P.A. Cox - Inorganic chemistry (793955), страница 31
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The pH scale is related to these concentrations. Water at pH 7is neutral, that with pH<7 is acidic (H3O+ dominating) and withpH>7 alkaline or basic (OH− dominating).The acidity constant and the related pKa value give the equilibriumconstant for protolysis. This reaction goes nearly to completion withstrong acids, which are leveled to H3O+, the strongest acid possiblein water.
Weak acids have incomplete protolysis. Strong and weakbases show analogous behavior, the former being leveled to OH−.The acid strength of nonmetal hydrides increases towards the rightand to the bottom of the periodic table. Acid strengths of oxoacidscan be predicted approximately from their formulae by Pauling’srules.
Metal cations with polarizing character are acidic in water, andsome form amphoteric oxides or hydroxides.Lewis acids and bases (C9)Hydrogen (F2)Solvent types and properties(E1)DefinitionsA Brønsted acid is a proton donor, and a Brønsted base a proton acceptor. In this definition an acid-base reactioninvolves protolysis:HA is called the conjugate acid to A−, and A− the conjugate base to HA; HB+ and B form another conjugate acid-basepair.
Examples of some conjugate pairs (with the acid given first) are:134SECTION E—CHEMISTRY IN SOLUTIONWater is both an acid and a base, and this also happens with polyprotic (or polybasic) acids such as H2SO4, whichcan undergo successive protolysis steps to giveandthusis the conjugate base of H2SO4 but theconjugate acid of.This definition of acids and bases should not be confused with the Lewis definition (Topic C9) although there is aconnection: H+ is a Lewis acid, and Brønsted bases are also Lewis bases, but in general Lewis acids such as BF3 are notBrønsted acids, and Brønsted acids such as HCl are not Lewis acids.Brønsted acidity is solvent dependent. Substances such as HCl are covalent molecules that undergo protolysis only insolvents polar enough to solvate ions, and when a base is present (which may be a solvent molecule).
The followingdiscussion concentrates on water, the commonest solvent in which protolytic reactions are studied. (See Topics E1, F5and F8 for some other protic solvents.)pHBeing simultaneously acidic and basic, water undergoes autoprotolysis, also called self-ionization:The equilibrium constant isH3O+ in these equations is often simply written H+. In pure water and in solutions that do not provide any additionalsource of H+ or OH− both ions have molar concentrations equal to 10−7.
Addition of an acid increases [H3O+] andhence decreases [OH−]; addition of a base has the reverse effect.The pH scale is defined byNeutral water has a pH of 7, acidic solutions have lower values (typically 0–7), and alkaline or basic solutionshigher values (7–14). In alkaline solutions [OH−] is thus greater than [H+].Strong and weak behaviorThe equilibrium constant of the protolysis reactionis known as the acidity constant or the acid dissociation constant (Ka) of HA:It is often expressed as a pKa value, defined as(Note that a larger Ka value corresponds to a smaller pKa.) A selection of Ka and pKa values is given in Table 1. If pKa<0(i.e. Ka>1) the equilibrium lies strongly to the right, and HA is called a strong acid.
Acids with pKa>0 (i.e. Ka<1) areE2—BRØNSTED ACIDS AND BASES135Table 1. Some acidity constants in water at 25°CaSeeTopic F4 for the anomalous behavior of carbonic acid.weak acids and undergo only partial protolysis. Strong acids in water include HCl and H2SO4, whereas HF andare weak acids.In a similar way it is possible to define the basicity constant Kb and the corresponding pKb from the equilibriumWe can distinguish strong bases with pKb<0 and the equilibrium lying to the right-hand side (examples being O2− and) and weak bases with pKb>0 (e.g. NH3 and F−, with pKb equal to 4.75 and 10.55, respectively). However, theuse of pKb is unnecessary, as the base reaction above may be combined with autoprotolysis to show thatwhere pKa refers to the conjugate acid BH+.
