Acid-Base Concepts

[Aurelius]

The Bronsted-Lowry Concept of Acids and Bases

The early definitions of acids and bases were strictly experimental.  An acid was a substance whose water solution: (1) turns blue litmus red, (2) neutralizes bases, (3) reacts with active metals with evolution of hydrogen, and (4) tastes sour.  On the other hand, a base was a substance which is water solution: (1) turns red litmus blue, (2) neutralizes acids, (3) tastes bitter, and (4) feels soapy.

Limitations of the Arrhenius Concept

The brilliant Swedish chemical pioneer, Svante Arrhenius, whose ideas were far ahead of his time as they are behind ours, proposed in 1887 that the characteristic properties of acids in water solution are the properties of the hydrogen ion, H⁺, which hydrolyzes (reacts with water) in water to form the hydronium ion, H₃O⁺ as seen in the table in

Post 445267 (missing)

(Aurelius: "Conjugate Acid-Bases and Their Relative Strengths", Chemistry Discourse)
and those of bases, the properties of hydroxide ion, OH^-.  An acid was, therefore, a substance whose water solution contains an excess of , H⁺ ions and a base, of OH^- ions.

It is apparent that the Arrhenius concept, applying as it does to aqueous solutions only, is exceedingly limited.  His idea that the characteristic properties of the following bases, NaOH- sodium hydroxide, Ca(OH)₂- calcium hydroxide, Mg(OH)₂- magnesium hydroxide and NH₄OH- ammonium hydroxide* are actually those of the hydroxide ion are correct.  He assumed, however, that such ionic bases as sodium and potassium hydroxide were ionized to produce hydroxide ion only when dissolved in water.  Today, since we know that ionic compounds exist as ions even in the crystalline state, this view is no longer tenable.

*Ammonium Hydroxide is a theoretical species believed to only exist in aqueous solution as a result of its reactions indicating such a species.  It is formed by dissolving Ammonia gas, NH₃ in water in a hydrolysis reaction analogous to that of the formation of carbonic acid as shown in

Post 445267 (missing)

(Aurelius: "Conjugate Acid-Bases and Their Relative Strengths", Chemistry Discourse).

Arrhenius assumed further that the excess hydrogen ions in a water solution (aqueous) of an acid were formed by a simple equilibrium ionization of the acid as it dissolved in the water.  Thus, for HCl, hydrogen chloride, Arrhenius postulated that the reaction consisted of a dissociation of some of the molecules into positive hydrogen ions and negative chloride ions, which existed in equilibrium with undissociated hydrogen chloride molecules.

HCl + H₂ gives (in equilibrium) H⁺ + Cl^-

In the light of more recent knowledge, however, this interpretation is suspect on two counts.  First, a hydrogen ion, being nothing more than a proton, is unique among cations (positive ions); it contains no electrons at all and its effective radius is only about 10^-13cm compared to 10^-8 for several other simple cations.  It is certain that a proton, with its enormously high ratio of charge to radius, could not exist unhydrated (having gone though hydrolysis with water) in a water solution, completely surrounded by water molecules, whose oxygen atoms, of course, bear a partial negative charge, as well as two unshared pairs of electrons.

In fact, approximate calculations show that the union of a proton with a water molecule would be so exothermic (thermodynamically favorable, giving off heat) to the extent of about 300,000 calories per mole (a standard chemical equivalent used for calculations involving more than one compound or ion) and that the fraction of protons which would remain unhydrated in water solution would be roughly 10^-190.  This is equivalent to saying that free hydrogen (unhydrated) ions (or protons) simply do not exist in aqueous solution.  The same conclusion holds for any other solvent whose molecules have unshared pairs of electrons, and these are the only solvents which normally give conducting solutions (electricity) even with the strongest acids dissolved in them.  It appears certain, therefore, that although the bare proton can be produced in a discharge tube or in nuclear reactions, and can exist in gaseous solutions at very low pressures, it cannot be responsible for the properties of acids in solutions- especially aqueous solutions.

Secondly, the Arrhenius postulation that acids are ionized to form proton flies in the face of all we now know about atomic an molecular structure.  The filled first shell, with its single orbital, holds two electrons (electronic configuration of helium).  This is the stable electronic configuration for hydrogen.

Hence, there could be no possible driving force in chemical reaction that would induce a hydrogen atom in hydrogen chloride to part with its share in the electron pair simply to form a bare proton.  The simple ion which hydrogen tends to form through chemical reactions is not proton, H⁺, but the hydride ion, H^-.

For the sake of convenience, knowledgeable chemists today sometimes represent the ionization of acids as forming simple H⁺, but as with the taking of liberties in the use of good grammar or in the exhibition of good manners, such a practice is suspect unless we are certain that the offender really knows better.

The Ionization of Covalent Acids and Bases

If the hydrogen ion or proton cannot exist unhydrated in water solution, we can most logically represent the ionization of hydrogen chloride in water as involving actual reaction with water molecules:

HCl + H₂O to give H₃O⁺ + Cl^-

The H₃O⁺ ion, because of resemblance to the ammonium ion (NH₄⁺) is called the hydronium ion.  Because water molecules are themselves associated (form weak bonds with each other, called hydrogen bonds), it appears certain that each proton is associated with a variable number of water molecules.  Actually, the hydronium ions in solution are probably (H₅O₂⁺), (H₇O₃⁺), (H₉O₄⁺), etc.  The average extent of hydration being dependent upon the concentration of the acid and the temperature to provide for the most stable configuration of atoms and molecules.  The formula H₃O⁺, however, is customarily used, as it is the simplest formula that denotes a hydrated hydrogen cation.  In fact, the existence of this ion was proved conclusively in 1957 by means of infrared spectroscopy.

On this basis, the ionization in water of any monoprotic (only one hydrogen that can be considered acidic in the molecule) acid, expressed as HA, can be represented by the general equation:

HA + H₂O to give H₃O⁺ + A^-

After review of the material presented, we can now conclude that the properties listed for acids earlier are actually characteristics of the hydronium ion, not acids. 

There is also a similar reaction of water with bases (hydrolysis) to cause the formation of the hydroxide ion from bases that are not ionic, rather covalent.  The ionic bases merely release hydroxide ions from themselves into the water solution rather than react with water to cause the formation of the hydroxide ion.  Covalent compounds cannot do this in such a direct manner.  The most common example, and probably the simplest is the use of ammonia in water to form ammonium ions and hydroxide ions.  The equilibrium reaction is shown below.

NH₃ + H₂O to give NH₄⁺ + OH^-

Thus, a general reaction formula for the hydrolysis of a covalent compound becomes:

B + H₂O to give BH⁺ + OH^-

[Aurelius]

Relative Strengths of Acids and Bases

The strength* of an acid, according to the Bronsted concept, is measured by its tendency to donate a proton; a strong acid is simply one which has a strong tendency to donate a proton.  Similarly, the strength of a base is measured by its tendency to accept a proton; a strong base is simply on which[/b]has a strong tendency to accept a proton[/b].  The different strengths of acids and bases vary between extremely wide limits. 

*Strong acids and bases both exhibit strong conductances for electricity when in solution in addition to the definition given above.

Water solutions of a few acids such as perchloric acid, hydrogen chloride, and nitric acid not only exhibit the charateristic properties of acids very strongly, but also have equivalent conductances and mole numbers comparable to those of ionic compounds at the same concentrations.  It is apparent then that these acids at ordinary concentrations in water solutions are essentially completely ionized (this effect is what causes the high conductance observed in the solutions); in other words, the reactions in which they donate a proton to water proceed virtually to completion.  These acids have a relatively strong tendency to donate a proton and are, therefore, relatively* strong acids. 

*It is always important in chemistry to remember that everything is relative.  Therefore, one acid can be considered a strong acid when compared to a base, but compared to another acid, may appear weak, even so far as to appear to be a base.  In some cases, compounds assumed to be acidic can become the basic compound in a much more strongly acidic environment.

