快猫短视频

THE STORY began in 1666 when Isaac Newton shone a thin beam of sunlight through a triangular glass prism. As the light emerged, it fanned out into the familiar rainbow spectrum of colours, from red to violet. When the German optician Joseph von Fraunhofer repeated this famous experiment in 1814, he produced a much sharper spectrum by first filtering the sunlight through a narrow slit before it reached the prism. On closer inspection, Fraunhofer noticed that the spectrum was crossed with hundreds of dark lines (later called Fraunhofer lines) corresponding to specific colours or frequencies of light missing from the sunlight itself.

The colours were missing because of a chemical barrier between the source of the radiation at the centre of the Sun and the prism. Atoms of elements such as hydrogen in the outer layers of the Sun absorb light of certain discrete frequencies (or wavelengths 鈥� see 鈥淓lectro-magnetic radiation鈥�) and prevented the colours from reaching the Earth. When Fraunhofer noticed the dark lines, he had observed the first atomic absorption spectrum, and the science of spectroscopy was born.

Most modern spectroscopic techniques depend on shining electromagnetic radiation into a chemical sample and seeing which of its 鈥渨avelengths鈥� become blocked on the way through. The techniques are no different in principle from Fraunhofer鈥檚 experiment. Measurement of an absorption spectrum requires a source of electromagnetic radiation, a sample of atoms or molecules in the gaseous, liquid or solid state, a dispersing element to split the radiation transmitted through the sample into its constituent wavelengths and a detector able to measure the intensity of the transmitted radiation. In Fraunhofer鈥檚 experiment, the Sun was the source of radiation and the outer layers of the Sun provided the sample 鈥� a mixture of atomic gases. He used a prism to disperse the Sun鈥檚 radiation into its constituent wavelengths, detecting the resulting spectrum by looking at the radiation through a theodolite telescope.

In modern spectrometers, chemists can select standard sources of radiation, such as lamps or lasers, that enable them to shine a particular range of wavelengths through the sample (see Figure 1). They might scan the sample with wavelengths from the infrared region of the spectrum, for example, which range from 1 to 100 micrometres. Alternatively, they might shine radio waves through the sample, or ultraviolet and visible radiation. Each band of radiation shone through the sample tells the spectroscopist something new about the molecules present.

Spectral clues

Peaks and plots

DETECTORS are set up to accept and process radiation after it has passed through the sample. Usually, the resulting spectrum is displayed as a graphical plot that has the appearance of a series of peaks. It takes a trained spectroscopist, however, to interpret the spectrum.

Spectroscopists usually make the source as close to monochromatic (that is, single wavelength) as possible and continuously variable, or tunable, to sweep over the required range of wavelengths. They may also provide an arrangement for comparing the intensity of radiation transmitted through the sample with a reference beam, so that any unwanted contributions from a solvent or the glass chamber containing the sample are cancelled out.

A molecule can absorb electromagnetic radiation either from the wave鈥檚 electric or magnetic components. However, most molecular spectroscopy focuses on how molecules interact with the electric component. Molecules absorb electric or magnetic radiation in the form of discrete 鈥減ackets鈥� of energy called photons (see 鈥淓lectro-magnetic radiation鈥�). Because the molecules scavenge these photons from the radiation, the corresponding wavelengths are missing or less intense when the radiation reaches the detector. At the same time, the packet of additional energy from the photon raises each molecule to a so-called 鈥渆xcited state鈥�. Radiation containing photons of different energies produces different kinds of excited states. Visible and ultraviolet light, with wavelengths measured in hundreds of nanometres, can excite a molecule to higher electronic energy states, usually weakening the chemical bonds between its atoms. This is photochemistry, and occurs every time the sun shines on green grass (see Inside Science No. 58).

Lower energy infrared or heat radiation, with wavelengths measured in micrometres (millionths of a metre) can trigger different patterns of vibrational motions in molecules. Like a collection of tiny balls (atoms) linked by springs (chemical bonds), a molecule consisting of many atoms may undergo a complicated set of vibrational motions that are picturesquely called stretching, bending, scissoring, wagging and twisting (see Figure 2). Superimposed on this vibrational motion may be rotational motions; the molecules tumbling through space as they vibrate. A molecule with two atoms, for example, will undergo rotational motions which resemble a dumbbell falling through space. Moving to still lower energies, the absorption of microwave radiation (wavelengths of centimetres or millimetres) simply raises molecules to higher levels of rotational energy.

Molecules do not soak up every single wavelength of energy thrown their way as the radiation source in the spectrometer sweeps through its preset range of wavelengths. This much is obvious simply from the fact that there are Fraunhofer lines: absorption occurs at discrete wave lengths rather than continuously over broad ranges of the spectrum. Instead, energy is absorbed in sharply defined amounts called quanta and the whole process is governed by the laws of quantum mechanics.

