Education in Chemistry - January 2004 - A brief history of acidity

Education in Chemistry - January 2004 - A brief history of acidity

John W. Nicholson

An excursion through the theories of acids and bases, from 13th to 21st centuries, can enhance students' understanding of the nature of acids.

Acids and bases have been recognised since antiquity.1,2 The alchemical works of Raymund Lully (ca 1232-1316), for example, describe the preparation of nitric acid and aqua regia, the 'water of kings'because it dissolved gold.3 The term 'acid', originates from the Latin acetum, meaning sour. Tasting was the earliest and most readily available chemical test and so, not surprisingly, the term acid was used to describe sour-tasting substances. With time, chemists realised that acids had other properties in common - for example, they changed the colour of the plant dye litmus from blue to red, they corroded metals and they precipitated sulphur from alkaline solution.

Alkalies, too, have been known for centuries. The name, alkali, comes from the Arabic word for plant ash, al kalja.1 The term was originally used for the water-soluble extracts obtained from the ash of certain plants. Early chemists were quick to appreciate that alkalies also had distinctive properties: they were slippery, they removed fats and oils from fabrics, and they turned litmus from red to blue.4 At first the term was used to describe hydroxides and carbonates of sodium and potassium. However, by 1744, Ruoelle had extended the concept to include the alkaline earth analogues and coined the term base for them, though the reason for this choice of word is not clear.4

Early classifications

The importance of salt formation in classifying acids and bases can be traced back to Johann Rudolph Glauber in 1648,5 who stated that acids and alkalies were, in some sense, opposites and that salts were formed by reacting these two components. This led to the ideas that salts can be broken down into an acid and an alkali, and that salts are characterised by having different properties from those of the parent acid or alkali. Many distinguished scientists, however, did not agree with these notions, notably Robert Boyle and Georg Stahl. Boyle went so far as to postulate a physical feature to account for the sour taste of acids, suggesting that they were composed of sharp angular particles. He reasoned that the loss of the sour taste on salt formation was because the acid particles became embedded in soft, rounded particles of alkali.5

The question of the composition of acids was taken up by Lavoisier. In 1777, he proposed that oxygen was the 'universal acidifying principle',5,6 and that an acid should be defined as a compound of oxygen with a non-metal. He classified substances such as silica, SiO2, as acids, a distinction still recognised today among geologists, who refer to rocks with high silica contents as being acidic. Indeed, the name Lavoisier gave to the newly discovered gas oxygen (or oxygene) means 'acid-former'.

In the early 19th century, the English chemist, Humphry Davy (1778-1829) dismissed Lavoisier's theory. Working at the Royal Institution in London, Davy was interested in what was then known as muriatic acid (in modern terms, HCl). He also studied 'oxymuriatic acid', the green gas that was obtained by oxidising muriatic acid. In his experiments, Davy exposed oxymuriatic acid to white-hot carbon in an attempt to remove it as CO2.7 When he was unable to do this, he concluded that oxymuriatic acid was in fact an element, which he called chlorine, from the Greek word chloros, meaning pale green.2 Davy reasoned that muriatic acid also contained no oxygen but was a binary compound of hydrogen and the new element, chlorine.

As for acidity, Davy was convinced that in some vague way, this was a function of the arrangement of the elements present, rather than the presence of a particular element. It took agricultural chemist Justus von Liebig (1803-73) to acknowledge the role of hydrogen in an acid when, in 1838, he defined organic acids as 'compounds containing hydrogen in which the hydrogen can be replaced by a metal'.8

A chemical milestone

Over 40 years later Svante Arrhenius, in 1884, suggested, as part of his doctoral thesis, that certain substances become ionised in solution. The idea was not well received, and his thesis was awarded a fourth class degree. By 1903, opinion had so changed that he won the third Nobel prize in chemistry for this work. Even by that time his critics had not entirely disappeared, however. Writing a few years after the award of the Nobel prize, the redoubtable Henry Armstrong noted in an obituary of Raphael Meldola that ' ... Meldola never speculated in ions and knew what was required of the chemist ...'.9

By 1887, Arrhenius had extended his ideas, suggesting that the characteristic properties of acids might be explained in terms of dissociation. He proposed the definition of an acid as a substance that dissociates in water into hydrogen ions and anions.4 He defined a base as a substance that dissociated in water into hydroxyl ions and cations. For the first time, a base was considered on its own terms, and not simply as a vague opposite of an acid. Arrhenius went on to say that hydrogen ions and the hydroxyl ions were able to form further water molecules, and the remaining dissociated anions and cations were the components of the newly formed water-soluble salt.4

Although Arrhenius' theory marked a milestone in chemistry - for the first time acids and bases were defined in terms other than by the reaction between them - it was limited to water-soluble substances and to aqueous systems. Arrhenius had thus failed to recognise the basicity of insoluble substances, or of oxides or silicates. Nonetheless, his theory paved the way for what is still regarded as the definitive explanation of acids and bases in aqueous systems, ie the Brønsted-Lowry theory.