There are a number of acids, such as acetic acid and hydrogen cyanide, whose water solutions display the characteristic acid properties only weakly.  Except in extremely dilute solutions, equivalent conductances are very small, and their mole numbers are only slightly greater than unity.  In other words, these acids are only slightly ionized; they tend to donate a proton to water only to limited extent.  Because of their relatively weak proton-donating tendencies, these acids are considered relatively weak acids.

Other acids, such as sulfurous, nitrous and phosphoric acids are intermediate in strength between perchloric, nitric, hydrochloric and sulfuric (the common strong acids), on the one hand, and acetic acid and hydrogen cyanide, on the other.  They are regarded as moderately weak acids.

Two Important Axioms of the Bronsted Concept

Hydrogen chloride is a relatively strong acid since it tends to give up proton readily; conversely, chloride ion must necessarily be a weak base since it has little tendency to hold on to the proton it already has.  This relationship suggests a highly significant axiom of the Bronsted concept:

The stronger an acid, the weaker its conjugate base and
The stronger a base, the weaker its conjugate acid

Further study of this axiom in some specific cases helps understanding of the simple concept.  As a strong acid, hydrogen chloride is highly ionized, even in concentrated aqueous solution.  At equilibrium, the reaction has proceeded far to the right, with most of the hydrogen chloride ionized to form hydronium and chloride ions.  We can emphasize this fact by using arrows of unequal length to designate the forward and reverse reactions respectively. 

In the equilibrium mixture, two acids, hydrogen chloride and hydronium ion, are competing to donate protons to a base.  The hydrogen chloride wins; it is, therefore, the stronger acid.  Similarly, two bases, water and chloride ion, are competing to accept protons.  The water wins; it is, therefore, the stronger base.  We see that the stronger acid, hydrogen chloride, has the weaker conjugate base, chloride ion.  The stronger base, water, has the weaker conjugate acid, hydronium ion.

In the ionization of acetic acid in water, equilibrium is reached when the reaction has proceeded to the right only to a slight extent, with only a small fraction of the acetic acid present in the form of ions (approximately 6% ionizes, 94% is left as molecular acid). 

Here, acetic acid and hydronium ion are the two acids competing to donate protons to a base.  Hydronium ion triumphs; it is, therefore, the stronger acid.  The bases, water and acetate ion, are competing to accept protons.  Acetate ion triumphs; it is, therefore, the stronger base.  Again, the stronger acid, hydronium ion, has the weaker conjugate base, water.  The stronger base, acetate ion, has the weaker conjugate acid, acetic acid.  Or conversely, we may say that acetate ion is a stronger base than water or chloride ion, because its conjugate acid, acetic acid, is a weaker acid than hydrogen chloride or hydronium ion.

It can be observed, also, that in each reaction when proton transfer has reached equilibrium, the reaction is found to favor the formation of the weaker acid and the weaker base.  When the stronger acid donates protons, the conjugate base formed is the weaker of the two bases.  When the stronger base accepts protons, the conjugate acid formed is the weaker of the two acids.  In fact, it is axiomatic that:

All proton transfer reactions run downhill to form predominantly the weaker acid and the weaker base.

This important generalization applies not only to ionization but also to all other types of proton transfer reactions, such as neutralization and hydrolysis.  In fact, whenever a potential proton donor and a potential acceptor come into contact with each other in solution, proton transfer or protolysis occurs to form a second acid and second base.  The extent of the pyrolysis, however, is determined by the relative strengths of the acids and bases involved.  If the starting acid and base are stronger than the products, protolysis will proceed far to the right before equilibrium is attained.

If the starting acid and base are much weaker than the products, the reaction will proceed only slightly to the right

[Aurelius]

Acid-Base Charts

The gamut of acidity ranges all the way from such powerful acids as perchloric acid to such feeble potential acids as  molecular hydrogen and methane.  We can condense a large body of useful and interesting information by listing a series of acids, intermediate in strength between these two extremes, in order of relative strength.  Such a list of acids, together with their conjugate bases, constitutes an acid-base chart, such as that shown in

Post 445267 (missing)

(Aurelius: "Conjugate Acid-Bases and Their Relative Strengths", Chemistry Discourse).

At least two general methods are widely employed in the comparison of relative acidity.  The first of these is a comparison of proton-donating tendencies of different acids toward the same base.  For moderately strong acids, water is commonly used as the base.  As examples, the acid ionization constant or acidity constant, K_a (a for acidity), for acetic acid at 25*C is 1.8x10^-5; that for hydrogen cyanide is 4.0x10^-10.

CH₃COOH + H₂O –reverse equilibrium-> H₃O⁺ + CH₃COO^-   K_a = 1.8x10^-5

HCN + H₂O –reverse equilibrium-> H₃O⁺ + CN^-  K_a = 4.0x10^-10

(The larger K_a value for acetic acid shows it is a stronger acid than hydrocyanic acid.  This direct relationship between K_a and the strength of an acid is true for any acid pair.)

Acetic acid is, therefore, a stronger acid than hydrogen cyanide; cyanide ion is a stronger base than acetate ion.

A second method is the competitive protolysis method.  Here, one acid is added to the conjugate base of another and the equilibrium concentrations are determined experimentally.  For example, when sodium ethoxide is added to water, we can demonstrate experimentally that the conjugate base of ethyl alcohol, ethoxide ion, reacts fairly completely with the water to form ethyl alcohol and hydroxide ion:

H₂O + C₂H₅O^- –forward reaction-> C₂H₅OH + OH^-

Ethoxide ion is therefore a stronger base than hydroxide ion, and water, a stronger acid than ethyl alcohol. 

An accurate experimental comparison of the relative strengths of two acids as feeble as molecular hydrogen and methane is extremely difficult, and, in fact, the order of acidity for those two exceedingly weak acids is a matter of dispute.

Perchloric acid is the strongest acid in the table in

Post 445267 (missing)

(Aurelius: "Conjugate Acid-Bases and Their Relative Strengths", Chemistry Discourse) ; its conjugate base, perchlorate ion, is consequently the weakest base in the list of bases.  Methane and molecular hydrogen are the weakest of the listed acids; their conjugate bases, methide ion and hydride ion, are consequently the strongest bases.  The remaining acids are listed between the two extremes in order of decreasing acidity, the bases in order of increasing basicity.

The most sweeping generalization implicit in the acid-base chart is this: The reaction of any acid with an equivalent quantity of any base below its conjugate base in the chart (stronger than its conjugate base) will proceed over 50% to the right, to an extent depending qualitatively upon the slope of the line joining the acid and the base.  Perchloric acid, and the base methide ion would react (don’t try it!) with explosive violence, and essentially to completion, to give the infinitely weaker acid, methane, and weaker base, perchlorate ion.  On the other hand, no significant reaction would be expected, for example, between the weak acid ammonia and the very weak base nitrate ion.

Phenol would be expected to react extensively with hydroxide ion, but not with bicarbonate ion; acetic acid, on the other hand, would give fairly complete reaction with either.  This difference is used widely in organic chemistry to distinguish between carboxylic acids, RCOOH (R = H or an alkyl or aryl groups), and to separate compounds, ArOH (where Ar = an aryl group), and to separate compounds of the two types.  Both carboxylic acids and phenols react with sodium hydroxide to form water-soluble, ether-insoluble salts:

RCOOH + OH^- –forward equilibrium-> H₂O + RCOO^-
ArOH + OH^- –forward equilibrium-> H₂O + ArO^-

With sodium hydrogen carbonate (bicarbonate), however, only the carboxylic acids from the carboxylic salts:

RCOOH + HCO₃^- –forward equilibrium-> H₂CO₃ + RCOO^-
ArOH + HCO₃^- –reverse equilibrium-> H₂CO₃ + ArO^-

Study of the acid-base chart serves to emphasize that the Bronsted acid-base concept is truly an extension of that of Arrhenius.  The properties of acids and bases in water solution are the properties of hydronium and hydroxide ions, respectively.  Compounds whose water solutions turn blue litmus red, neutralize alkalis, react with active metals with evolution of hydrogen, and taste sour are then definitely proton donors—in fact, they must be sufficiently powerful to donate proton, at least to a measureable extent, to water thus giving an excess of hydronium ion.  All such compounds are, of course, stronger acids than water; the stronger acids which ionize extensively in water to give hydronium ion must be stronger acids than hydronium ion.  Similarly, covalent compounds such as ammonia and amines, which give the properties associated with bases in the classical sense are then actually proton acceptors—in fact, they must be at least sufficiently strong to accept protons to some extent from the weak acid water to give an excess of hydroxide ion over hydronium ion.