These laws explain why there is quantisation of energy in atoms and molecules. Within an individual molecule, such as a water molecule, only certain molecular energy levels are possible, characterised by integer numbers, called quantum numbers. Transitions between energy levels take place in instantaneous quantum jumps. The difference in energy between the two levels involved in the transition (鈻礒) determines the frequency of the photon that can be absorbed through a simple relationship first deduced by the Danish physicist Niels Bohr. This is 鈻礒 = h v, where h is Planck鈥檚 constant and v is the frequency of the photon.

Energy levels

Climbing the ladder

WE CAN imagine that molecules possess a ladder of individual, quantised energy levels (see Figure 3). Energy levels associated with translational motion 鈥� the movement of the molecule through three-dimensional space 鈥� sit at the foot of the ladder but are very close together and thus of little use to spectroscopists. Neglecting translation, the bottom rungs of the ladder (the lowest energy levels) are those associated with pure rotational motion of the molecule. As we move up the ladder we begin to encounter its vibrational levels, with each vibrational level possessing its own subset of rotational levels. Moving higher still, we find electronic levels associated with different patterns or distributions of electron 鈥渃louds鈥� surrounding the atomic nuclei and which determine how the molecule is held together. Each electronic level has its own subset of vibrational levels which, in turn, have their own subsets of rotational levels. Spectroscopists can infer the structure of molecules by mapping out parts of their ladders of energy.

Spectral bands

From lines to humps

THE NUMBER of molecules present in a sample at a particular energy level is called the population of that level. When the molecules absorb radiation from the source in the spectrometer, a proportion of the population of the lower level involved in the transition rises to a higher level.

In principle, each transition between two energy levels gives rise to a line in an absorption spectrum, its intensity governed by the laws of quantum mechanics. In practice, limitations imposed by nature and by the spectrometer being used to generate the spectrum mean that it may not be possible to distinguish individual lines. However, spectroscopists are often able to make use of the information, even if it appears as a broad peak or 鈥渉ump鈥� in the spectrum rather than as a series of discrete lines.

For example, large molecules have a correspondingly large number of different vibrational and rotational energy levels, many of which may be populated at room temperature. Although the energy levels are discrete, they may be too close together to distinguish individually. Instead of the radiation raising the population of one energy level to a single higher level, the populations of many levels are transferred to many higher levels. Furthermore, if the molecules are present in solution, collisions and interactions with solvent molecules such as water can drown out individual features in an absorption spectrum. This effect is most noticeable in visible and ultraviolet spectroscopy, which is carried out using dissolved samples.

The patterns formed by the differently spaced rungs on the ladder of energy levels reveal details of the symmetries, structures and properties of the molecules as well as their identities. Spectroscopists use quantum mechanics to get further clues from the appearance of the spectra 鈥� the sizes and shapes of spectroscopic features. Anyone can examine spectra, but it takes a trained eye to see what鈥檚 there, and a trained mind to interpret the results.

Today a large number of spectroscopic techniques exist for studying molecules, at many different levels of detail. However, three techniques have become routine. Only rarely does a single form of spectroscopy reveal everything 鈥� it often takes a combination of techniques to solve the problem.

Radiation in the range spanning the visible and ultraviolet parts of the electromagnetic spectrum can affect electrons in molecules directly. Absorption by molecules alters electron clouds associated with chemical bonds, and these affect the spectrum produced. Patterns of electron clouds vary widely from one type of molecule to another, depending on the atoms in the molecule, their arrangement in the molecule and its structure, and the nature of the bonds that bind the atoms together.

Visible/ultraviolet spectroscopy is most useful in the investigation of molecules with bonds rich in electrons, as these will have clouds most heavily influenced by the radiation. For example, the ultraviolet spectra of alkenes (hydrocarbon molecules which contain an electron-rich carbon-carbon 鈥渄ouble bond鈥�, C=C) appear quite different from those of aldehydes and ketones (molecules which contain an electron-rich carbonyl group, C=O). The spectra of alkenes exhibit a distinct absorption band resembling a large hump and which peaks when the molecule absorbs radiation with a wavelength of 180 nanometres. This corresponds to the excitation of an electron in the C=C double bond to a higher energy. Aldehydes and ketones exhibit distinct absorption bands by absorbing wavelengths around 280 nanometres. These correspond to the excitation of an electron on the oxygen atom (and is not involved directly in bonding) up to a higher energy level.