Thus the pKa values in Table 1 can be used to calculate the pKb values for theconjugate bases A−.As a strong acid such as HCl is fully protolyzed it is impossible to study this species itself in water. H3O+ is effectivelythe strongest acid possible there, and any stronger acid is said to be leveled. In a similar way, strong bases such asare leveled to the strongest base possible in water, OH−. Solvent leveling limits the range of acid-base behavior thatcan be observed in a given solvent, and is one reason for using other solvents with different leveling ranges. Forexample, liquid ammonia is very basic compared with water, and H2SO4 is very acidic (see Topics F5 and F8).Trends in pK valuesA complete thermodynamic analysis of protolysis requires a cycle that includes the solvation of both HA and A−.Entropy is important because of the ordering of water molecules around small ions and species with strongly localizedcharges (see Topic E1).
Entropy changes will therefore tend to reduce the acid strength of any species giving a conjugatebase with strongly localized negative charge. For positive ions protolysis reduces the charge and entropy contributions willincrease the acid strength. Although solvation effects make a rigorous analysis difficult, some straightforward trends canbe observed.136SECTION E—CHEMISTRY IN SOLUTION• AHn compounds: acid strength increases from left to right in the periodic table, for example,This trend is most simply related to the electronegativity increase of the element attached to hydrogen, which givesmore bond polarity in the direction Aδ−−Hδ+ (see Topic B1).
Acid strength also increases down the group, forexample,which is not the order expected from electronegativity. Changes of solvation are important, but one simple contributionto the trend is the decreasing H-X bond strength down the group (see Topic C8).The oxides of nonmetallic elements are generally acidic and give oxoacids in water (e.g. HNO3 and H2SO4).Oxides and hydroxides of metals tend to be basic and form aqua cations. However, metals in high oxidation statescan also form oxoacids (see Topics B2 and F7).• The strengths of oxoacids can be predicted roughly by Pauling’s rules.(i) Writing the formula as XOp(OH)q the first pKa depends largely on the value of p, being roughly equal to 8−5pirrespective of q. Examples with their pKa values are: p=0 (HOCl, 7.2), p=1 (H3PO4, 2.1), p=2 (H2SO4, −2)and p=3 (HClO4, −10).(ii) For polyprotic acids, pKa increases by about five units for each subsequent protolysis step (e.g.7.4;12.7).Although solvation plays a role in these trends, the simplest explanation of rule (i) is that larger values of pgive more scope for the negative charge to be delocalized over the anion.
For example, in ClO− (1) the formalcharge is confined to one oxygen atom, whereas in(2; only two of the four equivalent resonance structuresare shown) it is spread equally over four.• Aqua cations: many metal cations are acidic in water. Table 1 shows that aqueous Fe3+ is a stronger acid than HF.Acidity may be correlated with the ‘polarizing’ power of a cation associated with deviations from the ionic model(see Topic B1). Strongly acidic cations have either a high charge/size ratio (e.g.
Be2+, Al3+, Fe3+) or are derivedfrom metals with low electropositive character (e.g. Hg2+). Salts containing these ions form rather acidic solutions,and if the pH is increased successive protolysis may lead to polymerization and precipitation of an insoluble oxide orhydroxide such as Al(OH).3 Some of these compounds show amphoteric behavior and dissolve in alkalinesolution to give oxoanions.
Thus Al(OH)3, which is very insoluble in a neutral pH range, dissolves at pH>10 to form[Al(OH)4]− (see Topic G5).Section E—Chemistry in solutionE3COMPLEX FORMATIONKey NotesEquilibrium constantsHard and soft behaviorChealates andmacrocyclesEffect of pHRelated topicsComplexes are formed in aqueous solution when a ligand moleculeor ion replaces solvating water molecules. Successive ligands may beattached, giving a series of step wise formation (equilibrium)constants.Class a (hard) cations complex more strongly with smallelectronegative ligands whereas class b (soft) cations have moreaffinity for less electronegative and more polarizable ligands. Thedifference involves entropic and enthalpic solvation terms.Polydentate or chelating ligands have more than one atom availablefor coordination to the metal, and form stronger complexes thanmonodentate ligands.
The effect is enhanced in macrocyclic ligands,which have more rigid structures.Basic ligands become protonated at low pH and complex formation issuppressed.Lewis acids and bases (C9)3d series: aqueous ions (H3)Group 12: zinc, cadmiumComplexes: structure andand mercury (G4)isomerism (H6)Equilibrium constantsA complex in general is any species formed by specific association of molecules or ions by donor-acceptor interactions(see Topic C9). In aqueous solution the most important complexes are those formed between a metal cation andligands, which may be ions (e.g.
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