[Aurelius]

Amphiprotic Substances

Another significant generalization emerges from the acid-base chart: many substances may act as acids in certain reactions and as bases in others.  They are said to be amphiprotic. 

The most familiar, and perhaps most interesting, amphiprotic compound is water (as mentioned in

Post 445301 (missing)

(Aurelius: "More acid-base concept chemistry", Chemistry Discourse)).  It acts as a base (proton acceptor) toward perchloric acid, hydrogen chloride, acetic acid and phenol, and as an acid (proton donor) toward ammonia, amines, ethoxide ion, and hydride ion.  In fact, the amphiprotic nature of water is well illustrated in the extremely slight dissociation of self-ionization of water:

H₂O + H₂O –reverse equilibrium-> H₃O⁺ + OH^-  K_w = 1.0x10^-14

In pure water and in all aqueous solutions this equilibrium exists and must satisfy the equation:

([H₃O⁺][OH^-]*) / ([H₂O]²) = K

*The bracketed formulas represent molar concentrations (moles/liter); for precise work a correction factor known as an activity coefficient must be applied, to give the corrected concentration or activity of the species in question.

But because is dilute solution the molar concentration of water [H₂O] can be considered constant (approx. 55.4 moles/liter), the constant value [H₂O]² can be combined with the constant K to give K_w:

K[H₂O]² = K_w = [H₃O⁺][OH^-]

The very important constant K₂ is called the dissociation constant, ionization constant or ion product, of water.  At 25*C it has the value of 1.0x10^-14.

In pure water, all the hydronium and hydroxide ions must arise from the dissociation of water.  This means that [H₃O⁺] = [OH^-], and if x mole of hydronium is formed per liter, x mole of hydroxide is also formed.

[H₃O⁺][OH^-] = 1.0x10^-14
(x)(x) = 1.0x10^-14
(x)² = 1.0x10^-14

x = 1.0x10^-7

In pure water then the concentration of hydronium ion = the concentration of the hydroxide ion = 1.0x10^-7 molar.  The ratio of H₂O concentration to that of hydronium or hydroxide ion is, therefore, :

(55.4 molar) / (1.0x10^-7) = 5.54x10⁸

On the average, then, there is one hydronium ion and one hydroxide ion for each 5.54x10⁸ or 554 million water molecules.  In aqueous solutions which show the characteristic acid properties, the hydronium ion concentration is greater than 1.0x10^-7 molar; in those which exhibit the familiar basic properties, the hydroxide ion concentration is greater than 1.0x10^-7 molar.

Many ions, such as hydrogen sulfate (bisulfate) ion, HSO₄^-, and hydrogen carbonate (bicarbonate) ion, HCO₃^-, are also amphiprotic.  As indicated in the acid-base chart, sulfuric acid ionizes in two steps:

H₂SO₄ + H₂O –equilibrium-> H₃O⁺ + HSO₄^-
HSO₄^- + H₂O –equilibrium-> H₃O⁺ + SO₄⁼

In the first step, hydrogen sulfate ion, HSO₄^-, acts as a base; in the second, as an acid.

[Aurelius]

The Leveling Effect

Since all protolysis reactions tend to run downhill, all the acids which are stronger than hydronium ion are extensively ionized in water to form the weaker acid hydronium ion.  All the acids above hydronium ion in the acid list, for example, are essentially completely ionized in dilute aqueous solution.  It is not surprising, therefore, that such strong acids as perchloric, sulfuric, and nitric acids and hydrogen chloride all appear to have equal strengths in water solution, for the same acid- hydronium ion- is common to all such solutions. 

This means that the apparent acidity in water solutions of all very strong acids is reduced to the same mediocre level- that of the hydronium ion.  This phenomenon is called the leveling effect.  When a stronger acid than hydronium ion is added to water, it instantaneously donates a proton to water to form hydronium ion.  Thus, hydronium ion is the strongest acid which can exist in appreciable concentration in water solution.

The strongest acid that can exist in appreciable concentration in any solvent that can act as a base is the conjugate acid of that solvent.  In acetic acid, that acid is CH₃C(OH)₂⁺. 

As an amphiprotic solvent, water also exerts a leveling effect on strong bases.  Consider what takes place when a stronger base than hydroxide ion is added to water.  In order to select such a base, consider the fact that according to the Bronsted concept sodium hydroxide itself is not a base; it is an ionic compound or salt* whose anion, hydroxide ion, happens to be a relatively strong base ( the stongest base, in fact, which can exist in water solution).  In all of the reactions of sodium hydroxide involving hydroxide ion, sodium ion, Na⁺, is a “spectator ion”- a mere bystander, existing as sodium ion both before and after the reaction.

* If this idea seems startling consider the following series of salts: Na₂SO₄, NaCl, CH₃COONa, C₆H₅ONa, NaNH₂, and NaH.  Why should NaOH be singled out to be designated as a base?

Where shall we look to find a base stonger than the hydroxide ion?  Simply select the conjugate base of a weaker acid than water.  Since ethyl alcohol, ammonia and hydrogen are all weaker acids than water, their conjugate bases, ethoxide ion (C₂H₅O^-), amide ion (NH₂^-), and hydride ion (H^-) are all stronger bases than the conjugate base of water (OH^-).  This fact is highlighted in the acid-base chart.

Sodium amide, NaNH₂, is a crystalline, ionic, solid compound, readily prepared by reaction of sodium with liquid ammonia, just as sodium hydroxide can be obtained from sodium and water.  In the Bronsted concept, sodium amide is viewed as simply a salt whose anion is a strong base – a stronger base than hydroxide ion.  Now, when sodium amide is added to water, a vigorous reaction ensues with liberation of ammonia.  The resulting solution shows all the classical basic properties; in fact, it is simply a solution of sodium hydroxide (and ammonia, because some remains dissolved in the water and does not escape).  The amide ion, as a stronger base than hydroxide ion, simply reacts almost completely with water to form the hydroxide ion. 

H₂O + NH₂^- –forward equilibrium-> NH₃ + OH^-

The strong bases ethoxide ion, in sodium ethoxide (NaOC₂H₅), and hydride ion (H^-) react with water in an analogous manner to form hydroxide ion in the same way.  In fact, hydroxide ion is the strongest base that can exist in any appreciable concentration in aqueous solution.  When a stronger base is added to water, it accepts a proton from water and its apparent basicity is reduced to the level of the hydroxide ion.  This is an example of the base leveling effect.

Again, we can generalize the fact by saying that the strongest base available in any solvent that can act as an  acid is the conjugate base of that solvent.  For acetic acid that base is acetate ion (C₂H₅O^-), for ethyl alcohol it is the ethoxide ion (C₂H₅O^-), and for liquid ammonia it is the amide ion (NH₂^-).

