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Acidity, Basicity, and Pka 8 Connections

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Acidity, Basicity, and Pka 8 Connections Acidity, basicity, and pKa 8 Connections Building on: Arriving at: Looking forward to: • Conjugation and molecular stability • Why some molecules are acidic and • Acid and base catalysis in carbonyl ch7 others basic reactions ch12 & ch14 • Curly arrows represent delocalization • Why some acids are strong and others • The role of catalysts in organic and mechanisms ch5 weak mechanisms ch13 • How orbitals overlap to form • Why some bases are strong and others • Making reactions selective using conjugated systems ch4 weak acids and bases ch24 • Estimating acidity and basicity using pH and pKa • Structure and equilibria in proton- transfer reactions • Which protons in more complex molecules are more acidic • Which lone pairs in more complex molecules are more basic • Quantitative acid/base ideas affecting reactions and solubility • Effects of quantitative acid/base ideas on medicine design Note from the authors to all readers This chapter contains physical data and mathematical material that some readers may find daunting. Organic chemistry students come from many different backgrounds since organic chemistry occu- pies a middle ground between the physical and the biological sciences. We hope that those from a more physical background will enjoy the material as it is. If you are one of those, you should work your way through the entire chapter. If you come from a more biological background, especially if you have done little maths at school, you may lose the essence of the chapter in a struggle to under- stand the equations. We have therefore picked out the more mathematical parts in boxes and you should abandon these parts if you find them too alien. We consider the general principles behind the chapter so important that we are not prepared to omit this essential material but you should try to grasp the principles without worrying too much about the equations. The ideas of acidity, basicity, and pKa values together with an approximate quantitative feel for the strength and weakness of acids and bases are at least as central for biochemistry as they are for organic chemistry. Please do not be discouraged but enjoy the challenge. Introduction This chapter is all about acidity, basicity, and pKa. Acids and bases are obviously important because many organic and biological reactions are catalysed by acids or bases, but what is pKa and what use is it? pKa tells us how acidic (or not) a given hydrogen atom in a compound is. This is useful because, if the first step in a reaction is the protonation or deprotonation of one of the reactants, it is obviously necessary to know where the compound would be protonated or deprotonated and what strength acid or base would be needed. It would be futile to use too weak a base to deprotonate a compound but, equally, using a very strong base where a weak one would do would be like trying to 182 8 . Acidity, basicity, and pKa crack open a walnut using a sledge hammer—you would succeed but your nut would be totally destroyed in the process. The aim of this chapter is to help you to understand why a given compound has the pKa that it does. Once you understand the trends involved, you should have a good feel for the pKa values of commonly encountered compounds and also be able to predict the values for unfamiliar com- pounds. Originally, a substance was identified as an acid if it exhibited the properties shown by other acids: a sour taste (the word acid is derived from the Latin acidus meaning ‘sour’) and the abilities to turn blue vegetable dyes red, to dissolve chalk with the evolution of gas, and to react with certain ‘bases’ to form salts. It seemed that all acids must therefore contain something in common and at the end of the eighteenth century, the French chemist Lavoisier erroneously proclaimed this common agent to be oxygen (indeed, he named oxygen from the Greek oxus ‘acid’ and gennao ‘I produce’). Later it was realized that some acids, for example, hydrochloric acid, did not contain oxygen and soon hydrogen was identified as the key species. However, not all hydrogen-containing compounds are acidic, and at the end of the nineteenth century it was understood that such compounds are acidic only if they produce hydrogen ions H+ in aqueous solution—the more acidic the compound, the more hydro- gen ions it produces. This was refined once more in 1923 by J.N. Brønsted who proposed simple def- í initions for acids and bases. Other definitions of acids and bases are useful, the most notable