CHAPTER 1
What is Flame Spectrometric Analysis?
1 The Nature of Light and Heat
Light and heat are both forms of energy, and as such can interact with matter to produce physical changes in the matter. Sometimes the nature of such interactions is obvious, or even dramatic, to onlookers; for example, many ceramic materials may shatter violently if heated rapidly by a flame. Often the effects are more subtle, such as when light from a burglar's torch activates an alarm or when light reflected from objects produces changes within the eye which the brain ultimately interprets as colour and brightness.
Frequently energy is transformed from one form to another. Use of a telephone involves a large number of such transformations, from conversion of chemical potential energy stored in the body, via movement, to sound waves, which in turn pass kinetic energy (energy associated with movement) on via the mouthpiece to produce electrical signals which pass along a wire, and so on. Although they are complex, because of their familiarity we tend to take such energy transfers and transformations for granted.
2 Interaction of Light and Heat with Atoms and Molecules
When visible light energy interacts with coloured molecules, or in other words when it is absorbed, this physical change takes the form of redistribution of some of the electrons of the molecule to different electronic orbitals. In this way the molecule temporarily stores the absorbed energy. We say it is 'electronically excited'. Atoms may become 'electronically excited' in a similar fashion, although this involves the absorption of ultraviolet (UV) light rather than visible light in most instances. Similarly, if molecules or atoms are subjected to high temperature environments, some of the heat energy may likewise be converted to electronic excitation energy.
The electronically excited molecules and atoms generally are not very stable, and soon lose their extra energy. They mostly do this by passing the energy on to other molecules by interaction with vibrating and rotating atoms bound within those molecules. For molecules in solution, for example, the energy transfer is mainly to solvent molecules, but such collisional deactivation will also occur in systems in the gaseous phase. Some molecules or atoms may be able to get rid of some of their extra energy by the emission of light energy, known as luminescence. The latter takes two forms, depending upon whether metastable intermediate electronically excited species are produced. If they are, the emission of light, then known as 'phosphorescence', is delayed by anything up to a few seconds. 'Fluorescence' emission, in the absence of metastable intermediate states, is much more rapid, generally occurring within a small fraction of a microsecond.
3 A Simple Quantitative Description of Light
In many situations, the behaviour of light can be best explained if the light is treated as an energetic wave. The wavelength of the light, λ is a measure of its energy. Expressed mathematically, the energy (E) is proportional to the speed of light (c) divided by the wavelength (λ), or E = hc/λ where h is a constant known as Planck's constant. Light at the red end of the spectrum has much less energy available than light at the violet end of the spectrum. Thus less energy is required to thermally excite atoms to electronically excited states which are likely to lead to the emission of red light than is required to excite them to states likely to lead to the emission of violet or UV light. The significance of this will become clearer in Chapter 2, when the instrumental requirements of flame emission spectrometry (FES) are discussed.
Visible light represents only a very small part of a much broader spectrum of energies. At wavelengths below the violet (i.e. with higher energies) come the ultraviolet (UV), and then X-rays and gamma rays. At wavelengths longer than red light come infrared (IR), and then radiowaves and microwaves. Only the UV–visible region is of interest in the present context, because this is the light which causes the physical changes in atoms or molecules which are exploited in analytical flame spectrometry.
4 The Absorption of Light for Quantitative Measurement
For coloured solutions, it is intuitively obvious that there is some sort of relationship between concentration of a soluble coloured species and the colour of the solution; the darker your tea or coffee, the stronger it is. Such solutions appear coloured because they absorb visible light. If they did not, they would be transparent. If a solution looks green, it must be transparent to green light. It may also transmit some blue and yellow light. It must, however, be selectively absorbing orange, red, and purple light. Measurement of the amount of red, orange, or purple light absorbed can provide a means of measuring the concentration of the absorbing species.
The human eye is not normally sensitive to UV light, and therefore notices nothing if UV light only is being absorbed. Atoms of many elements absorb light mainly or only in the UV region of the spectrum, and this region of the spectrum is widely exploited in atomic absorption spectrometry (AAS). At high concentrations especially, absorption of visible light may become significant for some elements. For example, the spectrum of our sun shows a number of dark absorption bands, where the continuum emitted from the high temperature solar surface is selectively absorbed by free atoms of elements such as sodium present in the solar atmosphere. These dark lines, the Fraunhoffer lines, are perhaps the oldest and best-known example of atomic absorption.