The positions of the peaks of these bands, and other bands characteristic of other chemical groups in organic molecules, can depend on neighbouring chemical groupings present in the molecule, such as chlorine atoms or methyl groups (CH₃). However, despite these influences, absorption features due to electron-rich groups such as C=C and C=O usually occur within a relatively narrow range of wavelengths. Because such groups give the molecules both a distinctive absorption spectrum and a distinctive photochemistry, they are known as chromophores (from the Greek meaning 鈥渇lour bringer鈥�).

The extent of absorption at a particular wavelength is related to the amount of absorbing substance present in the sample. Visible/ultraviolet spectroscopy is therefore also a useful means of measuring the concentrations of molecules in solution.

Good vibrations

Infrared spectroscopy

VISIBLE/ultraviolet light affects electron clouds in molecules. Molecules that absorb infrared radiation undergo vibrations and rotations. Patterns of lines in an infrared absorption spectrum are easier to interpret for diatomic molecules such as hydrogen chloride (HCl) containing two atoms than for polyatomic molecules such as water containing three atoms or more. The infrared spectrum of HCl, for example, shows a pattern of lines with an obvious centre and with lines marching out to lower and higher wavelengths (Figure 4).

Each line represents a transition from one combination of vibrational and rotational energy levels to another. The spacing between the lines in the spectrum reflects the molecule鈥檚 so-called moment of inertia 鈥� its resistance to rotational acceleration. This, in turn, depends on the length of the H-Cl bond. Each line in the HCl spectrum appears to be split in two. This is because chlorine has two naturally-occurring isotopes, ³⁵Cl and ³⁷Cl, with a natural abundance of about 3:1. The isotope ³⁷Cl has two additional neutrons in its nucleus and so is a little heavier than ³⁵Cl. Consequently, H-³⁷Cl has a greater moment of inertia. It needs more energy than H-³⁵Cl to force it through a change in vibrational or rotational motion.

The infrared spectra of most polyatomic molecules are not as simple as those of HCl, but yield similar information about the lengths of chemical bonds and the angles between them. Infrared spectroscopy is used more routinely, however, as a means of identifying chemicals or analysing their composition because absorption at certain wavelengths tell the spectroscopist which chemical groupings exist in the molecule and, sometimes, how they are connected. Although the patterns of vibrational motions in a large polyatomic molecule are complex, particular chemical groups produce characteristic bands in an infrared spectrum. These are known as group vibrations and usually reflect the stretching motions of individual bonds such as C=C, C=O, C-H, C-Cl, and so on. Infrared spectra of alkenes, for example, exhibit bands characteristic of C=C stretching vibrations. Aldehydes and ketones have C=O stretching vibrations. As with visible/ultraviolet spectroscopy, the ability to identify the presence of key chemical groups means that infrared spectroscopy provides a powerful tool for chemical analysis. As well as showing what is there, it helps to eliminate things that aren鈥檛 there. For spectroscopists struggling with a mystery organic compound, this greatly reduces the list of suspect structures.

Magnetic attractions

Nuclear resonance

WHILE visible/ultraviolet spectroscopy reveals clues about the structure of the electron-rich molecules, and infrared spectroscopy shows which chemical groupings are there or not, nuclear magnetic resonance (NMR) spectroscopy betrays the way that atoms and groupings in the sample connect together.

This kind of spectroscopy is based on the fact that some atomic nuclei behave like tiny magnets, as though they had 鈥渘orth鈥� and 鈥渟outh鈥� poles. When placed in a magnetic field, nuclear energy levels which previously had the same energy are split apart, and transitions between them become possible. These levels correspond to different alignments of the subatomic nuclear magnets in the magnetic field, and the nuclei can be made to flip between these different alignments by absorbing radiation at specific wavelengths in the radiowave region.

The precise frequencies at which these 鈥渕agnetic resonance鈥� transitions occur will depend, as before, on the size of the energy gap between the levels involved (remember, 鈻礒 = hv). This, in turn, depends on the extent to which the magnetic nuclei are 鈥渟creened鈥� from the effects of the magnetic field by neighbouring chemical groupings, and hence on the exact chemical environments of the nuclei. Although these are energy levels of nuclei tucked away inside the atoms that make up the molecule, they are nevertheless subtly affected by aspects of the structure and bonding of the molecule as a whole.

The most common form of NMR spectroscopy is based on the hydrogen nucleus, ¹H (which is just a proton), but other magnetic nuclei such as ¹³C and ¹⁹F can be used because they too have nuclei that can be made to flip in a magnetic field. The extent to which the ¹H nucleus is screened from neighbouring atoms determines its position in a ¹H-NMR spectrum. Thus, ¹H nuclei attached to different chemical groups (鈭扖H₃, 鈭扖0₂H and so forth) require different magnetic field strengths to cause their nuclear spins to flip when bathed in radio wave radiation at a fixed frequency. Peaks thus appear in correspondingly different positions in the NMR spectrum. Furthermore, subtle interactions with other magnetic nuclei present in the same molecule can give rise to complex splitting of the peaks in the spectrum.