being those of Lewis, also proposed in 1923. Brønsted definitions of acids and bases However, for this chapter, the Brønsted • definition is entirely adequate. • An acid is a species having a tendency to lose a proton • A base is a species having a tendency to accept a proton Acidity + An isolated proton is incredibly reactive—formation of H3O in water Hydrochloric acid is a strong acid: the free energy ∆G° for its ionization equilibrium in water is –40 kJ mol–1. + – ∆ -1 HCl (aq) H (aq) + Cl (aq) G°298K = –40 kJ mol Such a large negative ∆G° value means that the equilibrium lies well over to the right. In the gas phase, however, things are drastically different and ∆G° for the ionization is +1347 kJ mol–1. + – HCl (g) H (g) + Cl (g) ∆G°298K = +1347 kJ mol-1 This ∆G° value corresponds to 1 molecule of HCl in 10240 being dissociated! This means that HCl does not spontaneously ionize in the gas phase—it does not lose protons at all. Why then is HCl such a strong acid in water? The key to this problem is, of course, the water. In the gas phase we would have to form an isolated proton (H+, hydrogen ion) and chloride ion and this is energetically very unfavourable. In contrast, in aqueous solution the proton is strongly + attached to a water molecule to give the very stable hydronium ion, H3O , and the ions are no longer isolated but solvated. Even in the gas phase, adding an extra proton to neutral water is highly exothermic. + + ∆ –1 H2O (g) + H (g) H3O (g) H ° = – 686 kJ mol O H H H In fact, an isolated proton is so reactive that it will even add on to a molecule of methane in the gas H H O + O phase to give CH5 in a strongly exothermic reaction (you have already encountered this species in H H mass spectrometry on p. 52). We are therefore extremely unlikely to have a naked proton in the gas O H H phase and certainly never in solution. In aqueous solution a proton will be attached to a water mole- cule to give a hydronium ion, H O+ (sometimes called a hydroxonium ion). This will be solvated a structure for a solvated 3 hydronium ion in water just as any other cation (or anion) would be and hydrogen bonding gives rise to such exotic species as the dashed bonds represent H O+ (H O+·3H O) shown here. hydrogen bonds 9 4 3 2 Acidity 183 Every acid has a conjugate base In water, hydrogen chloride donates a proton to a water molecule to give a hydronium ion and chlo- ride ion, both of which are strongly solvated. + – HCl (aq) + H2O (l) H3O (aq) + Cl (aq) In this reaction water is acting as a base, according to our definition above, by accepting a proton from HCl which in turn is acting as an acid by donating a proton. If we consider the reverse reaction (which is admittedly insignificant in this case since the equilibrium lies well over to the right), the chloride ion accepts a proton from the hydronium ion. Now the chloride is acting as a base and the hydronium ion as an acid. The chloride ion is called the conjugate base of hydrochloric acid and the + hydronium ion, H3O , is the conjugate acid of water. •For any acid and any base AH + B BH+ + A– where AH is an acid and A– is its conjugate base and B is a base and BH+ is its conjugate acid, that is, every acid has a conjugate base associated with it and every base has a conjugate acid associated with it. For example, with ammonia and acetic acid + – CH3COOH + NH3 NH4 + CH3COO + the ammonium ion, NH4 , is the conjugate acid of the base ammonia, NH3, and the acetate ion, – CH3COO , is the conjugate base of acetic acid, CH3COOH. Water can behave as an acid or as a base If a strong acid is added to water, the water acts as a base and is protonated by the acid to become + H3O . If we added a strong base to water, the base would deprotonate the water to give hydroxide ion, OH–, and here the water would be acting as an acid. Such compounds that can act as either an acid or a base are called amphoteric. With a strong enough acid, we can protonate almost anything and, O OH likewise, with a strong enough base we can deprotonate almost anything. + HCl + Cl This means that, to a certain degree, all compounds are amphoteric. For Me OH Me OH example, hydrochloric acid will protonate acetic acid. In this example acetic acid is acting as a base! Other compounds need acids even stronger than HCl to protonate them. Remember that, in chemical ionization mass spectrometry (p. 52), proto- + nated methane, CH5 , was used to protonate whatever sample we put in to the machine in order to + give us a cation; CH5 is an incredibly strong acid. basic group acidic group The amino acids you encountered in Chapter 2 are amphoteric.
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