KEY POINT The surfaces of many catalytic materials and minerals have Brứnsted and Lewis acid sites.
Some of the most important reactions involving the Lewis and Brứnsted acidity of inorganic compounds occur at solid surfaces. For example, surface acids, which are solids with a high surface area and Lewis acid sites, are used as cata- lysts in the petrochemical industry for the interconversion of hydrocarbons. The surfaces of many materials that are important in the chemistry of soil and natural waters also have Brứnsted and Lewis acid sites.
Silica surfaces do not readily produce Lewis acid sites because –OH groups remain tenaciously attached at the surface of SiO2 derivatives; as a result, Brứnsted acidity is dominant. The Brứnsted acidity of silica surfaces them- selves is only moderate (and comparable to that of acetic acid). However, as already remarked, aluminosilicates dis- play strong Brứnsted acidity. When surface OH groups are removed by heat treatment, the aluminosilicate surface possesses strong Lewis acid sites. The best-known class of aluminosilicates is the zeolites (Section 14.15) which are widely used as environmentally benign heterogeneous cata- lysts.The catalytic activity of zeolites arises from their acidic nature and they are known as solid acids. Other solid acids include supported heteropoly acids and acidic clays. Some reactions occurring at these catalysts are very sensitive to the presence of Brứnsted or Lewis acid sites. For example, toluene can be subjected to Friedel–Crafts alkylation over a bentonite clay catalyst:
CH2X
+ + HX
When the reagent is benzyl chloride Lewis acid sites are involved in the reaction and when the reagent is benzyl alco- hol Brứnsted sites are involved.
Surface reactions carried out using the Brứnsted acid sites of silica gels are used to prepare thin coatings of a wide variety of organic groups using surface modification reac- tions such as
OH Si
O O O
+ HOSiR3
OSiR3
Si O O O
+ H2O
OH Si
O O O
+ ClSiR3
OSiR3 Si
O O O
+ HCl
Thus, silica gel surfaces can be modified to have affinities for specific classes of molecules. This procedure greatly expands the range of stationary phases that can be used for chroma- tography. The surface –OH groups on glass can be modified similarly, and glassware treated in this manner is sometimes used in the laboratory when proton-sensitive compounds are being studied.
Solid acids are finding new applications in green chemis- try. Traditional industrial processes generate large volumes of hazardous waste during the final stages of the process when the product is separated from the reagents and by- products. Solid catalysts are easily separated from liquid products and reactions can often operate under milder con- ditions and give greater selectivity.
FURTHER READING
W. Stumm and J.J. Morgan, Aquatic chemistry: chemical equilibria and rates in natural waters. Wiley, New York (1995). The classic text on the chemistry of natural waters.
N. Corcoran, Chemistry in non-aqueous solvents. Kluwer Academic Publishers (2003). A comprehensive account.
J. Burgess, Ions in solution: basic principles of chemical interactions.
Ellis Horwood, Chichester (1999).
T. Akiyama, Stronger Brứnsted acids. Chem. Rev., 2007, 107, 5744.
G. Cavallo, P. Metrangolo, R. Milani, T. Pilati, A. Priimagi, G.
Resnati, and G. Terraneo, The halogen bond. Chem. Rev., 2016, 116, 2478.
G.-J. Zhao and K.-L. Han, Hydrogen bonding in the electronic excited state. Acc. Chem. Res., 2012, 45, 404.
S.J. Grabowski, What is the covalency of hydrogen bonding?
Chem. Rev., 2011, 111, 2597.
E.J. Corey, Enantioselective catalysis based on cationic oxazaborolidines. Angew. Chem. Int. Ed., 2009, 48, 2100.
D.W. Stephan, ‘Frustrated Lewis pairs’: a concept for new reactivity and catalysis. Org. Biomol. Chem., 2008, 6, 1535.
D.W. Stephan and G. Erker, Frustrated Lewis pairs: metal-free hydrogen activation and more. Angew. Chem. Int. Ed., 2010, 49, 46.