The intrinsic acidity and basicity of solvents which can engage in proton transfer profoundly affect the usefulness of the solvent for reactions demanding strong acids or bases.  In many organic reactions, for example, strong acid or base catalysts are required.  Water cannot be used as a solvent in such reactions; in fact, its mere presence must be scrupulously avoided. (Such is the case with the Grignard Reaction, in which the reagent bearing the reaction’s name-sake acts so strongly as a base as to react with even minute amounts of water, ruining the effectiveness of the reaction.)  For reactions involving or catalyzed by very strong bases or acids, the organic chemist often uses so-called aprotic solvents, which are relatively indifferent to proton transfer.  Hexane (C₆H₁₄), benzene (C₆H₆), toluene (C₆H₅CH₃), Chlorobenzene (C₆H₅Cl), Nitrobenzene (C₆H₅NO₂), and Carbon Tetrachloride (CCl₄) are widely used aprotic solvents.

[Aurelius]

Hydrolysis

The fact that the solution resulting from the addition of sodium amide to water turns red litmus blue constitutes experimental proof that some base must have been introduced into the water.  The base is the amide ion, which accepts a proton from water to form ammonia and hydroxide ion.  This reaction is a specific example of the hydrolysis of a salt, or more strictly, of an anion.  The hydrolysis of amide ion, like that of hydride and ethoxide ions, is an extreme case in that amide ion is a stronger base than hydroxide ion; its reaction with water is essentially complete and an equilvalent amount of hydroxide ion is formed.  Complete hydrolysis is characteristic of anions whose conjugate acids are weaker than water.

There are, however, many salts whose anions are weaker bases than hydroxide ion but which nevertheless give alkaline water solutions.  These are the salts, such as sodium hydrogen carbonate (bicarbonate), potassium cyanide, sodium acetate, and potassium carbonate, the conjugate acids of whose anions are relatively weak, but at least somewhat stronger acids than water.  All these salts dissolve in water without any directly observable reaction, yet the resulting solution in each case turns red litmus blue.  The ion responsible for this behavior is the anion; in each case the anion, although a weaker base than hydroxide ion, must accept a proton from water, at least to a measureable extent, to afford an excess of hydroxide ion.  Thus, when sodium acetate is dissolved in water, the salt (or, more strictly speaking, the acetate ion) undergoes very partial hydrolysis to form acetic acid and hydroxide ion.

H₂O + CH₃COO^- <-reverse equilibrium- CH₃COOH + OH^-

This reaction is the exact reverse of the neutralization of acetic acid with an ionic hydroxide.  In fact, for any salt, hydrolysis (protolysis between an ion and water) is the exact reverse of neutralization (protolysis in which water is one of the products).  We have already seen that the neutralization of acetic acid does not proceed essentially to completion  (and the reversal of this phenomenon is an example of the the LeChatlier principle in action- to be explained later).  We can now explain this fact from another standpoint by saying that the reverse reacion, the hydrolysis, occurs to at least some measureable extent.  Since acetate ion is a weaker base than hydroxide ion—just as water is a weaker acid than acetic acid—the hydrolysis of acetate ion in sodium acetate proceeds to a limited extent.

A water solution of sodium chloride is neutral to litmus.  Why does this example (the chloride anion) not undergo hydrolysis?  Simply because, like the conjugate bases of other strong acids, it is too weak as a base to accept a proton from water to any significant or measureable extent.  Only anions that are relatively strong bases undergo hydrolysis.  In fact, the effect of sodium acetate solution on red litmus is a direct indication that acetate ion in a relatively strong base, and constitutes the simplest experimental proof that acetic acid is a weak acid.

There are salts, such as ammonium and alkylammonium chlorides and sulfates, whose cations contain hydrogen, that exhibit weak acid characteristics in water.  These salts must contain an ion that donates a proton to water to significant or measureable level to form hydronium ion.  In other words, ammonium and alkylammonium salts are slightly hydrolyzed:

NH₄⁺ + H₂O <-reverse equilibrium- H₃O⁺ + NH₃
RNH₃⁺ + H₂O <-reverse equilibrium- H₃O⁺ + RNH₂

The ammonia (or amine) is competing with water for protons.  Ammonia (or amines) as the stronger base retains the greater share of protons.  But the basicity of water is at least comparable to that of ammonia, and water retains a few protons, thus affording a sufficient excess of hydronium ions to give an acid test toward litmus.  In fact,  the litmus test applied to ammonium chloride constitutes simple experimental proof that ammonia is, after all , a rather weak base.

Salts whose cations are relatively strong acids and whose anions are relatively weak bases are easily decomposed, because of protolysis between the two types of ions.  Ammonium acetate and ammonium carbonate (smelling salts) give an appreciable vapor pressure of ammonia even at room temperatures (in other words, you can smell the ammonia).  This is due to the reactions.

NH₄⁺ + CH₃COO^- -equilibrium- CH₃COOH + NH₃
NH₄⁺ + CO₃⁼ -equilibrium- HCO₃^- + NH₃

[Aurelius]

Hydrolysis of Hydrated Salts

If aluminum chloride is added to water, hydrolysis occurs and the resulting solution is strongly acid.  This is true for many salts and psuedo-salts, such as zinc chloride, stannic chloride, cupric sulfate, and platinic chloride, that contain a polyvalent metal.  In fact, in general, the higher the oxidation state of the metal, the more extensive the hydrolysis.

Now it is well known that aluminum ion exists as a hydrate, probably Al(H₂O)₆⁺⁺⁺, in water.  The slight tendency of water molecules to donate protons is undoubtedly greatly increased when the water molecules are attached to a highly charged cation, since there is a strong force of repulsion between the ion and the similarly positively charged proton.  The hydrated aluminum ion would be expected to be a considerably stronger acid than water and to undergo partial hydrolysis:

Al(H₂O)₆⁺⁺⁺ + H₂ –equilibrium- H₃O⁺ + Al(H₂O)₅(OH)⁺⁺

Similarly, for the hydrated zinc ion we may write:

Zn(H₂O)₄⁺⁺ + H₂O –equilibrium-  H₃O + Zn(H₂O)₃(OH)⁺

Indicators

Most acid-base indicators such as litmus, methyl orange, and bromothymol blue are themselves conjugate acid-base systems, in which the conjugate acid differs sharply in color from its conjugate base.  Each indicator undergoes a color change in water solution at a pH dependent upon the strengths of the acid and base in the conjugate system of the indicator.  Often indicators are used to measure relative strengths of acids and bases in water solution.

Buffer Solutions

Solutions which resist marked changes in hydronium ion concentration when diluted or when acids or bases are added to them are called buffer solutions.   An equimolar solution of acetic acid—sodium acetate is a buffer solution.  When strong acid is added to the solution, the hydronium ions react with acetate ions to form acetic acid; when ionic hydroxide is added, the hydroxide ions react with acetic acid to form acetate ions:

H₃O⁺ + CH₃COO^- –forward equilibrium-> CH₃COOH + H₂O
CH₃COOH + OH^- –forward equilibrium-> H₃O⁺ + CH₃COO^-

In either case, the added hydronium ion or hydroxide ion is effectively neutralized, and the changes in hydronium and hydroxide concentrations are slight.  Actually a mixture of any weak acid, somewhat stronger as an acid than water, with its conjugate base, acts as a buffer.

[Aurelius]

Acidity, Basicity and Structure

(“Need Two, From You”—“Have Pair, Will Share”)

Any hydrogen-containing molecular or ionic species is a potential proton donor or Bronsted acid.  But, as we have seen, acidity runs the gamut from powerful acids like H₃SO₄⁺ to such feeble acids as CH₄.  And potentially, any molecular or ionic species with an unshared electron pair is a proton acceptor or Bronsted base.  But basicity runs the gamut from potent bases like CH₃^- to such weak bases as H₂SO₄.

What makes an acid an acid or a base a base?  What structural features confer proton-donating or proton-accepting qualities or tendencies?  There is a direct link between structure and chemical properties.

Structure of the Hydrogen Chloride Molecule

Consider the state of the hydrogen within a hydrogen chloride molecule.  For hydrogen, the stable electronic configuration is that of helium, with two electrons in the valence shell (K-shell).  And in the HCl molecule, hydrogen does indeed, have two electrons in its valence shell. But the situation is not exactly as the electronic formula (as draw using Lewis-dot conventions) would seem to imply.  Chlorine has a much stronger power of attraction for a pair of shared electrons than the hydrogen does.  This quality is known as electrotivity, therefore, chlorine is more electronegative than hydrogen. 