Trained spectroscopists can often work backwards from the splitting patterns to deduce the structure of the molecules.

Many modern NMR spectrometers work in a different way, although the basic principles are the same. Instead of using radiowaves at a fixed frequency and sweeping the magnetic field, modern spectrometers fire a short burst of highpower radio wave radiation at the sample. After this, the nuclei flip back towards an equilibrium population, emitting radiation detectable as a function of time. The resulting complex, oscillating signal contains the whole NMR spectrum which can be obtained from a standard mathematical transformation using a computer.

Molecules light up

Emission spectroscopy

WE HAVE so far focused on techniques which involve the absorption of radiation, essentially looking for moves up the energy ladder. Sometimes, emission of radiation from molecules is just as revealing, indicating leaps down the ladder. Once in some excited level or other, a molecule may lose its excess energy in one or more ways, sometimes by emitting radiation. All the same quantum rules apply. By dispersing the emitted radiation, emission spectra can be produced which add further details of a molecule鈥檚 energy level map.

Emission spectroscopy is, in principle, possible over as much of the electromagnetic spectrum as we have considered here for absorption. In practice, however, it is confined mainly to emission of visible and ultraviolet radiation. For reasons beyond the scope of this supplement, the emission spectrum of a molecule in this region does not match that molecule鈥檚 visible/ultraviolet absorption spectrum. Similar energy levels are involved, but the final energy levels in absorption transitions are not always the starting energy levels for the reverse emission transitions. The two forms of spectroscopy are complementary, often with clues available in an emission spectrum that may be absent from an absorption spectrum.

The ways that scientists use spectroscopy at many of the frontiers of modern chemistry and physics are different only in the sophistication that comes from decades of research effort. With sophisticated techniques for cooling molecules to very low temperatures, for example, spectroscopists can drag the molecules down to their lowest energy levels, greatly simplifying otherwise complex spectra. Combined with lasers able to deliver radiation with very well-defined wavelengths, or narrow pulses of light lasting millionths of a billionth of a second, the result is a precision spectrometer. Chemists are using these techniques to study a variety of unstable, reactive molecules and to probe the very act of chemical reaction itself.

Electromagnetic radiation

ABOUT fifty years after Fraunhofer鈥檚 discovery, the British physicist James Clerk Maxwell capitalised on the remarkable experimental and theoretical work of Michael Faraday to devise a theory of electromagnetic radiation. Maxwell described radiation as two transverse waves 鈥� essentially, waves which move up and down at right angles, both to each other and to their direction of travel. One set of waves represents an oscillating electric field, and the other an oscillating magnetic field. The combination is an electromagnetic wave.

The fundamental characteristics of a wave include its wavelength 鈥� the distance the wave travels to complete one up-and-down cycle, denoted as 位; its frequency 鈥� the number of up and down cycles in a given time interval, denoted as v; and its speed. In a vacuum, electromagnetic radiation moves at only one speed 鈥� the speed of light, c (2.9979248 脳 10⁸ metres per second), and v = c/位. The wavelengths and frequencies of electromagnetic radiation vary considerably, from the radio region, through microwaves and infrared to visible light, ultraviolet light and beyond to X-rays and gamma rays.

Although applied routinely today, Maxwell鈥檚 theory enjoyed fame as a complete theory of electromagnetic radiation for only a relatively short time. In 1990, the German physicist Max Planck introduced a new concept into physics which was adapted to very great effect by Albert Einstein. This was the concept of the light quantum. We now understand that radiation can be described in two ways: either in terms of waves, or in terms of light quanta, discrete particles or 鈥減ackets鈥� of energy which we call photons. Which kind of behaviour we see depends on what kind of experiment we select. The connection between these two apparently quite contradictory types of behaviouor is expressed in a mathematical relationship first discovered by Planck. This connects the frequency of the radiation, v with E, the energy of its photons, whereby E = hv, and h is Planck鈥檚 constant (6.6260755 脳 10^鈭�34 Joule-seconds). The higher the frequency of the radiation, the higher the energy of its photoons. Radio waves therefore consist of very low energy photons, whereas X-rays and gamma rays consist of very high energy photons.

Spectroscopists therefore have a very broad 鈥減alette鈥� of radiation that they can use to study molecules, from the low energy microwaves through to the high energy X-rays (see Figure). Molecules respond very differently depending on the wavelength of the radiation they absorb. The changes are so diverse that regions of the spectrum have spawned distinct areas of spectroscopy. The ones that this supplement has focused on are infrared spectroscopy, ultraviolet/visible spectroscopy and nuclear magnentic resonance spectroscopy.