Independently, in 1923, in two separate publications, Johannes N. Brønsted of Copenhagen10 and T. Martin Lowry of the University of Cambridge11 published the definition of an acid as '... a substance which gives up a proton ...' (which is similar to that of Arrhenius' definition), and that of a base, as '... a substance capable of accepting protons'. In this definition of a base, the relationship between an acid and a base is a reciprocal one, and an acid-base reaction involves the transfer of a proton from the acid to the base. This removes both the restriction to aqueous solutions and the necessity for ionisation, the two major limitations to Arrhenius' view of acids and bases.

In fact, though Lowry's contribution was similar to Brønsted's, he failed to recognise the conjugate relationships existing in acid-base systems. He therefore failed to recognise NH4+ as an acid or CH3CO2- as a base.8 Lowry himself attributed this insight to Brønsted in a paper published in 1927.12

However, in Lowry's favour, he did make an important observation about the state of hydrogen in solution. Brønsted had simply assumed it was H+, whereas Lowry proposed the form that is commonly used today (ie H3O+), though he actually wrote it as OH3+. He clearly understood the driving force for this hydration, as is shown by the following quotation:11

It is a remarkable fact that strong acidity is apparently developed only in mixtures and never in pure compounds. Even hydrogen chloride only becomes acidic when mixed with water. This can be explained by the extreme reluctance of the hydrogen nucleus to lead an isolated existence.

The concept of pH

Once the crucial role of hydrogen in acidity had been established, the way became clear for the development of a quantitative scale of acid strength based on the concentration of hydrogen. This is the basis of the pH scale (pH = -log10[H+]) introduced by Sören Sörensen (1868-1939) in 1909.

Sören Sörensen
Sörensen was born at Havreber, Denmark, on 9 January 1868, the son of a farmer. In 1886 he matriculated at the University of Copenhagen, and studied chemistry under S. M. Jorgensen, the inorganic chemist. In 1899, he submitted his doctorate for work on cobalt oxalates. At the early age of 33, he was appointed director of the chemical department of the Carlsberg Laboratory. In this role, he became interested in the behaviour of proteins in solution, and effectively a pioneer in the field of colloid chemistry. However, the need to supervise analyses, as well as his early training in preparative chemistry, meant he was meticulous in his work, and willing when necessary to explore beyond the immediate field of colloid chemistry. This explains his interest in the properties of acid solutions, and his development of the pH scale as a simple means of measuring their acidity.

Later in his career, his wife Margrethe worked with him at the Carlsberg Laboratory, studying interactions of carbon monoxide with haemoglobin. Sörensen was well known in his lifetime, and numerous honours were conferred on him, of which the most important was the presidency of the Danish National Academy of Sciences. He died at the age of 71 in 1939.

Sörensen became interested in this problem when he was director of the chemical department of the Carlsberg Laboratory in Copenhagen (see Box). He was studying the behaviour of proteins in aqueous solution. Since this behaviour is strongly influenced by the acidity of their environment, Sörensen first considered the means by which this environment could be described quantitatively.13 In 1909, he published a paper entitled Studies on enzymes, II, in which he discussed hydrogen ion concentration in depth.14

Sörensen began by observing that the rate at which enzymic splitting of proteins occurs depends, among other things, on the degree of acidity or alkalinity of the solution. He recognised that the total concentration of acid was not the important factor, but the extent to which that acid was dissociated. However, he applied this idea to 0.1 normal HCl, which he suggested was not fully dissociated. He thus assumed the hydrogen ion concentration to be '... somewhat less than 0.1 normal'.14 He went on to say:14

The value of the hydrogen ion concentration will accordingly be expressed by the hydrogen ion based on the normality factor of the solution used, and this factor will have the form of a negative power of 10. Since in the following section I usually refer to this, I will explain here that I use the name 'hydrogen ion exponent' and the designation PH for the numerical value of this power.