P. Raveendran, Y. Ikushima, and S.L. Wallen, Polar attributes of supercritical carbon dioxide. Acc. Chem. Res., 2005, 38, 478.
F. Jutz, J.-M. Andanson, and A. Baiker, Ionic liquids and dense carbon dioxide: a beneficial biphasic system for catalysis. Chem.
Rev. 2011, 111, 322.
D.R. MacFarlane, J.M. Pringle, K.M. Johansson, S.A. Forsyth, and M. Forsyth, Lewis base ionic liquids. Chem. Commun., 2006, 1905.
R. Sheldon, Catalytic reactions in ionic liquids. Chem. Commun., 2001, 2399.
I. Krossing, J.M. Slattery, C. Daguenet, P.J. Dyson, A. Oleinikova, and H. Weingọrtner, Why are ionic liquids liquid? A simple explanation based on lattice and solvation energies. J. Am.
Chem. Soc., 2006, 128, 13427.
G.A. Olah, G.K. Prakash, and J. Sommer, Superacids. Wiley, New York (1985).
R.J. Gillespie and J. Laing, Superacid solutions in hydrogen fluoride.
J. Am. Chem. Soc., 1988, 110, 6053.
E.S. Stoyanov, K.-C Kim, and C.A. Reed, A strong acid that does not protonate water. J. Phys. Chem. A, 2004, 108, 9310.
EXERCISES
5.1 Sketch an outline of the s and p blocks of the periodic table and indicate on it the elements that form (a) strongly acidic oxides and (b) strongly basic oxides, and (c) show the regions for which amphoterism is common.
5.2 Identify the conjugate bases corresponding to the following acids: [Co(NH3)5(OH2)]3+, HSO4−, CH3OH, H2PO4−, Si(OH)4, HS−. 5.3 Identify the conjugate acids of the bases C5H5N (pyridine), HPO42−, O2−, CH3COOH, [Co(CO)4]−, CN−.
5.4 Calculate the equilibrium concentration of H3O+ in a 0.10 M solution of butanoic acid (Ka = 1.86 × 10−5). What is the pH of this solution?
5.5 The Ka of ethanoic acid, CH3COOH, in water is 1.8 × 10−5. Calculate Kb of the conjugate base, CH3CO2−.
5.6 The value of Kb for pyridine, C5H5N, is 1.8 × 10−9. Calculate Ka for the conjugate acid, C5H5NH+.
5.7 The effective proton affinity Ap′ of F− in water is 1150 kJ mol−1. Predict whether it will behave as an acid or a base in water.
5.8 Draw the structures of chloric acid and chlorous acid and predict their pKa values using Pauling’s rules.
5.9 Aided by Fig. 5.5 (taking solvent levelling into account), identify which bases from the following lists are (a) too strong to be studied experimentally; (b) too weak to be studied experimentally; or (c) of directly measurable base strength.
(i) CO32−, O2−, ClO4−, and NO3− in water; (ii) HSO4−, NO3−, ClO4− in H2SO4.
5.10 The aqueous solution pKa values for HOCN, H2NCN, and CH3CN are approximately 4, 10.5, and 20 (estimated), respectively. Explain the trend in these cyano derivatives of binary acids and compare them with H2O, NH3, and CH4. Is the CN group electron donating or withdrawing?
5.11 H3PO4, H3PO3, and H3PO2 all have a pKa value of 2, but the pKa values of HOCl, HClO2, and HClO3 are 7.5, 2.0, and −3.0, respectively. Explain this observation.
5.12 Arrange the following ions in order of increasing acidity in aqueous solution:
Fe3+, Na+, Mn2+, Ca2+, Al3+, Sr2+.
5.13 Use Pauling’s rules to place the following acids in order of increasing acid strength: HNO2, H2SO4, HBrO3, and HClO4 in a nonlevelling solvent.
5.14 Which member of the following pairs is the stronger acid?
Give reasons for your choice. (a) [Fe(OH2)6]3+ or [Fe(OH2)6]2+, (b) [Al(OH2)6]3+ or [Ga(OH2)6]3+, (c) Si(OH)4 or Ge(OH)4, (d) HClO3 or HClO4, (e) H2CrO4 or HMnO4, (f) H3PO4 or H2SO4.
