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8 Acids and bases8.1 Reactions of acids and bases

The reactions of acids and basesAcids and bases are substances that are familiar, both from the laboratory and from everyday life. A few everyday acids are shown in Figure 8.1. Examples of acids include hydrochloric acid (HCl) and sulfuric acid (H2SO4), whereas bases include sodium hydroxide (NaOH) and ammonia (NH3).

If we look at the formulas of some acids (Table 8.1), we can see that all acids contain at least one H atom. It is this H atom that leads to the characteristic properties of an acid. In the simplest de nition of acids and bases, an acid is de ned as a substance that produces hydrogen ions in solution, and a base is a substance that produces hydroxide ions in solution. The characteristic reactions of acids are described below. In each case we can write an ionic equation where the acid is represented by H+(aq).

Reactions of acids with metalsAcids react with reactive metals such as magnesium to produce a salt and hydrogen gas. For example:

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

Ionic equation: Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)

Unreactive metals such as copper do not react with dilute acids.

Reactions of acids with carbonates and hydrogencarbonates

Figure 8.1 These all contain acids. Cola is a solution of phosphoric acid, among other things.

Acids react with carbonates and hydrogencarbonates to produce a salt, carbon dioxide and water. For example:

Na2CO3(aq) + H2SO4(aq) Na2SO4(aq) + H2O(l) + CO2(g)

(ionic equation: CO32(aq) + 2H+(aq) H2O(l) + CO2(g))

and NaHCO3(aq) + HCl(aq) NaCl(aq) + H2O(l) + CO2(g)sodium hydrogencarbonate

Reactions of acids with bases and alkalisBases are metal oxides, such as copper(II) oxide. These react with acids to form a salt and water. This is called a neutralisation reaction.

acid + carbonate/hydrogencarbonate salt + carbon dioxide + watermetal + acid salt + hydrogen

Acid Formula

hydrochloric HCl

sulfuric H2SO4

nitric HNO3

carbonic H2CO3

ethanoic CH3COOH

benzoic C6H5COOH

Table 8.1 The formulas for some common acids, with the H that is lost as H+ shown in red.

Learning objectives

Understand how to distinguish between an acid and a base using an indicator

Give equations for reactions of acids with metals, bases, alkalis, carbonates and hydrogencarbonates

3138 ACIDS AND BASES

For example:

CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l)

Alkalis are bases that dissolve in water, and the reaction between an acid and alkali is the same as that between an acid and a base. For example:

NaOH(aq) + HNO3(aq) NaNO3(aq) + H2O(l)

Ionic equation: OH(aq) + H+(aq) H2O(l)The reactions of acids with ammonia solution, an alkali, are often

written slightly di erently, e.g.

NH3(aq) + HCl(aq) NH4Cl(aq)

Ammonia solution is equivalent to ammonium hydroxide (NH4OH); in some laboratories bottles are labelled ammonia solution and in others they are labelled ammonium hydroxide. Ammonia is in equilibrium with the ammonium ion and the hydroxide ion:

NH3(aq) + H2O(l) NH4+(aq) + OH(aq)

The reaction with hydrochloric acid could also have been written:

NH4OH(aq) + HCl(aq) NH4Cl(aq) + H2O(l)

which corresponds to the general equation of alkali + acid salt + water.

8.2 Defi nitions of acids and bases

The BrnstedLowry defi nition of acids and basesWe have considered one de nition of acids and bases above, but a slightly more general de nition is that due to Brnsted and Lowry.

The BrnstedLowry de nition of acids and bases is:Acid = proton (H+) donorBase/alkali = proton (H+) acceptor

Consider the reaction of ethanoic acid with water (Figure 8.2, overleaf ). In the forward direction, the CH3COOH donates a proton (H+) to the H2O the CH3COOH is thus an acid, as it donates a proton. The H2O accepts a proton and therefore acts as a base. In the reverse direction, H3O+ donates a proton to CH3COO; H3O+ is thus an acid and CH3COO is a base.

When CH3COOH acts as an acid and donates a proton, it forms a base: CH3COO. CH3COO is called the conjugate base of CH3COOH. CH3COOH and CH3COO are called a conjugate acidbase pair. Similarly, when H2O acts as a base and accepts a proton, it forms H3O+, which acts as an acid in the reverse direction. H3O+ is the conjugate acid of H2O, and H3O+ and H2O are a conjugate acidbase pair.

Learning objectives

Understand the BrnstedLowry and Lewis de nitions of acids and bases

Recognise BrnstedLowry and Lewis acids and bases

Conjugate acid and base pairs always di er by one proton (H+).

base + acid salt + water

alkali + acid salt + water

An alkali is obtained when a base dissolves in water.

