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Ionic Equilibria

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IONIC EQUILIBRIA1. Concepts of Acid/Base 1.1 Arrhenius Theory 1.1.1 Definition + 1. An acid is a substance which dissolves in water to give hydrogen ions, H (aq). 2. A base is a substance which dissolves in water to give hydroxide ions, OH (aq). + 3. An acid-base reaction (i.e. neutralization) is the combination of H (aq) and OH (aq) to form H2O(l). + Neutralization : H (aq) + OH (aq) H2O(l) Summary : H+ (aq)

and OH (aq) are responsible for acidic and basic properties respectively.

-

1.1.2 Drawbacks of Arrhenius Theory 1. It confines to aqueous solution only. However, many acids are capable of giving hydrogen ions in solvents other than water. e.g. nitric acid in ethanol ionizes to give hydrogen ions 2. Some compounds are not classified as acids because they do not contain the element hydrogen explicitly, although they are capable of releasing hydrogen ions after reacting with water. 2+ e.g. SO2(g) + H2O(l) 2H (aq) + SO3 (aq) 3. Some substances are not classified as bases because they do not contain hydroxide explicitly, although they are capable of releasing hydroxide ions by reacting with water. + e.g. NH3(g) + H2O(l) NH4 (aq) + OH (aq) 1.2 Bronsted-Lowry Theory 1.2.1 Introduction + This theory emphasizes on the role of hydrogen ion, H (i.e. proton). Definition : + 1. An acid is a molecule or ion that can donate a proton, H (i.e. a proton donor). + 2. A base is a molecule or ion that can accept a proton, H (i.e. a proton acceptor). 3. An acid-base reaction is the transfer of proton from an acid to a base (i.e. a proton-transfer reaction). + Neutralization : H (aq) + OH (aq) H2O(l) + NH3(aq) + HCl(aq) NH4Cl(aq) or NH4 (aq) + Cl (aq) Implications : 1. An acid must possess the element hydrogen for donation. 2. A base can either be a negative ion or a neutral molecule with a lone pair of electrons to form dative covalent bond with proton. Reason : Since proton has no electrons, bonding between the base and proton must be dative covalent in nature, with the base providing both electrons a base must have a lone pair of electrons available for donation

H N H H + H+ H N H H H

+

3. Bronsted-Lowry definition includes all substances that fit the Arrhenius definition. In addition, it includes + a. as acids, ions such as HSO4 and NH4 that can donate proton in solution, even though they cannot exist alone; and 2b. as bases, all substances that can accept proton such as CO3 or CN , even though they do not contain hydroxide explicitly.

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1.2.2 Conjugate Acid-Base Pair Consider the ionization of ethanoic acid in water : + CH3COO (aq) + H3O (aq) CH3COOH(aq) + H2O(l) After CH3COOH(aq) has lost its proton, it becomes CH3COO (aq). From reverse reaction, CH3COO (aq) is capable of accepting a proton to form back the acid CH3COO (aq) is itself a CH3COO (aq) is called the CH3COOH(aq) + H2O(l) acid CH3COO (aq) + H3O conjugate base+ (aq) -

of CH3COOH(aq)

Similarly, consider the ionization of ammonia in water : + NH4 (aq) + OH (aq) NH3(aq) + H2O(l) After NH3(aq) has accepted a proton, it becomes NH4 (aq). + From reverse reaction, NH4 (aq) is capable of donating its proton to form back the base NH4 NH4+ (aq) + (aq) +

is itself an is called the NH4 (aq) + OH (aq) conjugate acid+ -

of NH3(aq)

NH3(aq) + H2O(l) base

Conclusion : 1. After an acid loses a proton, it becomes a base and is called the conjugate base of the given acid. 2. After a base accepts a proton, it becomes an acid and is called the conjugate acid of the given base. These pairs of species are known as conjugate acid-base pairs; they convert into each other by the transfer of , i.e. gain/loss of a . Remark : 1. An acid-base reaction must involve both acid and base. Q Both donor and acceptor of proton must be present for transfer of proton to take place. 2. An acid only shows its properties in the presence of a base. Similarly, a base only shows its properties in the presence of an acid. + H. Q For an acid to function, a base must be present to + H. For a base to function, an acid must be present to 3. In an acid-base reaction, since the given acid (acid1) will be converted into its conjugate base (base1) and the given base (base2) will be converted into its conjugate acid (acid2), the process can be represented as : As there are acids and bases on both sides of the equation, the process can go either way and is represented by an . The forward reaction is the transfer of a proton from acid1 to base2, acid + base 1 2 and the reverse reaction is the transfer of the proton from acid2 to base1.

conjugate acid-base pair2 base1 + acid2

Essentially, the process can be taken as base1 and base2 competing to accept proton an acid-base reaction is in fact the competition to gain/lose proton between bases alternatively, an acid-base reaction can also be taken as the competition to gain/lose proton between acids

conjugate acid-base pair1

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Example :

HCl(aq) acid

+

H2O(l) base

Cl (aq) conjugate base of HCl

-

+

H3O (aq) conjugate acid of H2O

+

Exercise : 1. Identify the conjugate acid-base pairs in the following neutralization reactions. + + H2O(l) H3O (aq) + CH3COO (aq) a. CH3COOH(aq)

b. NH4

+ (aq)

+

H2O(l)

H 3O

+ (aq)

+

NH3(aq)

c.