5 The Difference between Atomic and Molecular Absorption
It is appropriate at this point to consider the difference between the absorption of light by atoms and by molecules. Figure 1 compares the absorption spectrum of an element such as zinc with that of a simple triatomic gas, sulfur dioxide. The absorption spectra show how absorbance (a term explained more fully in section 1 which indicates how strongly light is being absorbed) varies with wavelength. Any particular electronic transition in an atom or molecule requires photons with an appropriate amount of energy to bring about a transition from a lower discrete (quantized) energy state to a higher quantized energy state. If the photons do not have enough energy, in other words, if the wavelength of the light is too long, the transition cannot occur. For atoms, the transition cannot occur if the wavelength is too short, either, for there is no mechanism by which the excess energy may be absorbed. Atomic absorption spectra therefore consist of isolated, very narrow bands, or lines, with one line for each possible electronic transition. This explains why the atomic absorption bands of sodium in the sun's atmosphere are sharp lines.
In the case of molecular absorption, the excess energy may be transferred into kinetic energy in the form of vibrations or rotations of atoms which constitute the molecules. Each particular rotational or vibrational excitation process requires a particular energy also, the energy needed being much less than that of an electronic transition. Thus molecular absorption spectra of gases consist of series of absorbing bands lying very close together. For molecules in solution, the situation is even more complicated because of interactions between the rotationally and vibrationally excited absorbing molecules and solvent molecules. This has the effect of smoothing out the spectra into broad bands, often covering 100 nm or more. The fine structure, such as that of sulfur dioxide in Figure 1, would not therefore be seen in a molecular absorption spectrum in solution at room temperature.
For exactly the same reasons, atomic emission spectra of elements consist of sharp lines, whereas molecular emission spectra consist of series of bands in the gas phase, or, in the solution phase, broad bands often with little or no obvious fine structure.
6 Quantitative AAS – What Should We Measure?
The wavelengths at which absorption or emission spectral peaks occur are characteristics of the particular atom or molecule giving rise to the peaks, and thus may be used for qualitative identification. In quantitative instrumental methods of analytical chemistry, we try to measure some property of atoms or molecules which varies linearly with the concentration of the species of interest. What parameter should we measure if we wish to exploit atomic absorption?
Consider a number of photons (I0) in a narrow, monochromatic beam passing through a cloud of (n) atoms: If some photons are absorbed by the atoms in the cloud, the number of transmitted photons (It) will be less than I0. Thus:
It = xI0, where 1 > x > 0 (1)
Consider now what happens if the concentration of atoms in the atom cloud is doubled. Suppose the probability of a photon being absorbed is independent of the number of photons, and depends only upon the number of atoms in the beam path. I0 – It photons will still be absorbed by the first n atoms, so we need to consider how the additional n atoms will interact with the xI0 photons which were not absorbed by the first n atoms. Thus for the second n atoms, once again a fraction 'x' of these xI0 photons will be absorbed. Thus when 'n' is increased to '2n', It is then given by x2I0. Similarly, for n, 2n, 3n 4n ... atoms, It would have the values xI0, x2I0, x3I0, x4I0. Thus the relationship between It, I0, and the atom concentration, c. is of the form:
It = xcI0 (2)
or: log It = c log x + log I0 (3)
or: log It = log I0 = kc (4)
or: log (It/I0) is proportional to c (5)
Since It< I0, It/I0< 1, and log (It/I0) < 0, if we define a parameter A, the absorbance, as:
A = - log (It/I0) (6)
the absorbance, A, will always be positive and proportional to concentration. Thus absorbance is the parameter which should be measured if straight line calibration plots are deemed desirable. Note that if 90% of the light is absorbed, It/I0 is equal to 10/100 or 0.1, and A is equal to 1. Similarly 99% absorption corresponds to an absorbance of 2, and 99.9% to an absorbance of 3, and so on. If you think about it, it should become apparent that precise measurement of absorbance values > 2 is likely to be difficult in practice.