5.15 Arrange the oxides Al2O3, B2O3, BaO, CO2, Cl2O7, SO3 in order from the most acidic through amphoteric to the most basic.
5.16 Arrange the acids HSO4−, H3O+, H4SiO4, CH3GeH3, NH3, HSO3F in order of increasing acid strength.
5.17 The ions Na+ and Ag+ have similar radii. Which aqua ion is the stronger acid? Why?
5.18 When a pair of aqua cations forms an M–O–M bridge with the elimination of water, what is the general rule for the change in charge per M atom on the ion?
5.19 Write balanced equations for the main reaction occurring when (a) H3PO4 and Na2HPO4 and (b) CO2 and CaCO3 are mixed in aqueous media.
5.20 Hydrogen fluoride acts as an acid in anhydrous sulfuric acid and as a base in liquid ammonia. Give the equations for both reactions.
5.21 Explain why hydrogen selenide is a stronger acid than hydrogen sulfide.
5.22 Explain why the Lewis acidity of the silicon tetrahalides follows the trend:
SiI4 < SiBr4 < SiCl4 < SiF4 whereas the trend for the boron trihalides follows the trend BF3 < BCl3 < BBr3 < BI3.
5.23 For each of the following processes, identify the acids and bases involved and characterize the process as complex formation or acid–base displacement. Identify the species that exhibit Brứnsted acidity as well as Lewis acidity.
(a) SO3 + H2O → HSO4− + H+
(b) CH3[B12] + Hg2+ → [B12]+ + CH3Hg+; [B12] designates the Co-porphyrin, vitamin B12 (Section 26.11).
(c) KCl + SnCl2 → K+ + [SnCl3]− (d) AsF3(g) + SbF5(l) → [AsF2]+[SbF6]−(s)
(e) Ethanol dissolves in pyridine to produce a nonconducting solution.
5.24 Select the compound on each line with the named characteristic and state the reason for your choice.
(a) Strongest Lewis acid:
BF3 BCl3 BBr3 BeCl2 BCl3 B(n-Bu)3 B(t-Bu)3
(b) More basic towards B(CH3)3 Me3N Et3N
2-CH3C5H4N 4-CH3C5H4N
5.25 Using hard–soft concepts, which of the following reactions are predicted to have an equilibrium constant greater than 1?
Unless otherwise stated, assume gas-phase or hydrocarbon solution and 25°C.
(a) R3PBBr3+ R3NBF3 R3PRF3+ R3NBBr3
(b) SO2+ (C6H5)3PHOC(CH3)3 (C6H5)3PSO2+ HOC(CH3)3 (c) CH3HgI + HCl CH3HgCl + HI
(d) [AgCl2]−(aq) + 2 CN−(aq) [Ag(CN)2]−(aq) + 2 Cl−(aq) 5.26 Identify the products from the reaction between the following pairs of reagents. In each case identify the species which are acting as a Lewis acid or a Lewis base in the reactions.
(a) CsF + BrF3 (b) ClF3+ SbF5 (c) B(OH)3+ H2O (d) B2H6+ PMe3
5.27 The enthalpies of reaction of trimethylboron with NH3CH3, NH2, (CH3)2NH, and (CH3)3N are −58, −74, −81, and
−74 kJ mol−1, respectively. Why is trimethylamine out of line?
5.28 With the aid of the table of E and C values (Table 5.5), discuss the relative basicities of (a) acetone and dimethyl sulfoxide, (b) dimethyl sulfide and dimethyl sulfoxide. Comment on a possible ambiguity for dimethyl sulfoxide.
5.29 Give the equation for the dissolution of SiO2 glass by HF and interpret the reaction in terms of Lewis and Brứnsted acid–
base concepts.
5.30 Aluminium sulfide, Al2S3, gives off a foul odour
characteristic of hydrogen sulfide when it becomes damp. Write a balanced chemical equation for the reaction and discuss it in terms of acid–base concepts.
