge lamps contain xenon gas or mercury vapour through which an
electric discharge is passed. The light they give out may be mainly at one
wavelength so they are essentially monochromatic. Mercury lamps at low
pressure emit 253.7 nm UV, and at medium pressure emit 184.9 nm. Xenon lamps
discharge at 147.0 nm. Discharge lamps can discharge pulses of light.
Lasers are particularly important in photochemical research because they
meet the requirements of being monochromatic, very intense and can deliver
femtosecond pulses. These can also be steered and polarised into one plane.
The commonly used lasers are neodymium:yttrium-aluminium-garnet (wavelength
1060 nm), helium-neon (633 nm), argon (458-514 nm), nitrogen (337 nm) and
excimer (UV). There are also tuneable dye lasers in the range 200-1100 nm,
and infrared diode lasers, of the type used in compact disc players, that go
down to 633 nm.
Skin and photochemistry
IN the Sun, our skin undergoes photochemical reactions. These may be
beneficial, cosmetic, damaging or even fatal. The Victorians thought sunlight
was beneficial; they knew it promoted healing and prevented rickets, the
vitamin D deficiency that affected inhabitants of smoky industrial cities.
This belief in the blessings of sunbathing lingers on, although most people
now do it for cosmetic reasons to get a golden tan, despite the fact that this
browning is the body鈥檚 response to a threat. Ultraviolet light which causes
tanning is very damaging. At best it weakens the skin and leads to excessive
wrinkling in old age, at worst it causes skin cancers. Every year, doctors
treat more than 25 000 people for this condition in Britain alone. Happily, 95
per cent of patients are cured.
UV-C light kills cells, and light of wavelength 260-265 nm is the most
deadly. This band is exactly that at which DNA and RNA absorb, and may explain
why cells are so vulnerable and strong sunlight is so dangerous. We can
identify pyrimidine and purine groups as the components of DNA and RNA which
absorb at 260 nm. When these are activated they undergo chemical reactions,
either with another group nearby, or with a water molecule, or with oxygen.
The danger of heavy damage is that it will swamp our body鈥檚 repair mechanisms,
and allow a mutant cell to survive. If that cell then replicates we start a
cancer.
We have 100 billion (1011) cells in our body, each of which is
damaged 100 times a day. Some of this damage is caused photochemically and
this is the research interest of Chris Wharton. His research group is using
the special laser facility of the Rutherford Appleton Laboratory in
Oxfordshire to study how UV and X-ray radiation damages cells and how these
then repair themselves.
Cells have evolved to deal with UV light. Many cells contain sunscreens to
prevent damage occurring, others use repair kits, others employ another
photochemical reaction which darkens our skin. Sunbathing starts to redden our
skin after about 12 hours, and this process peaks after 24 hours. Cells die or
deform, and the longer we are exposed, the greater the damage. UV-B is the
culprit; light of around 300 nm penetrates the clear flattened cells on the
surface of our skin and reaches the next layer of cells, which die. One
defence is tanning in which cells called melanocytes in this lower layer are
activated to produce deposits of brown melanin. This compound absorbs UV and
protects the DNA. A natural suntan cuts down UV penetration by 90 per cent.
Alternatively, we can apply a sunscreen lotion to our skin. This will contain
a substance which absorbs in the 290-320 nm region, such as benzoic acid,
menthyl anthranilate, or even zinc oxide.
Most skin cancers are easily noticed and dealt with by a doctor, but one
type called melanoma is dangerous and can be fatal if not treated early. In
Britain there has been an increase of 50 per cent in cases of malignant
melanoma since 1980, and there are about 2700 cases a year. The increase could
be due to stratospheric ozone depletion, the popularity of sunbathing, a
series of fine summers, or cleaner air which no longer cuts out UV over urban
areas.
Some skin conditions can be cured by photochemical therapy. Psoriasis, or
flaky skin, is caused by overactive skin cells. Sensitise these cells with a
compound called psoralen and bathe the skin with ultraviolet light of
wavelength 320-400 nm (UV-A) and you can destroy the overactive cells.
Dry in the blink of an eye
POLYMERS have a special relationship with light: in some cases light is
used to make them or modify them. Yet polymers are always vulnerable to light,
which can break them down. In some cases this is desirable.
Polymers can be modified photochemically by a process called cross-linking
in which strands of polymers are turned into a more rigid network by the
formation either of chemical bonds between them, or by joining them with short
molecular links.
The processes are used to harden plastics, and the same effect makes all
plastics exposed to sunlight brittle, unless they contain additives to
counteract the process. Deliberate cross-linking requires a sensitiser, a
compound such as benzophenone
(C6H5)2C=O, which can be excited by
light sufficiently to grab a hydrogen atom from a nearby strand of polymer.
This leaves behind a free radical centre on the polymer chain which binds with
another polymer chain to form a link. Sometimes, this process of cross-linking
can be amplified by attaching groups to the polymer chain that themselves are
susceptible to photoactivation.
An important commercial use of photochemistry is in the printing of
integrated circuits. Here the polymers exposed to light become cross-linked
and remain fixed in the desired pattern, whereas the polymer beneath masking
material and thus not exposed to light can be washed away.
If we want to cure a hard transparent surface coating with UV we use
benzoin, a derivative of benzophenone that is more sensitive to light. If we
are printing with an ink that contains a pigment that cuts out UV 鈥� such as
titanium dioxide 鈥� and therefore interferes with the rapid curing process, we
would chose a ketone-amine combination which can be photoactivated with
visible light of wavelength above 400 nm. Photolithography, as it is called,
permits a can of Coca-cola to be printed and dried in a hundredth of a second.
Generally, photochemical action damages plastics or polymers, so chemists
take steps to prevent it. Titanium dioxide protects against photo-degradation.
It absorbs the harmful UV light as far as 390 nm. Zinc oxide is almost as
good. Carbon black is ideal, absorbing all light which hits the surface of the
plastic and deactivating any free radicals formed.
But sometimes we want plastics to break down, and if microbes are unable to
do the job, sunlight may be able to achieve it instead provided the plastic
contains peroxide residues left behind from when they were made, or have
carbonyl groups, C=O, built into the polymer chain. These absorb UV
light and once excited tend to rupture the polymer chain. Such photodegradable
polymers can be reduced to fine particles after only a few weeks in the sun.
Further Reading
鈥淭he chemical effects of light鈥�, by G. Oster in Scientific American, volume
219, p158 (1968) is a useful introduction to the subject.
Light, chemical change and life: a source book in photochemlstry, edited by
J.D. Coyle, R.R. Hill and D.R. Roberts (Open University Press, Milton Keynes,
1982). 鈥淧hotodynamic therapy鈥�, by A.J. MacRobert and D. Phillips, Chemistry
and Industry (6 January 1992).
鈥淎 little light relief鈥�, by D. Phillips, Royal Institution Proceedings (1985).
If you want an advanced book on photochemistry then Principles and
Applications of Photochemistry, by Richard Wayne (Oxford University Press,
1989), is excellent. Even more specialised is Essentials of Molecular
Photochemistry, by Andrew Gilbert and Jim Baggott (Blackwell Scientific
Publications, 1991).