As a result, (speaking in terms of averages) the bonding electron pair in hydrogen chloride is not shared equally between the two atoms.  Rather, it is displaced away from the hydrogen and toward the more electronegative atom, chlorine.  Because of this electron displacement, the bond is said to be polar.  The chlorine end of the linear (straight-line) molecule has a high electron density and is, therefore, partially negatively charged (as electron pairs are what cause a negative charge within an object). 

The hydrogen chloride molecule as a whole is electrically neutral, because the positive and negative charges within each molecule are exactly balanced.  Nevertheless, the molecule is electrically unsymmetrical.  Such molecules, in which the centers of positive and negative charge do not coincide, but are separated by some finite distance, are called polar molecules or dipoles.  In the hydrogen chloride molecule the hydrogen atom is the positive end of the dipole and the chlorine atom is the negative end. 

The positively charged hydrogen atom is electrically deficient.  It is susceptible to attack by any species, which can share a pair of electrons.  Its motto is, “Need two, from you.” (this is common for acids) And such a hydrogen becomes transferable as a proton.  For when the possibility of a better share in an electron pair presents itself, the hydrogen accepts it, leaving the former bonding electron pair behind.

When hydrogen chloride comes into contact with ammonia, for example, the hydrogen accepts a share in the unshared pair on the nitrogen atom of the ammonia molecule to form ammonium ion, leaving the chlorine with the bonding pair behind as chloride ion.

General Structural Nature of Bronsted Acids and Bases

In general, molecules such as hydrogen chloride that behave as Bronsted acids are polar, and the hydrogen that is acidic (transferable as a proton) is at the positive end of the dipole.  In other words, the bonding electron pair is withdrawn from the hydrogen toward the atom or group (of atoms) to which the hydrogen is attached. 

In cations that contain hydrogen atoms, these hydrogen atoms are in an electron-poor or positive environment, because that total ionic unit within which they are contained is electron-deficient or positive.  Hence, the hydrogen atoms in cations tend to be transferred as protons, and all hydrogen-containing cations have acidic tendencies. 

To be a Bronsted base or proton acceptor, a molecule or ion must have at least one unshared pair.  If the base is to have appreciable proton-acceptor tendencies, this unshared pair must be located on an atom in a negative environment, either in an anion or at the negative end of a dipole.  For a base, it is a case of “Have pair, will share.”  If the base is to be strong, this unshared electron pair must not be held too tightly by the atom on which it is located, because of the inherent electronegativity either of the atom, itself, or of the atoms to which it is attached.  The strongest type of base, then, would be an anion with a large negative charge on an atom of low electronegativity.

Relationship of Acidity and Basicity to Charge

From the discussion thus far, we should conclude that, other things being equal, acidity should increase with positive charge and basicity with negative charge.  And this is true.  But it is difficult to keep other things constant.  Fe(H₂O)₆⁺⁺⁺ is indeed a stronger acid than Fe(H₂O)₆⁺⁺⁺ is indeed a stronger acid than Fe(H<sub>2O)6⁺⁺, and Ni(OH)4⁼ a stronger base than Ni(OH)4^-.

H4O⁺⁺, if it does indeed exist, is bound to be a powerful acid.  Oxide ion, on the other hand, is a potent base, as evidenced by the familiar conversion of ionic (basic) oxides to hydroxides water:

From the fact that this reaction proceeds essentially to completion to the right, it is safe to conclude that oxide ion is a considerably stronger base than hydroxide ion. 

Actually, the entire sequence serves to emphasize the obvious—that an acid is a stronger acid than its conjugate base and that a base is a stronger base than its conjugate acid.  Thus, H2PO4^-, dihydrogen phosphate ion, is a stronger acid than HPO4⁼, hydrogen phosphate ion, and a stronger base than H3PO4, phosphoric acid.</sub>

[Aurelius]

Acidity, Basicity and Relative Electronegativities

If we think of an acid, in unit particle terms, as a molecule or an ion containing a positively charged hydrogen, we might expect that the more positive the hydrogen, the stronger the acid.  On that basis, we would predict that for any acid HA, the more strongly A withdraws the bonding  electron pair from H, the more acid HA would be, i.e., the acidity of HA should increase with the electronegativity of the atom or group A.

When applied within proper limits, this idea will give us valuable mileage.  In general, the electronegativity of an atom increases with increasing nuclear charge and with decreasing atomic radius.  Hence the relative electronegativity in a horizontal period of elements in the periodic table increases with atomic number.  In the first horizontal period of eight (excluding the inert gas neon), for example, the relative electronegativity increases in the order:

Li<Be<B<C<N<O<F

We should predict, therefore, that the acidity of the hydrides should vary in the order:

LiH<BeH₂<CH₄<NH₃<H₂<HF*

*Simple BH₃ in unknown, it is only stable as B₂H₆, diborane.

As an active metal, lithium has both a strong tendency to give up its single 2s valence electron and also an extremely low electronegativity.  As a result, not only is lithium hydride devoid of acid properties, but it is actually a salt whose anion is the powerful base H:^-, hydride ion.  In hydride ion, the electron-rich hydrogen shares its pair of electrons with even weakly acid hydrogen atoms, such as those in water or ethyl alcohol, to form its feeble conjugate acid, molecular hydrogen.

H₂ + H:^- –forward equilbrium-> H:H (molecular hydrogen) + OH^-
C₂H₅OH + H:^- –forward equilibrium-> H:H + C₂H₅O

Berylium hydride is also basic, though less so than lithium hydride.

Methane, CH₄, has no unshared electron pair, and therefore, cannot be a base, but it is a very feeble acid.  Ammonia, though a weak acid, is considerably stronger than methane; water is stronger than ammonia, and hydrogen fluoride definitely stronger than water.  These experimentally observed relationships are exactly those predicted from relative electonegativities.

Considering ammonia, water and hydrogen fluoride now as bases, we should expect that the hetero atoms would exert an increasing strong pull on their valence electrons as the nuclear charge increases.  In other words, the electrons should be held increasingly tightly in the order N
NH₃<H₂<HF

Again, these predictions are in line with experimental findings.

Many of the acids and bases of interest to chemists are compounds in which the center is an oxygen or nitrogen atom and the acid hydrogen is one attached to the oxygen or nitrogen.  In making generalizations, we shall find it extremely helpful to consider such acids and bases as derivatives of water or ammonia in which one (or more) of the hydrogen atoms has been replaced by another atom or group.

Each of the parent compounds water and ammonia has its own characteristic acid and base strengths.  The electronegativity of hydrogen is approximately the same as that of carbon, much higher than that of active metals such as sodium and potassium, but considerably lower than that of active non-metals such as oxygen and chlorine.

Now, for any derivative of water or ammonia, YOH or YNH₂, the acid and base strengths can vary widely in either direction from those of the parent compounds, depending upon the electronegativity of Y.

If Y is more electronegative than hydrogen-
then Y will decrease the electron density or electron pair availability at an oxygen or nitrogen atom and will increase electron pair withdrawl from the hydrogen linked to the oxygen or nitrogen atom.  This means that the compound YOH or YNH₂ will be less basic and more acidic than its parent compound.
In fact, highly electronegative atoms or groups are often described as atoms or groups which increase the acidity and decrease the basicity of molecules into which they are introduced.

If Y is less electronegative than hydrogen-
then Y will increase the electron density or electron pair availability at the oxygen or nitrogen atom and will decrease electron pair withdrawl from the hydrogen bonded to the oxygen or nitrogen atom.  As a result, YOH or YNH₂ will be more basic and less acidic than its parent compound.