5.31 Describe the solvent properties that would (a) favour displacement of Cl– by I– from an acid centre, (b) favour basicity of R3As over R3N, (c) favour acidity of Ag+ over Al3+, (d) promote the reaction 2 FeCl3 + ZnCl2 → Zn2+ + 2 [FeCl4]−. In each case, suggest a specific solvent that might be suitable.
5.32 Catalysis of the acylation of aryl compounds by the Lewis acid AlCl3 was described in Section 5.7b. Propose a mechanism for a similar reaction catalysed by an alumina surface.
5.33 Use acid–base concepts to comment on the fact that the only important ore of mercury is cinnabar, HgS, whereas zinc occurs in nature as sulfides, silicates, carbonates, and oxides.
5.34 Write balanced Brứnsted acid–base equations for the dissolution of the following compounds in liquid hydrogen fluoride: (a) CH3CH2OH, (b) NH3, (c) C6H5COOH.
5.35 Is the dissolution of silicates in HF a Lewis acid–base reaction, a Brứnsted acid–base reaction, or both?
5.36 The f-block elements are found as M(III) lithophiles in silicate minerals. What does this indicate about their hardness?
5.37 Use the data in Table 5.5 to calculate the enthalpy change for the reaction of iodine with phenol.
5.38 In the gas phase, the base strength of amines increases regularly along the series NH3 < CH3NH2 < (CH3)2NH < (CH3)3N.
Consider the role of steric effects and the electron-donating ability of CH3 in determining this order. In aqueous solution, the order is reversed. What solvation effect is likely to be responsible?
5.39 The hydroxoacid Si(OH)4 is weaker than H2CO3. Write balanced equations to show how dissolving a solid M2SiO4 can lead to a reduction in the pressure of CO2 over an aqueous solution. Explain why silicates in ocean sediments might limit the increase of CO2 in the atmosphere.
5.40 The precipitation of Fe(OH)3 discussed in the chapter is used to clarify waste waters, because the gelatinous hydrous oxide is very efficient at the co-precipitation of some contaminants and the entrapment of others. The solubility constant of Fe(OH)3 is Ks = [Fe3+][OH−]3 ≈ 1.0 × 10−38. As the autoprotolysis constant of water links [H3O+] to [OH–] by Kw = [H3O+][OH–] = 1.0 × 10–14, we can rewrite the solubility constant by substitution as [Fe3+]/[H+]3 = 1.0 × 104. (a) Balance the chemical equation for the precipitation of Fe(OH)3 when iron(III) nitrate is added to water. (b) If 6.6 kg of Fe(NO3)3.9H2O is added to 100 dm3 of water, what is the final pH of the solution and the molar concentration of Fe3+, neglecting other forms of dissolved Fe(III)? Give formulas for two Fe(III) species that have been neglected in this calculation.
5.41 The frequency of the symmetrical M–O stretching vibration of the octahedral aqua ions [M(OH2)6]2+ increases along the series, Ca2+ < Mn2+ < Ni2+. How does this trend relate to acidity?
5.42 An electrically conducting solution is produced when AlCl3 is dissolved in the basic polar solvent CH3CN. Give formulas for the most probable conducting species and describe their formation using Lewis acid–base concepts.
5.43 The complex anion [FeCl4]− is yellow whereas the species [Fe2Cl6] trapped in an argon matrix is reddish. Dissolution of 0.1 mol FeCl3(s) in 1 dm3 of either POCl3 or PO(OR)3 produces a reddish solution that turns yellow on dilution. Titration of the red solution in POCl3 with Et4NCl solutions leads to a sharp
colour change (from red to yellow) at a 1:1 mole ratio of FeCl3/ Et4NCl. Vibrational spectra suggest that oxochloride solvents form adducts with typical Lewis acids by coordination of oxygen. Compare the following two sets of reactions as possible explanations of the observations.
(a) Fe2Cl6 + 2 POCl3 [FeCl4]− + 2 [POCl2]+ POCl2+ + Et4NCl Et4N+ + POCl3
(b) Fe2Cl6 + 4 POCl3 [FeCl2(OPCl3)4]+ + [FeCl4]− Both equilibria are shifted to products by dilution.