He notes that in a solution whose normality based on hydrogen ions equals 0.01, 'PH' can be written briefly as 10-2. Sörensen does not explain his notation any further, nor does he account for his choice of the letter 'P'. Others, though, have claimed that it comes from the German word potenz, meaning power or concentration.7

Sörensen was aware of the difference between concentration and activity, and in this first paper, explores in detail the relationship between them in solutions such as 0.1 M HCl and 0.01 M HCl with 0.09 M KCl. He also noted that the pH of neutral water was 7.07. Later he explored the small variation in pH of a solution caused by adding indicators, and he carried out numerous studies of the buffering ranges and effects of borates, citrates, phosphates and glycine (aminoethanoic acid).13

Modern views

The modern formulation of the equation defining pH is:15

pH = - log aH3O+ This builds on Lowry's recognition of the activity of the hydronium ion rather than of the hydrogen ion as the key to pH.15

The simple description of the solvated proton in terms of the hydronium ion, H3O+, is not fully satisfactory, because there are several ionic species that occur in acid solutions. The hydronium ion has been shown to exist using mass spectroscopy,16 but other ions of the type H+(H2O)n with values of n up to eight have also been shown to be stable by using this technique16 and calculations have demonstrated that values of n of three or four are stable.17 In reality, any of these ions has only a fleeting lifetime because the proton is very mobile and transfers quickly and easily among neighbouring clusters of water molecules. This makes it impossible to be sure of the relative ratios of the various possible hydrated proton species in water. Thus, provided the concept is used with care, the activity of the hydronium ion remains useful as the basis for a modern understanding of the concept of pH.

Our knowledge of acidity now extends well beyond simple aqueous systems. Gilbert Lewis of the University of California, Berkeley, for example, redefined an acid as a substance that can use a lone pair of electrons from another molecule to complete the stable shell of one of its own atoms, ie as an electron pair acceptor. Typical examples are metal compounds such as SnCl4 and ZnCl2, which form donor-acceptor complexes of the type Sn(NH3)2Cl4 and ZnCl2·2H2O.

Acidity may also extend well beyond pH values obtainable with common mineral acids. In the 1960s, George Olah of Case Western Reserve University developed powerful examples of 'superacids' by mixing together various substances. Superacids are several orders of magnitude more acidic than conventional acids, with pH values as low as 10-18. One example is 'Magic Acid': a mixture of antimony pentafluoride (SbF5) and fluorosulphuric acid (HSO3F), which was named for its apparently magical ability to dissolve candle wax.

However, for practical everyday chemistry, it is aqueous systems that are of most interest. For these, we cannot improve on pH and its underlying concepts as the basis of our understanding of that most important of chemical properties, which has been known for centuries, of acidity.

Professor John Nicholson is head of school of science at the University of Greenwich, Chatham Maritime, Chatham, Kent ME4 4TB.

References

  1. H. L. Finston and A. C. Rychtman, A new view of current acid-base theories. New York: John Wiley, 1982.
  2. J. Hudson, The history of chemistry. London: Macmillan, 1992
  3. J. R. Partington, A short history of chemistry. London: Macmillan, 1948.
  4. M. C. Day and J. Selbin, Theoretical inorganic chemistry. New York: Reinhold, 1969.
  5. P. Walden, Salts, acids and bases: electrolytes, stereochemistry. New York: McGraw-Hill, 1929.
  6. M. P. Crosland, Isis, 1973, 64, 325.
  7. H. M. Leicester and H. S. Klickstein (eds), Sourcebook in chemistry, p243. Cambridge MA: Harvard University Press, 1968.
  8. R. P. Bell, The proton in chemistry. London: Methuen, 1959.
  9. H. E. Armstrong, Analyst, 1916, 41, 27.
  10. A NAME=10>J. N. Brønsted, Rec. Travaux chimiques des Pay-Bas et de la Belgique, 1923, 42, 718.
  11. T. M. Lowry, Chem. & Ind., 1923, 42, 43.
  12. T. M. Lowry, J. Chem. Soc., 1927, 2562.
  13. E. K. Rideal, Sören Peter Lauritz Sörensen, memorial lectures delivered before The Chemical Society 1933-1942, Vol IV, p159. London: The Chemical Society, 1951
  14. S. P. L. Sörensen, Biochem. Zeitschrift, 1909, 23, 131.
  15. P. Atkins and J. de Paula, Physical chemistry, 7th edn. Oxford: OUP, 2002.
  16. A. J. C. Cunningham, J. D. Payzant and P. Kebarle, J. Am. Chem. Soc., 1972, 94, 7627.
  17. E. Kochanski, J. Am. Chem. Soc., 1985, 107, 7869.

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