Relative Acidities and Basicities of Water Derivatives

Let’s look at the derivatives of water, YOH, in which one of the hydrogen atoms of water has been replaced by an atom or group of atoms Y.  Compounds of this type vary from such ionic hydroxides as potassium hydroxide, KOH, and sodium hydroxide, NaOH, to such powerful acids as sulfuric acid SO<sub>2(OH)2 and perchloric acid, ClO3OH.

Before we consider these formulas, let’s look at the electronic structure of water itself.
Oxygen is much more electronegative than hydrogen, and, on the average, the electron pairs shared between the oxygen atom and the hydrogen atoms in the water molecule are displaced away from the hydrogens and toward the oxygen.  The angle between the two H-O bonds is 105*.  Since the two bonds are polar and the molecule is unsymmetrical, the water molecule as a whole is polar with the hydrogens at the positive end of the dipole and the oxygen at the negative end.

The hydrogen atoms are in a somewhat electron-deficient or positive environment and are, therefore, somewhat acidic.  The oxygen atom, with its unshared electron pairs and its negative environment, is a somewhat basic center.

If Y is sodium, which as an active metal has an extremely low electronegativity, the electron pair is released completely to the oxygen with the formation of sodium and hydroxide ions.  This is, of course, the extreme situation in which the oxygen, as part of a negative ion, becomes a highly basic center, and the hydrogen, as part of an anion, loses almost all of its acidity.  We should expect YOH to be ionic whenever Y is a metal with a very small attraction for electrons—in general, an active metal of low ionization potential.

But if Y is chlorine, the highly electronegative chlorine atom will decrease the electron density at the oxygen atom and will increase the effective electronegativity of the oxygen atom in withdrawing the electrons from the hydrogen.  Another way of expressing this idea is to say that the group ClO is more electronegative than HO. As a result, hypochlorous acid is a considerably weaker base and stronger acid than water.

All shades of intermediate acid and base strengths between those of hydroxide ion, on the one hand, and hypochlorous acid, on the hand, are observed in compounds YOH, where the electronegativity of atom Y is greater than that of sodium but less than that of chlorine.  In general, however, we may say that the hydroxides of metals are basic, whereas the hydroxyl compounds of non-metals are acids.

As another example of the effect of an increase in the relative electronegativity of Y in increasing acidity and decreasing basicity in YOH, consider a series of compounds in which all the central atoms to which the OH groups are attached belong to the same family and exhibit the same oxidation state.  In any family of elements in the periodic table, relative electronegativity increases with increasing proximity of the valence electrons to the nucleus—in other words, from bottom to top in the periodic table.  Hence, we should expect the acidity of any family of hydroxyl compounds to decrease and the basicity to increase with increasing atomic number.  For the known hypohalous acids and for the halic acids (where Y is a group of atoms, rather than a single atom), we should predict that increasing acidity should occur as follows:

HOI<HOBr<HOCl

Similarly, for the acids of the nitric acid family, we should predict the increase in acidity with increase in electronegativity as follows:

HNO3>H3PO4>H3AsO4>HSb(OH)6>HBiO3

In all cases, our predictions correspond to experimental facts.

For compounds containing the OH group attached to the same central atom, the acidity increases and the basicity decreases with an increase in the number of electronegative atoms attached to the central atom, i.e., with an increase in the oxidation number of the central atom.  The fact is well illustrated by the oxygen acids of chlorine as follows in order of increasing acidity:

HOCl<HO2Cl<HO3Cl<HO4Cl

The linking of highly electronegative oxygen atoms to chlorine increases the effective electronegativity of the chlorine, so that the electronegativity of Y groups increases in the order Cl2Cl3</sub>Cl.

Now take a look at series of hydroxyl compounds, YOH, of the elements in a horizontal period in the periodic table, such as the following (written with Y bracketed in each case): [Na]OH, [HOMg]OH, [(HO)₂Al]OH, [(HO)OSi]OH, [(HO)₂OP]OH, [(HO)O₂S]OH, and [O₃Cl]OH.  Not only does the electronegativity of the central atom here increase with atomic number, but so does the oxidation number.  Hence, the electronegativity of the entire group Y increases sharply with atomic number, and you can predict with confidence the increase in acidity and decrease in basicity from sodium hydroxide through perchloric acid in the series shown.

Sodium Hydroxide, NaOH (strong base, ionic OH^-); Magnesium Hydroxide, Mg(OH)₂ (strong base, ionic OH^-); Aluminum Hydroxide, Al(OH)₃ (Amphiprotic); Silicic Acid, SiO(OH)₂ (weak acid); Phosphoric Acid, PO(OH)₃ (moderately strong acid); Sulfuric Acid, SO₂(OH)₂ (strong acid); Perchloric Acid, ClO₃OH (very strong acid).

The experimental tests affirm the predictions above.

Predictions about some Ammonia Derivatives

These same principles apply to derivatives of ammonia, YNH₂.  However, because ammonia is basic and less acidic than water, YNH₂ for any given Y is more basic and less acidic than YOH.  For a thorough comparison of the ammonia and water systems of compounds, see Sisler, H.H., “Chemistry in Non-Aqueous Solvents,” Reinhold Publishing Corp, N.Y., (1961), p.26.

Sodamide, NaNH₂, like sodium hydroxide, is ionic, and amide ion is considerably more basic than hydroxide ion.  Chloramine, ClNH₂, is considerably less basic and more acidic than ammonia, but is definitely less acidic and more basic than hypochlorous acid.  Nitramine, NO₂NH₂, is much less basic and more acidic than ammonia, but less acidic and more basic than nitric acid.  Hydrazine, N₂H₂, as would be expected from the fact that the NH₂ is more electronegative than hydrogen is somewhat less basic and more acidic than ammonia.  On the other hand, hydrazine is less acidic and more basic than hydroxylamine, NH₂OH.  In keeping with the generalizations concerning the relative acidities and basicities of oxygen versus nitrogen compounds, the more basic center in hydroxylamine is the nitrogen atom, and the more acidic hydrogen is that attached to oxygen.

[Aurelius]

The Hydrogen Halide Problem
 
Applied within proper limits, the generalization that in any series of compounds, HA, the acidity of H increases with the electronegativity of A is most useful.  We have just seen that this rule is valid for any series of acids in which H is attached, in every case, to the same element or to elements in the same horizontal period.  But electronegativity in not the whole story.
 
Fluorine, at the extreme right and top of the periodic table (excluding inert gases), is the most electronegative element in the periodic table.  But hydrogen fluoride is certainly not the strongest acid of the hydrogen halides.  In fact, acidity of the hydrogen halides increase rapidly in the sequence:
 
HF<<HCl<HBr<HI
 
exactly the reverse of that predicted solely on the basis of electronegativity.
 
The anomaly is, of course, not peculiar to the hydrogen halides, for acidity in the water family of hydrides increases in the order:
 
H₂O<H₂S<H₂Se<H₂Te
 
and in the ammonia family:
 
NH₃<PH₃<AsH₃<SbH₃.
 
In each case, the acidity order is the reverse of the electronegativity order.
 
These anomalies arise because the electronegativity principle of acidity is actually an oversimplification.  It is concerned only with the structure of the acid itself.  But the structure of the conjugate base is also important in determining the acidity of any acid.  In fact, it is not the nature of the acid alone, nor the nature of the base alone, but rather the free energy relationship between the conjugate base and the acid that determines the acidity of the acid.  In other words, for any acid, HA, the more stable A^- is as compared with HA, the stronger HA is as an acid.  A strong acid, then, is one whose conjugate base is relatively stable compared to the acid itself.  Any factor that stabilizes the anion more than it stabilizes the acid should increase the acidity; any factor that stabilizes the acid more than the anion should decrease acidity.
 