5.44 In the traditional scheme for the separation of metal ions from solution that is the basis of qualitative analysis, ions of Au, As, Sb, and Sn precipitate as sulfides but redissolve on addition of excess ammonium polysulfide. By contrast, ions of Cu, Pb, Hg, Bi, and Cd precipitate as sulfides but do not redissolve. In the language of this chapter, the first group is amphoteric for reactions involving SH− in place of OH−. The second group is less acidic. Locate the amphoteric boundary in the periodic table for sulfides implied by this information. Compare this boundary with the amphoteric boundary for hydrous oxides in Fig. 5.5. Does this analysis agree with describing S2− as a softer base than O2−? 5.45 The compounds SO2 and SOCl2 can undergo an exchange of radioactively labelled sulfur. The exchange is catalysed by Cl− and SbCl5. Suggest mechanisms for these two exchange reactions with the first step being the formation of an appropriate complex.
5.46 Pyridine forms a stronger Lewis acid–base complex with SO3 than with SO2. However, pyridine forms a weaker complex with SF6 than with SF4. Explain the difference.
5.47 Predict whether the equilibrium constants for the following reactions should be greater than 1 or less than 1:
(a) CdI2(s) + CaF2(s) CdF2(s) + CaI2(s)
(b) [CuI4]2−(aq) + [CuCl4]3−(aq) [CuCl4]2−(aq) + [CuI4]3−(aq) (c) NH2−(aq) + H2O(l) NH3(aq) + OH−(aq)
5.48 For parts (a), (b), and (c), state which of the two solutions has the lower pH:
(a) 0.1 M Fe(ClO4)2(aq) or 0.1 M Fe(ClO4)3(aq) (b) 0.1 M Ca(NO3)2(aq) or 0.1 M Mg(NO3)2(aq) (c) 0.1 M Hg(NO3)2(aq) or 0.1 M Zn(NO3)2(aq)
5.49 Why are strongly acidic solvents (e.g. SbF5/HSO3F) used in the preparation of cations such as I2+ and Se82+, whereas strongly basic solvents are needed to stabilize anionic species such as S42−
and Pb94− ?
5.50 A standard procedure for improving the detection of the stoichiometric point in titrations of weak bases with strong acids is to use acetic acid as a solvent. Explain the basis of this approach.
5.51 Explain the following observations:
(a) The heat of adduct formation of BEt3 with NMe3 is
−72.3 kJ mol−1, whereas that of B(OMe)3 with the same amine is −31.5 kJ mol−1.
(b) The heat of adduct formation of BMe3 with NMe3 is
−75.2 kJ mol−1, whereas with N(SiH3)3 it is about +4 kJ mol−1. 5.52 Write equations to account for the acidity of CO2 and boric acid (B(OH)3), each dissolved in water.
TUTORIAL PROBLEMS
5.1 A paper by Gillespie and Liang entitled ‘Superacid solutions in hydrogen fluoride’ (J. Am. Chem. Soc., 1988, 110, 6053) discusses the acidity of various solutions of inorganic compounds in HF. (a) Give the order of acid strength of the pentafluorides determined during the investigation. (b) Give the equations for the reactions of SbF5 and AsF5 with HF. (c) SbF5 forms an F-bridged dimer (Sb2F11−) in HF: give the equation for the equilibrium between the monomeric and the dimeric species.
5.2 In the reaction of t-butyl bromide with Ba(NCS)2, the product is 91 per cent S-bound tBu–SCN. However, if Ba(NCS)2 is impregnated into solid CaF2, the yield is higher and the product is 99 per cent t-Bu–NCS. Discuss the effect of alkaline earth metal salt support on the hardness of the ambident nucleophile SCN−. (See T. Kimura, M. Fujita, and T. Ando, J. Chem Soc., Chem.
Commun., 1990, 1213.)