H2O(l)

+

CO3

2(aq)

HCO3 (aq) +

-

OH (aq)

-

2. Write an equation to show that HPO4 Solution :

2(aq)

can act as a Bronsted base in water.

(94, I)

3. Outline the Bronsted-Lowry theory of acids and bases with respect to the following reaction : + H2O(l) + HCl(aq) H3O (aq) + Cl (aq) Solution : H2O(l) + HCl(aq) H 3O+ (aq)

(89, II)

+

Cl (aq)

-

2 conjugate acid-base pairs : 4. Which of the following is a conjugate acid-base pair ? A. NH4+(aq) and NH3(aq) B. H2SO4(aq) and SO42-(aq) C. H3O+(aq) and OH-(aq) D. HCl(aq) and NaOH(aq)

&

(1 mark, 06, I, 3c) 1.2.3 Amphoteric Substances Some substances can both act as an acid to donate a proton, or act as a base to accept a proton. Such behaviour is known as amphoteric. Example : 1. waterH H OH H 2 O H 3 O ++ +

2. HSO4 (aq). As an acid As a base

H H SO4 2 HSO 4 H 2 SO 4

+

+

: :

HSO4 (aq) + OH (aq) HSO4 (aq) + H3O+ (aq)

-

-

SO4

2(aq)

+ H2O(l)

H2SO4(aq) + H2O(l)

For an amphoteric substance, whether it behaves as an acid or a base depends on the other substance present. If a stronger acid is present, the amphoteric species is forced to act as acid/base to accept/donate a proton. If a stronger base is present, the amphoteric species is forced to act as acid/base to accept/donate a proton. Example : Give the equation for the reaction between HCO3 (aq) and (a) NaOH(aq) (b) HCl(aq)

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1.2.4 Acid Strength & Base Strength The strength of an acid is a measure of its tendency to donate a proton. The strength of a base is a measure of its tendency to accept a proton. Both are related to the extent of equilibrium. Example : HCl(aq) + H2O(l) Cl (aq) + H3O+ (aq)

complete ionization/dissociation

This equilibrium lies very far to the right most HCl(aq) dissociate HCl(aq) has a very high tendency to lose proton HCl(aq) is a strong acid Example : HCN(aq) + H2O(l) H 3O+ (aq)

+ CN (aq)

-

incomplete ionization/dissociation

This equilibrium lies very far to the left few HCN(aq) dissociate HCN(aq) has a low tendency to lose proton HCN(aq) is a weak acid The strength of a base can be defined accordingly. Example : Strong base : equilibrium favours the gain of proton CN (aq) + H2O(l) Weak base :-

HCN(aq) + OH (aq)

-

equilibrium far to right

equilibrium not favoured by gain of proton NH3(aq) + H2O(l) NH4+ (aq)

+ OH (aq)

-

equilibrium far to left

1.2.5 Fact :

Strength within Conjugate Acid-Base Pair The conjugate base of a strong acid is a weak base.

Reason : For a strong acid, its tendency to lose proton must be high/low the ability for its conjugate base to accept a proton to form back the strong acid must be high/low, its conjugate base must be strong/weak Example : HCl(aq) + H2O(l) Cl (aq) + H3O+ (aq)

HCl(aq) is a strong acid the above equilibrium lies far to left/right + Cl has little/strong tendency to accept H to form back HCl Cl is a strong/weak base Exercise : Convince yourself with the following statement. The conjugate base of a weak acid is a strong base. The conjugate acid of a strong base is a weak acid. The conjugate acid of a weak base is a strong acid.

Example : Example : Example :

HCN vs CN-

-

CN vs HCN Cl and HCl-

Remark : Strength is a relative term. It takes two acids or two bases to tell which is stronger and which is weaker.

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1.3 Hydrogen Ion + Fact : In aqueous solution, isolated hydrogen ion H (aq) never exists. Reason : Hydrogen ion is just a bare proton with extremely high/low charge density (due to exceptionally H seeks out positive/negative centre immediately and will bond to + H is therefore hydrated with molecules of water+

) in water

Remark : + 1. The most predominant hydrated species is believed to be hydroxonium ion (or hydronium ion) H3O (aq), with + bond with H . Other higher hydrated ions like the lone pair on H2O forming a + + [H(H2O)2] (aq) and [H(H2O)3] (aq) etc. are also believed to exist.

+ O H H+ (aq)

+ H+or H3O+ (aq)

O H

H H+

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