Usually, the base-acid relative stability viewpoint of acid strength leads to the same predictions as does the relative electronegativity viewpoint.  For in general, the more strongly the atom on group A withdraws from H in HA (the greater the electronegativity of A), the more stable is A^-, and the greater is the tendency of HA to go to the anionic form A^-.  This is logical, because the more strongly A tends to be the negative end of the dipole in HA, the more easily A should be able to accommodate the pair of electrons left behind when H⁺ leaves, i.e., the more stable the anion A:^- should be.  Thus, for the series, CH₄, NH₃, H₂O and HF, the relative stability of the anions would be expected to increase markedly with increasing electronegativity of the central ion, as indeed it does (left to right, increasing stability):

CH:₃^-<NH:₂^-<OH:^-<F:^-

In fact, the stability of the halide ions also increases with increasing electronegativity in the expected order (left to right, increasing stability):

I:^-<Br:^-<Cl:^-<F:^-

But, as the electronegativity of the halogen atom increases, so also does the strength of the HA bond, * and therefore the stability of HA:

HI<HBr<HCl<HF

* One of the standard methods from the calculation of relative electronegativities of various elements is based on the relative bond strengths of the bonds in the hydrides of the elements.

This, too is expected.  But the real anomaly in the case of the hydrogen halides is that somehow the increase in the stability of the halide ions does not keep pace with the increase in the stability of the hydrogen halides themselves.  In other words, for the hydrogen halides, HA, (and the water and ammonia families, as well), the combined effect of an increase in the electronegativity and a decrease in the size of A is to stabilize the undissociated acid HA more than the anion, A:^-.  The over-all result, therefore, is that there is a decrease of the stability of the halide ion A^- compared to its conjugated acid HA in the sequence HI>HBr>HCl>HF.  This is, therefore, the order of decreasing acidity.

[Aurelius]

Acidity, Basicity and Ion Size

These facts certainly suggest that the stability of an anion A^- is dependent upon other factors in addition to the electronegativity of A.  Certainly, this is the case.  A number of complex factors are involved.  Some of them, such as the solvation of the anion, are rather indeterminate.  In the special case of hydrogen fluoride, the complexing of fluoride ions with undissociated HF molecules to form ions of the type HF₂^- is undoubtedly important.  But there is one simple and clear-cut factor which we should expect to be especially significant in all cases—the size of the anion, or, more specifically, the extent to which the negative charge can be distributed or diffused in the anion.

According to the laws of physics, the stability of any charged system is increased by dispersal of the charge.  We have already seen that bare proton, with its tremendously high ratio of charge to size, is unstable.  With the addition of a water molecule to form the much larger hydronium ion, the charge is effectively dispersed, and a much more stable ion results.  Dispersal or smearing of charge in an ion over a larger volume tends to stabilize the ion.

In the halide ions, the negative charge is spread or dispersed over the entire ion.  The larger the ion, the more effectively the negative charge in spread or dispersed, and the lower is the charge density at any given point in the ion.  Definitely, then, we would expect the sharp increase in ion size in the sequence

F:^-<Cl:^-<Br:^-<I:^-

to serve as a stabilizing factor for the ions in the same order.  We see then, that in this series increasing size of the ion works in exactly the opposite direction from increasing electronegativity in favoring increased ion stabilization.

Similarly, for the hydrides of the water and ammonia series, it is the size of the central atom which determines almost completely the extent to which the negative charge in the conjugate base can be dispersed or diffused.  Once more, we would expect increased ion size to favor increased ion stability in the order:

OH^-<HS^-<The^-<HNH^-<HPH^-<HAsH^-<HSbH^-

Again, the effect of an increase in the size of the central atom in promoting ion stability runs counter to the effect of increasing electronegativity.  The total result is that, in each of these series of hydrides, as the size of the central atom decreases, the stability of the anion, A:^-, inceases less rapidly than that of the acid, HA, itself.  The observed result is that the acidity of acids in any one the families decreases from bottom to top in the periodic table.  These conclusions are summarized by the strengths of the hydrogen halides.

General Applicability of Charge Dispersal Principle

The principle that ion stability increases with charge dispersal has wide applicability.  For a series of compounds containing the OH group attached to the same central atom, the acidity increases with an increase in the oxidation number of the central atom (with the number of electronegative atoms such as oxygen or fluorine linked to the central atom).  Undoubtedly, this due in large part to increasing stabilization (through charge dispersal) of the conjugate base as the number of oxygen or fluorine atoms attached to the central atom increases.  This is exemplified in the Chorine-based oxy-acids with hypochlorous acid being the weakest acid with only one oxygen atom for charge dispersal and perchloric acid being an extremely strong acid having four electronegative oxygen atoms attached to the central chlorine atom for more effective charge dispersal.

[Aurelius]

The Lewis Concept of Acidity and Basicity

By paraphrasing Bronsted slightly, we can set up the following definitions for Bronsted acids and bases:

Acid: any species containing hydrogen which (as a proton) can accept a share in a pair of electrons.

Base: any species which can share an electron pair with a proton.

Now, inspect these definitions from the standpoint of their conceptual breadth.  One word in each definition severely limits and restricts them.  The words are “hydrogen” and “proton”.  You might ask, “Why limit electron-pair-sharing to protons? And why only hydrogen?”

Or, more elegantly, you might quote the versatile chemist G.N. Lewis (1875-1946), for many years an outstanding professor of chemistry at the University of California at Berkeley.  Lewis maintained that “any instructive extension of the idea of acids has been prevented by what I am tempted to call the cult of the proton.”

The Lewis Concept as an Extension of the Bronsted Concept

The logical conceptual extensions of the Bronsted definitions suggest themselves automatically:

Acid: any species containing any atom which can accept a share in an electron pair.

Base: any species which can share an electron pair with any electron pair acceptor.

Or, as Lewis proposed:

An acid is an electron-pair acceptor.
A base is an electron-pair donor.
An acid-base reaction is the sharing of an electron pair with an acid by a base.

These three simple definitions constitute the heart of what is now known as the Lewis concept of acids and bases.  Experimentally and conceptually, they are an extension of the Bronsted definitions.  In one broad sweep, they unleash the concept from its dependence upon hydrogen and open a broad new vista of acid-base chemistry.

The Fundamental Lewis Acid-Base Reaction

The fundamental acid-base reaction in the Lewis sense is the formation of a coordinate covalent bond between an acid and a base.  The base is the electron-pair donor and the acid is the acceptor.  The process is called neutralization, or simply coordination.  The product is a coordinated compound, coordinated complex, or adduct, made up of an acid portion and a base portion.  A typical, and oft-cited example is the reaction of the acid boron trifluoride with the base ammonia to form the cooridinated complex or adduct BF₃NH₃.

The coordinated molecule may be thought of as being made up of the acid portion BF₃ and the base portion NH₃.

Classification of Lewis Acids

Bases in the Lewis system are essentially the same as those in the Bronsted system, because molecules or ions which share a pair of electrons with any other electron-pair acceptors, or Lewis acids, as a rule do so also with proton.

The range of acids, however, is greatly extended by the Lewis concept; in fact, according to the Lewis definition the proton, although an important and powerful acid, is but one of many possible acids.  It is apparent that for Lewis acidity there is but one requirement—at least one available unfilled orbital in a valence shell.  Any species is potentially a Lewis acid which has at least one available unfilled orbital in the valence shell of one of its atoms.  Lewis acids are of several types.

Simple Cations

Theoretically, all simple cations are potential Lewis acids, although their strength as acids varies within wide limits.  Potassium ion (K⁺) is a very weak Lewis acid, aluminum ion (Al⁺⁺⁺) a powerful Lewis acid.  In general, we can expect the acid strength or coordination ability to increase with (a) an increase in positive charge on the ion, (b) an increase in nuclear charge for atoms in any horizontal period, (c) a decrease in ionic radius, and (d) a decrease in the number of shielding electron shells.

