Professor of Analytical Chemistry. He was a pioneer in introducing Atomic Spectrometry methods in Spain. He is a world-reputed Atomic Spectroscopist who introduces/develops many novel, state-of-the-art atomic techniques. He is a member of the editorial board of ABC and also of the editorial board of the RSC for a series of monographs on Atomic Spectrometry (leaded by Prof. Neil Burnett). He currently develops and tests cutting-edge instruments, even for some commercial firms. He published many papers and books. In 2007 he won the Robert Kellner award (The Robert Keller Lecture will be given at the Euroanalysis 2007).

Flame Spectrometry in Environmental Chemical Analysis: A Practical Guide

By Malcolm S. Cresser

The Royal Society of Chemistry

Copyright © 1994 The Royal Society of Chemistry
All rights reserved.
ISBN: 978-0-85186-734-2

Contents

Chapter 1 What is Flame Spectrometric Analysis?, 1,
Chapter 2 Instrumentation for Analytical Flame Spectrometry, 9,
Chapter 3 Interferences and How to Overcome Them, 31,
Chapter 4 Optimization in Flame Spectrometry, 43,
Chapter 5 Sample Preparation, 59,
Chapter 6 Some Useful Accessories, 69,
Chapter 7 Choice of Technique for Commonly Determined Elements, 79,
Chapter 8 How Do I Know I'm Getting the Right Answer?, 95,
Chapter 9 Safety in Flame Spectrometry, 99,
Subject Index, 103,

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...