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where yl is the chemical potential of the solute, xx its mole fraction and x{) the mole fraction of the solvent. In the subsequent text, wherever possible, the quantities jU and pf will not be distinguished by separate symbols: only the symbol $ will be employed. In real mixtures and solutions, the chemical potential ceases to be a linear function of the logarithm of the partial pressure or mole fraction. Consequently, a different approach is usually adopted. The simple form of the equations derived for ideal systems is retained for real systems, but a different quantity a, called the activity (or fugacity for real gases), is



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introduced. Imagine that the dissolved species are less 'active' than would correspond to their concentration, as if some sort of loss' of the given interaction were involved. The activity is related to the chemical potential by the relationship [i^tf + RTlna, (1.1.8)





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In addition to these tools to aid your macro work with the G11, keep these ideas in mind no matter what equipment setup you work with: Use the Self-timer in lieu of a Remote Switch If you don t have a Remote Switch, use your camera s Self-timer to trip the shutter to avoid any camera shake when depressing the Shutter Release button with your nger Macro photography magni es any unintentional camera movement, so do everything you can to keep the camera rock steady Check the composition Remember to compose for the entire frame, not just the subject Macro photography is about more than just getting as close as you can to a subject Frame the subject, but be aware of any background elements that may be distracting.

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As in electrochemical investigations low pressures are usually employed, the analogy of activity for the gaseous state, the fugacity, will not be introduced in the present book. Electrolyte solutions differ from solutions of uncharged species in their greater tendency to behave non-ideally. This is a result of differences in the forces producing the deviation from ideality, i.e. the forces of interaction between particles of the dissolved substances. In non-electrolytes, these are short-range forces (non-bonding interaction forces); in electrolytes, these are electrostatic forces whose relatively greater range is given by Coulomb's law. Consider the process of concentrating both electrolyte and nonelectrolyte solutions. If the process starts with infinitely dilute solutions, then their initial behaviour will be ideal; with increasing concentration coulombic interactions and at still higher concentrations, van der Waals non-bonding interactions and dipole-dipole interactions will become important. Thus, non-ideal behaviour must be considered for electrolyte solutions at much lower concentrations than for non-electrolyte solutions. 'Respecting non-ideal behaviour' means replacing the mole fractions, molalities and molar concentrations by the corresponding activities in all the thermodynamic relationships. For example, in an aqueous solution with a molar concentration of 10~3 mol dm~3, sodium chloride has an activity of 0.967 x 10~3. Non-electrolyte solutions retain their ideal properties up to concentrations that may be as much as two orders of magnitude higher, as illustrated in Fig. 1.1. Thus, the deviation in the behaviour of electrolyte solutions from the ideal depends on the composition of the solution, and the activity of the components is a function of their mole fractions. For practical reasons, the form of this function has been defined in the simplest way possible: ax = yxx (1.1.9)

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where the quantity yx is termed the activity coefficient (the significance of the subscript x will be considered later). However, the complications connected with solution non-ideality have not been removed but only transferred to the activity coefficient, which is also a function of the concentration. The form of this function can be found either theoretically (the theory has been quite successful for electrolyte solutions, see Section 1.3) or empirically. Practical calculations can be based on one of the theoretical or semiempirical equations for the activity coefficient (for the simple ions, the activity coefficient values are tabulated); the activity coefficient is then multiplied by the concentration and the activity thus

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