This means that Lewis acidity of simple cations tends to increase for the elements from left to right and from bottom to top in the periodic table.  On the basis of the generalizations, we can predict the following representative sequences of increasing acid strength:

Fe⁺⁺<Fe⁺⁺⁺<K⁺<Na⁺<Li⁺<Be⁺⁺<B⁺⁺⁺

From the beginning to end of each of the various series of transition elements, there is a build-up of nuclear charge with simultaneous contraction of ionic radius and no increase in the number of shielding shells.  As a result, many of the transition element cations are strong Lewis acids and tend to form a variety of complex ions.

Actually, not all complexes formed by the arrangement of complexing groups around a simple central cation are held together by true covalent bonds.  In many cases the bonds are of the ion-ion or ion-dipole type.  These complexing groups are called ligands.  Frequently, more than one kind of ligand can be present in a single complex, and the different possible combinations into which cations and ligands can enter yield thousands of known complexes.  Common ligands are such molecular bases as water, ammonia, and ethylene diamine and such ionic bases as hydroxide, cyanide, nitrate, carbonate, oxalate, sulfide, thiosulfate, thiocyanate and halide ions.  In accounting for the orderly grouping of ligands about a central ion, and also for the magnetic and spectral properties of complexes, chemists have achieved remarkable success in recent years through the application of the ligand field theory, an adaptation of the crystal field theory, originally developed by the physicists to account for the properties of crystals.

Of particular interest in organic chemistry are a group of reactive positive ions, thought to be active intermediates in many types of organic reactions.  Among these are the nitronium ion, brominium ion, and the various carbonium and acylium ions. 

Compounds Whose Central Atom has an Incomplete Octet

Among the most Lewis acids are compounds whose central atom has less than a full octet of electrons.  Strength of these Lewis acids, in a general way* increases with (a) increase in nuclear charge of the central atom (for central atoms in the same horizontal period), (b) increase in the number and relative electronegativity of electronegative atoms attached to the central atom, (c) decrease in atomic radius of the central atom, (d) decrease in number of shielding electron shells in the central atom.

·   Actually, the order of acidity (of basicity) in any series of Lewis acids (or bases) often varies toward different bases (or acids).  Frequently, the ease, purely from a spatial standpoint, with which the central atom can accommodate an additional group becomes the critical factor.  In such cases, acidity, especially toward large bases, tends to increase with increasing size of the central atom.

These rules are by no means completely general; some anomalies, such as the increased acidity toward most bases of trimethyl boron over boron trifluoride, are difficult to explain.  Coordination reactions of hydrogen fluoride acts as a base toward aluminum trifluoride suggests that the latter is an extremely strong acid.

Often Lewis coordinations of this type are accompanied by a proton shift, as in the reaction of sulfur trioxide with water to form sulfuric acid.

Although the over-all reaction here is more complex, the key step is simply a Lewis acid-base reaction.

Compounds in which the Octet of the Central Atom can be Expanded

Although carbon and silicon belong to the same family of elements, silicon tetrafluoride and silicon tetrachloride are tremendously more reactive than their carbon analogs, carbon tetrafluoride and carbon tetrachloride.  The explanation is straightforward—the silicon, with its vacant d orbitals, can act as a Lewis acid by expanding its octet.  This is illustrated by the reaction of silicon tetrafluoride with fluoride ion to form fluosilicate ion.

With no available d orbitals, carbon cannot do this, in keeping with the fact that the elements in the first period of eight can accommodate no more than eight electrons in their valence shell.

Actually, the silicon halides typify a large group of halides which, with vacant d orbitals, can expand their octets.  These halides tend to form adducts with halide ions and with organic bases such as ethers.  Halides of this type are vigorously hydrolyzed to form an oxy-acid (or oxide) of the central atom and the appropriate hydrogen halide.  This reaction depends upon the ability of the halides to act as Lewis acids.  The first step in the removal of each halogen atom is undoubtedly the acid-base cooridination of the acid halide with the base water.  This is followed by elimination of the hydrogen halide from the adduct.

All three chlorines are replaced by the hydroxyl group in water, the chloride ion formed accepts the remaining hydrogen atom with the ultimate formation of three molecules of hydrogen chloride, and, after a proton shift, phosphorus acid. 

Nitrogen trichloride, with a central atom (nitrogen) whose octet cannot be expanded, is hydrolyzed quite differently to ammonia and hydrochlorous acid:

NCl₃ + H₂ –forward equilibrium-> NH₃ + 3HOCl

One of the most important chemical reactions in high temperature metallurgical processes is that of silica (the gangue) with basic oxides (the flux) to form silicates (the slag).  This is an acid-base reaction of the acid silica (SiO<sub>2[/sub polymer) with the base oxide ion.

The first step in the mechanism of this reaction is undoubtedly the coordination of oxide ion with silicon in the SiO2 polymer; this requires an expansion of the silicon octet.  Similar Lewis acid-base reactions are of fundamental importance in the manufacture of glass.

Compounds having Multiply-Bonded Acid Centers

There are many compounds, particulary organic, in which a multiply-bonded atom can accept a share in an electron pair with a synchronous shift in a pair of electrons of the multiple bond.  By a slight extension of the Lewis concept, we can classify such compounds as Lewis acids.  Although the atom involved does not, in a strict sense, have an unfilled orbital nevertheless, an orbital is made available as the incoming base forces the intramolecular electron-pair shift.

A familiar example is carbon dioxide.  Consider its neutralization by hydroxide to hydrogen carbonate ion:

CO2 + OH^- –forward equilibrium-> HCO3^-

Carbon dioxide accepts a pair of electrons from the base hydroxide ion in the process of coordination.  The base attacks the less electronegative (and therefore more positive) of the double-bonded atoms and pushes a pair of electrons to the moreelectronegative atomes.  The hydrogen carbonate ion is a resonance hybrid.  Writing of the second equivalent resonance structure will reveal that the increased electron density and negative charge supplied by the hydroxide ion are shared equally by the two oxygen atoms originally present in carbon dioxide.

All nucleophilic attacks on carbonyl compounds, such as cyanohydrin formation, are good examples of the organic side to this type of reaction and coordination.

Elements with an Electron Sextet

To the extent that oxygen and sulfur atoms participate directly in chemical reactions, they may be regarded as Lewis acids.  On this basis, the oxidation with sulfur or sulfite to thiosulfate and of sulfide to polysulfide ion can be classified as acid-base reactions.

Further Experimental Behavior of Acids and Bases

Besides neutralizing or coordinating bases, Lewis acids give additional experimental evidence that the relationship between Bronsted and Lewis acidity is not merely a formal one.  Lewis acid-base titrations can be carried out in a variety of solvents.  For example, boron trichloride and tin (IV) chloride, as acids, can be titrated against the bases pyridine and trimethylamine in chlorobenzene solution with crystal violet as indicator, just as hydrochloric acid can be titrated against sodium hydroxide in water solution.  In both cases, the crystal violet is violet in basic and yellow in acid solution.  A wide variety of other acids and base, as well as other indicators, such as butter yellow an thymol blue, can be used.  Many Lewis acids, including carbon dioxide, sulfur dioxide, sulfur trioxide, tin (IV) chloride, aluminum chloride, and phosphorus trichloride, when added to water to give solutions which test acid.

The Lewis concept has revolutionized not only the theory, but also the practice, of acid catalysis, particularly in organic chemistry.  Aluminum chloride, boron trifluoride, sulfur trioxide, and ferric bromide are all important acid catalysts, used interchangeably for many reactions with Bronsted acids such as sulfuric acid and hydrogen fluoride.  Often the Lewis acid catalysts are far superior, and in some cases are effective for reactions where Bronsted acids are useless.</sub>

[Lilienthal]

Hey Aurelius, maybe you could just post the links to good acid / base pages on the net - I mean, this is a discussion forum and not a textbook...  

[Aurelius]

I already had these available to me, and searching takes time- especially since the quality of web sites is,well, not as high as it should be. plus, i'm done with this topic.  I wouldn't have been so concerned with it, but acid-base chemistry is extremely important to this sites major topics. 

In the future, I'll stick to web sites.