5.3 In their paper ‘The strengths of the hydrohalic acids’ (J.
Chem. Educ., 2001, 78, 116), R. Schmid and A. Miah discuss the validity of literature values of the pKa values for HF, HCl, HBr, and HI. (a) On what basis have the literature values been estimated? (b) To what is the low acid strength of HF relative to HCl usually attributed? (c) What reason do the authors suggest for the high acid strength of HCl?
5.4 Superacids are well established but superbases also exist and are usually based on hydrides of Group 1 and Group 2 elements.
Write an account of the chemistry of superbases.
5.5 In a review article (Angew. Chem. Int. Ed., 2009, 48, 2100), E.J. Corey describes asymmetric catalysis by chiral boranes. Show how these chiral boranes, which are Lewis acids, are also able to direct Brứnsted acidity.
5.6 An article by Poliakoff and co-workers describes how a new industrial chemical process was initiated, in which conventional solvents were replaced by scCO2 (Green Chem., 2003, 5, 99).
Explain the advantages and challenges of introducing such changes to traditional technology.
5.7 The reversible reaction of CO2 gas with aqueous emulsions of long-chain alkyl amidine compounds has important practical applications. Describe the chemistry that is involved in this demonstration of ‘switchable surfactants’ (Science, 2006, 313, 958).
5.8 An article by Krossing and co-workers (J. Am. Chem. Soc., 2006, 128, 13427) explains the behaviour of ionic liquids in terms of a thermodynamic cycle approach. Describe the principles that are applied and summarize the predictions that are made.
Oxidation and reduction
A large class of reactions of inorganic compounds can be regarded as occurring by the transfer of electrons from one species to another. Electron gain is called reduction and electron loss is called oxidation; the joint process is called a redox reaction. The species that supplies electrons is the
reducing agent (or ‘reductant’) and the species that removes electrons is the oxidizing agent (or ‘oxidant’). Many redox reactions release a great deal of energy and they are exploited in combustion or battery technologies as well as biology.
Reduction potentials 6.1 Redox half-reactions
6.2 Standard potentials and spontaneity 6.3 Trends in standard potentials 6.4 The electrochemical series 6.5 The Nernst equation Redox stability
6.6 The influence of pH 6.7 Reactions with water
6.8 Oxidation by atmospheric oxygen
6.9 Disproportionation and comproportionation 6.10 The influence of complexation
6.11 The relation between solubility and standard potentials Diagrammatic presentation of potential data
6.12 Latimer diagrams 6.13 Frost diagrams
6.14 Proton-coupled electron transfer: Pourbaix diagrams 6.15 Applications in environmental chemistry: natural waters Chemical extraction of the elements
6.16 Chemical reduction 6.17 Chemical oxidation 6.18 Electrochemical extraction
Further reading Exercises Tutorial problems
Those figures with an in the caption can be found online as interactive 3D structures. Type the following URL into your browser, adding the relevant figure number: www.chemtube3d.com/weller7/[chapter number]F[figure number]. For example, for Figure 3 in Chapter 7, type www.chemtube3d.com/weller7/7F03.
Many of the numbered structures can also be found online as interactive 3D structures: visit www.chemtube3d.com/weller7/[chapter number] for all 3D resources organized by chapter.
Oxidation is the removal of electrons from a species; reduction is the addition of electrons. Almost all elements and their com- pounds can undergo oxidation and reduction reactions and the element is said to exhibit one or more different oxidation states.
In this chapter we present examples of this ‘redox’ chemistry and develop concepts for understanding why oxidation and reduc- tion reactions occur, considering mainly their thermodynamic aspects. We discuss the procedures for analysing redox reactions in solution and see that the electrode potentials of electrochemi- cally active species provide data that are useful for determining and understanding the stability of species and solubility of salts.
We describe procedures for displaying trends in the stabilities of various oxidation states, including the influence of pH. Next, we describe the applications of this information to environmental chemistry, chemical analysis, and inorganic synthesis. The discus- sion concludes with a thermodynamic examination of the condi- tions needed for some major industrial oxidation and reduction processes, particularly the extraction of metals from their ores and new applications in clean and efficient technologies.
6