Types of acids and bases

Strictly speaking, the term hydride refers to ionic compounds of hydrogen with the electropositive metals of Groups 1-2; these contain the hydride ion, H^–, and are often referred to as "true" hydrides. However, the term is often used in its more general sense to refer to any binary compound MH_n in which M stands for any element. The hydride ion is such a strong base that it cannot exist in water, so salts such as sodium hydride react with water to yield hydrogen gas and an alkaline solution:

NaH(s) + H₂O → Na⁺ + OH^– + H₂(1-1)

The more electronegative elements form covalent hydrides which generally react as acids, a well-known example being hydrogen chloride, a gas which dissolves readily in water to give the solution we know as hydrochloric acid

HCl(g) + H₂O → H₃O⁺ + Cl^–(1-2)

Most of the covalent hydrogen compounds are weak acids— in some cases, such as methane, CH₄, so weak that their acidic properties are rarely evident. Many, such as H₂O and NH₃, are amphiprotic. The latter compound, ammonia, is a weaker acid than H₂O, so it exhibits basic properties in water

NH₃(g) + H₂O → NH₄⁺ + OH^–(1-3)

but behaves as an acid in non-aqueous solvents such as liquid ammonia itself:

2 NH₃(l) → NH₄⁺ + NH₂^–(1-4)

In general, the acidity of the non-metallic hydrides increases with the atomic number of the element to which it is connected. Thus as the element M moves from left to right across the periodic table or down within a group, the acids MH become stronger, as indicated by the acid dissociation constants shown at the right. Note that

• The formulas shown in red represent "strong" acids (that is, acids stronger than H₃O⁺.)
• Hydrofluoric acid is the only weak member of the hydrohalogen acids.
• Acids weaker than water do not behave as acids in aqueous solution. Thus for most practical purposes, methane and ammonia are not commonly regarded as acids. H₂O itself is treated as an acid only in the narrow context of aqueous solution chemistry.

Attempts to explain these trends in terms of a single parameter such as the electronegativity of M tend not to be very useful. The reason is that acid strengths depend on a number of factors such as the strength of the M-H bond and the energy released when the resultant ions become hydrated in solution. This last factor plays a major role in making HF something of an anomaly amongst the strong acids of Group 17.

Ammonia is such a weak acid that its conjugate base, amide ion NH₂^–, cannot exist in water. In aqueous solution, NH₃ acts as a weak base, accepting a proton from water and leaving a OH^– ion.

An aqueous solution of NH₃ is sometimes called “ammonium hydroxide”. This misnomer reflects the pre-Brønsted view that all bases contain –OH units that yield hydroxide ions on dissociation according to the Arrhenius scheme  NH₄OH → NH₄⁺ + OH^–.

A solution of ammonia in water is more correctly referred to as "aqueous ammonia" and represented by the formula NH₃(aq). There is no physical evidence for the existence of a NH₄OH molecule, but the name seems to remain forever etched on reagent bottles in chemical laboratories and in the vocabularies of chemists.

Compounds containing the hydroxyl group –OH constitute the largest category of acids, especially if the organic acids (discussed separately farther on) are included. M–OH compounds also include many of the most common bases.

Whether a compound of the general type M–O–H will act as an acid or a base depends is influenced by the relative tendencies of the M–O and the O–H bonds to break apart in water. If the M–O bond cleaves more readily, then the –OH part will tend to retain its individuality and with its negative charge will become a hydroxide ion. If the O–H bond breaks, the MO-part of the molecule will remain intact as an oxyanion MO^– and release of the proton will cause the MOH compound to act as an acid.

Note: this is not solely a matter of the relative strengths of the two bonds; the energy change that occurs when the resulting ions interact with water molecules is also an important factor.

Metal hydroxides

In general, if M is a metallic element, the compound MOH will be basic. The case of the highly electropositive elements of Groups 1 and 2 is somewhat special in that their solid MOH compounds exist as interpenetrating lattices of metal cations and OH^– ions, so those that can dissolve readily in water form strongly alkaline solutions; KOH and NaOH are well known examples of strong bases. From the Brønsted standpoint, these different “bases” are really just different sources for the single strong base OH^–. As one moves into Group 2 of the periodic table the M-OH compounds become less soluble; thus a saturated solution of Ca(OH)₂ (commonly known as limewater) is only weakly alkaline. Hydroxides of the metallic elements of the p-block and of the transition metals are so insoluble that their solutions are not alkaline at all. Nevertheless these solids dissolve readily in acidic solutions to yield a salt plus water, so they are formally bases.