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We can label the conjugate acidbase pairs in the equation:

CH3COOH(aq) + H2O(l) CH3COO(aq) + H3O+(aq) acid 1 base 2 base 1 acid 2

Species with the same number form a conjugate pair.

Some more examples of acidbase reactionsLet us look at another example:

NH3(aq) + H2O(l) NH4+(aq) + OH(aq) base 1 acid 2 acid 1 base 2

In the forward direction: H2O donates a proton to NH3. H2O therefore acts as an acid, and as NH3 accepts the proton, it acts as a base.

When H2O donates the proton, it forms OH. OH is the conjugate base of H2O, and H2O and OH are a conjugate acidbase pair. We could also say that H2O is the conjugate acid of OH.

In the reverse direction: NH4+ donates a proton to OH. NH4+ therefore acts as an acid, and as OH accepts the proton, it acts as a base.

When NH3 accepts a proton it forms NH4+. NH4+ is the conjugate acid of NH3, and NH3 and NH4+ are a conjugate acidbase pair. We could also say that NH3 is the conjugate base of NH4+.

The Lewis theory of acids and basesThe BrnstedLowry de nition of acids and bases is only one of several de nitions of acids and bases, and we will now consider the Lewis de nition.

LewisAn ACid is an electron pair ACceptorA base is an electron pair donor

H HO

H

OH H++

HOH H+

H

dativecovalentbond

a

dativecovalentbond

b c

Figure 8.3 Three different representations of the structure of H3O+. (a,b) The electronic structure; (c) the trigonal pyramidal shape.

CH3COOH(aq) + H2O(I) CH3COO(aq) + H3O

+(aq)a

H + donate

d

CH3COOH(aq) + H2O(I) CH3COO(aq) + H3O

+(aq)b

H + donate

d

Figure 8.2 The reaction of ethanoic acid with water. (a) In the forward direction ethanoic acid acts as an acid, as it donates a proton to water; (b) in the reverse direction, H3O+ acts as an acid, as it donates a proton to CH3COO.

It was originally thought that all acids contain oxygen, and the names of this element in English, German (Sauersto ) and several other languages re ect this mistaken assumption.

The hydrated proton may be written as H+(aq) or H3O+(aq). H3O+ (called the hydronium ion, hydroxonium ion or oxonium ion) has the structure shown in Figure 8.3.

Water is sometimes described as amphiprotic, as it can accept or donate a proton.

BrnstedLowryAn acid is a proton (H+) donorA base/alkali is a proton (H+) acceptor

3158 ACIDS AND BASES

The Lewis de nition is more general than the BrnstedLowry de nition and can be applied to reactions that do not involve the transfer of a proton. It is also useful when considering acidbase reactions in solvents other than water.

The Lewis de nition of acids and bases encompasses all the BrnstedLowry reactions, as acceptance of a proton by a base must involve donation of an electron pair to the proton. Thus, if we consider the protonation of ammonia:

NH3 + H+ NH4+

NH3 is the Lewis base, as it donates an electron pair to H+, which is the electron pair acceptor, i.e. Lewis acid.

The reaction between BF3 and NH3 (Figure 8.4) is a Lewis acidbase reaction that does not involve the transfer of a proton:

NH3 + BF3 H3N:BF3

A dative covalent bond is always formed in a Lewis acidbase reaction.

Formation of a complex ion by a transition metal ion (Figure 8.5) is another example of a Lewis acidbase reaction. The transition metal ion is the Lewis acid, and the ligand is the Lewis base. The ligands bond to the transition metal ion through the formation of dative covalent bonds. For example:

Fe2+ + 6H2O [Fe(H2O)6]2+

The last two reactions would not be described as acidbase according to the BrnstedLowry theory.

For a substance to act as a Lewis base, it must have a lone pair of electrons. For a substance to act as a Lewis acid, it must have space to accept a pair of electrons in its outer shell.

For example, NH3 and H2O can both act as Lewis bases, as they have a lone pair of electrons, and H+ and BX3 (where X is a halogen) can act as Lewis acids because they have space in their outer shells to accept a pair of electrons.

In all further discussion of the behaviour and properties of acids, we will be using the BrnstedLowry de nition.

Lewisacid

Lewisbase

NH FB

H

H

F

F

Lewisacid

Lewisbase dative covalent

bond

NH H+

H

H

Figure 8.4 An adduct is formed when NH3 and BF3 react together.

Only Higher Level students should encounter transition metal complex ions.

OH

Fe2+

H

O O

O O

OHH

H

HH

H

H

HH

H

Figure 8.5 A transition metal complex ion.

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The relationship between depth and simplicity

There are various theories of acids and bases. We have encountered the BrnstedLowry and Lewis theories. The Lewis theory is a more sophisticated theory and extends the BrnstedLowry theory, which is limited to describi