7 The Sensitivity Problem in AAS
Although both the concept of absorbance and the nature of atomic absorption spectra had been understood for many decades by spectrophysicists by the early 1950s, atomic absorption had not been applied in quantitative analytical spectrometry at that time. The main limitation appeared to be the narrowness of atomic absorption lines. Monochromators could be used relatively easily (see Chapter 2) to provide a window to isolate bands of the UV–visible spectrum about 0.1 nm in width, while atomic absorption occurred over a much narrower spectral interval, typically 0.005 nm or less. Consider the situation if quite strong atomic absorbance occurred as a consequence of light passing through a cell such as a flame containing free atoms of the element of interest; there would be no change in 95% of the light passing through the monochromator 'window'. This situation is represented in Figure 2. Thus It would be only very slightly smaller than I0, the ratio It/ I0 would be close to 1, and, from equation 6, absorbance would be close to zero. Sensitivity would therefore always be very poor.
Sensitivity could, of course, be improved significantly by using a very high resolution monochromator, which could isolate regions of the spectrum of 0.005 nm or less. At that time, however, such monochromators were large and expensive, had a very low light throughput (which could result in signal stability problems), and sometimes required fairly precise temperature control to give good wavelength stability. Thus they were hardly ideal for routine use in busy and often congested laboratories.
A major breakthrough came in Australia when Alan Walsh realized that light sources were available for many elements which emitted atomic spectral lines at the same wavelengths as those at which absorption occurred. By selecting appropriate sources, the emission line widths could be even narrower than the absorption line widths (Figure 2). Thus the sensitivity problem was solved more or less at a stroke, and the modern flame atomic absorption spectrometer was born.
8 The Potential Selectivity of Flame AAS
The importance of Walsh's ideas should not be underestimated. Not only had he suggested a potentially highly sensitive method of analysis which would prove eventually to be suitable for the determination of many elements in the periodic table, but at the same time he had suggested a development which, theoretically at least, should lead to virtually specific analysis. The very narrowness of the absorption lines which had hitherto held back progress in AAS suddenly became its most powerful asset. It meant that the chances of spectral overlap of the absorption line of one element with the emission line of another were extremely small. Thus atomic spectral interferences should be, and indeed are, rare in AAS.
It is little wonder, then, that analysts working in environmental laboratories were amongst the first to grasp the importance of Walsh's benchmark papers. Hitherto the complex nature of their typical sample matrix and the nature of the elements they were interested in often resulted in the need for complex and time consuming separation and preconcentration by skilled chemists prior to the determination step. Suddenly they were apparently being offered a very sensitive and almost universally applicable technique where the only sample preparation needed seemed to be sample dissolution. Eventually problems did start to surface, as is invariably the case, but as Chapter 3 will show, few have proved insuperable.
9 Flame Emission Spectrometry
So far in this chapter, absorption techniques have been considered in preference to emission techniques, in spite of the much longer history of flame emission spectrometry (FES). This is deliberate, and reflects the far greater relative importance of AAS as a routine technique in most modern environmental analytical laboratories. Flame atomic fluorescence spectrometry (AFS) evolved a decade after flame AAS, but is still of only minor importance. A similar balance will be seen in the following chapters of this book, since many FES and AFS determinations are performed utilizing atomic absorption spectrometers.
The introduction of the first flame spectroscopes early in the second half of the 19th century was as impressive an achievement in its time as that of AAS a century later. Elements were introduced into a flame in the form of a salt and the spectrum excited in the flame was examined using a prism spectrometer. Using the eye as a detector, characteristic emission wavelengths of elements could be measured, and subsequently used as a basis for qualitative analysis. In the flames available at the time, only those elements such as K, Li, Na, and Ca which were easily thermally excited (elements with low excitation potentials) and which therefore emitted visible light were studied. However the technique was successfully applied to establishing the existence of other elements falling into this category, such as Cs, In, Rb, and Tl.
The potential development of FES as a quantitative technique was hampered until well into the 20th century by the use of the eye as a detector; this limited the reliability of quantification by the comparison of light intensity from samples and standard materials. The evolution of photocell detectors enhanced sensitivity and improved reliability, considerably extending the range of application of FES techniques.
The emission intensity depends upon the number of excited atoms, N*. which in turn depends upon the number of unexcited or ground-state atoms, No, the statistical probabilities of excitation and emission occurring, and exponentially upon the flame temperature, T, and the reciprocal of the excitation potential, E. Thus:
N* is proportional to No exp (-[E/kT)) (7)
where k is Boltzmann's constant, High emission intensity is favoured, as might be intuitively expected, by high flame temperature and when the excitation potential is low.