Hydroxy compounds of the nonmetals

The acidic character of these compounds, known collectively as oxyacids, is attributed to the displacement of negative charge from the hydroxylic oxygen atom by the electronegative central atom. The net effect is to make the oxygen slightly more positive, thus easing the departure of the hydrogen as H⁺ . The presence of other electron-attracting groups on the central atom has a marked effect on the strength of an oxyacid. Of special importance is the doubly-bonded oxygen atom. With the exception of the halogen halides, all of the common strong acids contain one or more of these oxygens, as in sulfuric acid SO₂(OH), nitric acid NO₂(OH) and phosphoric acid PO(OH)₃. In general the strengths of these acids depends more on the number of oxygens than on any other factor, so periodic trends are not so important.

Most of the halogen elements form more than one oxyacid. Fluorine is an exception; being more electronegative than oxygen, no oxyacids of this element are known. Chlorine is the only halogen for which all four oxyacids are known, and the K_a values for this series show how powerfully the Cl–O oxygen atoms affect the acid strength.

Binary oxides that contain no hydrogen atoms can exhibit acid-base behavior when they react with water.

The division between acidic and basic oxygen oxides largely parallels that between the hydroxy compounds. The oxygen compounds of the highly electropositive metals of Groups 1-2 actually contain the oxide ion O^2–. This ion is another case of a proton acceptor that is stronger than OH^–, and thus cannot exist in aqueous solution. Ionic oxides therefore tend to give strongly alkaline solutions:

O^2– + H₂O → 2OH^–(aq)

In some cases, such as that of MgO, the solid is so insoluble that little change in pH is noticed when it is placed in water. CaO, however, also known as quicklime, is suffciently soluble to form a strongly alkaline solution with the evolution of considerable heat; the result is the slightly-soluble slaked lime, Ca(OH)₂.

Oxygen compounds of the transition metals are generally insoluble solids having rather complex extended structures. Although some will dissolve in acids, they display no acidic properties in water.

The oxides and hydroxides of the metals of Group 3 and higher tend to be only weakly basic, and most display an amphoteric nature. Most of these compounds are so slightly soluble in water that their acidic or basic character is only obvious in their reactions with strong acids or bases. In general, these compounds tend to be more basic than acidic; thus the oxides and hydroxides of aluminum, iron, and zinc all dissolve in mildly acidic solutions, whereas they require treatment with concentrated hydroxide ion solutions to react as acids.

* FeO(OH) is the hydrous oxide of Fe(III); it is essentially the same as Fe(OH)₃.

The product ions in the second column are known as aluminate, zincate, and ferrate. Other products, in which only some of the –OH groups of the parent hydroxides are deprotonated, are also formed, so there are actually whole series of thse oxyanions for most metals.

The binary oxygen compounds of the non-metallic elements tend to produce acidic solutions when they are added to water. Such compounds are sometimes referred to as acid anhydrides (“acids without water”.)

CO₂ + H₂O → H₂CO₃

SO₂ + H₂O → [H₂SO₃]

SO₃ + H₂O → H₂SO₄

P₄O₁₀ + 6 H₂O → 4 H₃PO₄

In some cases, the reaction involves more than simply incorporating the elements of water. Thus nitrogen dioxide, used in the commercial preparation of nitric acid, is not an anhydride in the strict sense:

3 NO₂ + H₂O → 2 HNO₃

When sodium chloride is dissolved in pure water, the pH remains unchanged because neither ion reacts with water. However, a solution of magnesium chloride will be faintly acidic, and a solution of iron(III) chloride FeCl₃ will be distinctly so. How can this be? Since none of these cations contains hydrogen, we can only conclude that the protons come from the water.

The water molecules in question are those that find themselves close to any cation in aqueous solution; the positive field of the metal ion interacts with the polar H₂O molecule through ion-dipole attraction, and at the same time increases the acidity of these loosely-bound waters by making facilitating the departure H⁺ ion. In general, the smaller and more highly charged the cation, the more acidic will it be; the acidity of the alkali metals and of ions like Ag⁺(aq) is negligible, but for more highly-charged ions such as Mg²⁺, Pb²⁺ and Al³⁺, the effect is quite noticeable.

Most of the transition-metal cations form organized coordination complexes in which four or six H₂O molecules are chemically bound to the metal ion where they are well within the influence of the coulombic field of the cation, and thus subject to losing a proton. Thus an aqueous solution of "Fe³⁺" is really a solution of the ion hexaaquo iron III, whose first stage of "dissociation" can be represented as

Fe(H₂O)^₆3+ + H₂O → Fe(H₂O)₅(OH)²⁺ + H₃O⁺

As a consequence of this reaction, a solution of FeCl₃ is acts as stronger acid than an equimolar solution of acetic acid. A solution of FeCl₂, however, will be a much weaker acid; the +2 charge is considerably less effective in easing the loss of the proton.

It should be possible for a hydrated cation to lose more than one proton. For example, an Al(H₂O)₆³⁺ ion should form, successively, the following species:

AlOH(H₂O)₅²⁺ → Al(OH)₂(H₂O)₄⁺→ Al(OH)₃(H₂O)_3⁰ → Al(OH)₄(H₂O)_2^–→ Al(OH)₅(H₂O)^2–→ Al(OH)₆^3–

However, removal of protons becomes progressively more difficult as the charge decreases from a high positive value to a negative one; the last three species have not been detected in solution. In dilute solutions of aluminum chloride, the principal species are actually Al(H₂O)₆³⁺ (commonly represented simply as Al³⁺) and
AlOH(H₂O)₅²⁺, sometimes written as (AlOH)²⁺).

When salts dissolve in water, they yield solutions of anions and cations, so their effects on the pH of the solution will depend on the properties of the particular pair of ions. For a salt such as sodium chloride, the solution wil remain neutral because sodium ions have no acidic properties and chloride ions, being conjugate to the strong acid HCl have negligible proton-accepting tendencies. Ions of this kind are often referred to as "strong" ions (that is, derived from a strong acid and a strong base— HCl and NaOH in the case of NaCl.) The possible outcomes for the other three possibilities are shown below.

The reactions that cause salt solutions to have non-neutral pH values are sometimes still referred to by the older term hydrolysis (“water splitting”) — a reminder of times before the concept of proton transfer acid-base reactions had developed.

The carboxyl group –CO(OH) is the characteristic functional group of the organic acids. The acidity of the carboxylic hydrogen atom is due almost entirely to electron-withdrawal by the non-hydroxylic oxygen atom; if it were not present, we would have an alcohol –COH whose acidity is smaller even than that of H₂O. This partial electron withdrawal from one atom can affect not only a neighboring atom, but that atom’s neighbor as well. Thus the strength of a carboxylic acid will be affected by the bonding environment of the carbon atom to which it is connected. This propagation of partial electron withdrawal through several adjacent atoms is known as the inductive effect and is extremely important in organic chemistry.

A very good example of the inductive effect produced by chlorine (another highly electronegative atom) is seen by comparing the strengths of acetic acid and of the successively more highly substituted chloroacetic acids:

Phenols

The acidic character of the carboxyl group is really a consequence of the enhanced acidity of the –OH group as influenced by the second oxygen atom that makes up the –COOH group. The benzene ring has a similar although weaker electron-withdrawing effect, so hydroxyl groups that are attached to benzene rings also act as acids. The most well known example of such an acid is phenol, C₆H₅OH, also known as carbolic acid. Compared to carboxylic acids, phenolic acids are quite very weak, as indicated by the acid dissociation constants listed below:

Amines and organic bases

We have already discussed organic acids, so perhaps a word about organic bases would be in order. The –OH group, when bonded to carbon, is acidic rather than basic, so alcohols are not the analogs of the inorganic hydroxy compounds. The amines, consisting of the –NH₂ group bonded to a carbon atom, are the most common class of organic bases. Amines give weakly alkaline solutions in water:

CH₃NH₂ + H₂O → CH₃NH₃⁺ + OH^–

For another view of this topic, see the Tutorial on carboxylic acids by Jim Clark

Make sure you thoroughly understand the following essential ideas which have been presented above. It is especially imortant that you know the precise meanings of all the highlighted terms in the context of this topic.

• Given the formula of a binary compound of hydrogen with an element of Z<18, predict whether its aqueous solution be acidic or basic, and write an appropriate equation.
• Do the same for an oxygen- or hydroxy compound of a similar element.
• Define an oxyacid, and explain why many of these are very strong acids.
• Write an equation describing the amphoteric nature of zinc or aluminum hydroxide.
• Define an acid anhydride, and write an equation describing its behavior.
• Explain how metal cations can give acidic solutions.
• Write equations showing why aqueous solutions of some salts are acidic, while others are alkaline.
• Write the formulas of an organic acid and an organic base, and write an equation showing why the latter gives an alkaline solution in water.