Ionic Equilibrium : Acids and Bases

Acids and Bases

(1) Arrhenius concept : According to Arrhenius concept all substances which give H⁺ ions when dissolved in water are called acids while those which ionise in water to furnish OH^– ions are called bases.

$\[\underset{(Acid\,\,{{H}_{2}}O)}{\mathop{HCl}}\,\rightleftharpoons \underset{(aq)}{\mathop{{{H}^{+}}}}\,+\underset{(aq)}{\mathop{C{{l}^{-}}}}\,\,\,;\,\,\,\underset{(Base\,\,{{H}_{2}}O)}{\mathop{NaOH}}\,\rightleftharpoons \underset{(aq)}{\mathop{N{{a}^{+}}}}\,+\underset{(aq)}{\mathop{O{{H}^{-}}}}\,$

Some acids and bases ionise almost completely in solutions and are called strong acids and bases. Others are dissociated to a limited extent in solutions and are termed weak acids and bases.

HCl, HNO₃, H₂SO₄, HClO₄, etc., are examples of strong acids and NaOH, KOH, (CH₃)₄NOH are strong bases. Every hydrogen compound cannot be regarded as an acid, e.g., CH₄ is not an acid. Similarly, CH₃OH. C₂H₅OH, etc., have OH groups but they are not bases.

Actually free H⁺ ions do not exist in water. They combine with solvent molecules, i.e., have strong tendency to get hydrated.

$\[HX+{{H}_{2}}O\rightleftharpoons \underset{(Hydronium\,\,ion)}{\mathop{{{H}_{3}}{{O}^{+}}}}\,+{{X}^{-}}$

The proton in aqueous solution is generally represented as H⁺(aq). It is now known that almost all the ions are hydrated to more or less extent and it is customary to put (aq) after each ion.

The oxides of many non-metals react with water to form acids and are called acidic oxides or acid anhydrides.

$\[C{{O}_{2}}+{{H}_{2}}O\to {{H}_{2}}C{{O}_{3}}\rightleftharpoons 2{{H}^{+}}(aq)+CO_{3}^{2-}(aq)$

$\[{{N}_{2}}{{O}_{5}}+{{H}_{2}}O\to 2NH{{O}_{3}}\rightleftharpoons 2{{H}^{+}}(aq)+2NO_{3}^{-}(aq)$

Many oxides of metals dissolve in water to form hydroxides. Such oxides are termed basic oxides.

$\[N{{a}_{2}}O+{{H}_{2}}O\to 2NaOH\rightleftharpoons 2N{{a}^{+}}(aq)+2O{{H}^{-}}(aq)$

The substances like NH₃ and N₂H₄ act as bases as they react with water to produce OH⁻ions.

$\[N{{H}_{3}}+{{H}_{2}}O\to N{{H}_{4}}OH\rightleftharpoons NH_{4}^{+}(aq)+O{{H}^{-}}(aq)$

(i) Utility of Arrhenius concept : The Arrhenius concept of acids and bases was able to explain a number of phenomenon like neutralization, salt hydrolysis, strength of acids and bases etc.

(ii) Limitations

(a) For the acidic or basic properties, the presence of water is absolutely necessary. Dry HCl shall not act as an acid. HCl is regarded as an acid only when dissolved in water and not in any other solvent.

(b) The concept does not explain acidic and basic character of substances in non-aqueous solvents.

(c) The neutralisation process is limited to those reactions which can occur in aqueous solutions only, although reactions involving salt formation do occur in absence of solvent.

(d) It cannot explain the acidic character of certain salts such as AlCl₃ in aqueous solution.

(2) Bronsted–Lowry concept (The proton – donor Or acceptor concept): According to this concept,

“An acid is defined as a substance which has the tendency to give a proton (H⁺) and a base is defined as a substance which has a tendency to accept a proton. In other words, an acid is a proton donor whereas a base is a proton acceptor.”

The above definition may be explained with the help of the following examples,

$\[\underset{\text{Acid}}{\mathop{HCl}}\,\,\,+\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\rightleftharpoons {{H}_{3}}{{O}^{+}}+C{{l}^{-}}$ …..(i)

$\[\underset{\text{Acid}}{\mathop{C{{H}_{3}}COOH}}\,\,\,+\,\,\underset{\text{Base}}{\mathop{{{H}_{2}}O}}\,\rightleftharpoons {{H}_{3}}{{O}^{+}}+C{{H}_{3}}CO{{O}^{-}}$ …..(ii)

$\[\underset{\text{Base}}{\mathop{N{{H}_{3}}}}\,\,\,+\,\,\underset{\text{Acid}}{\mathop{{{H}_{2}}O}}\,\rightleftharpoons NH_{4}^{+}+O{{H}^{-}}$ …..(iii)

$\[\underset{\text{Base}}{\mathop{CO_{3}^{-2}}}\,\,\,+\,\,\underset{\text{Acid}}{\mathop{{{H}_{2}}O}}\,\rightleftharpoons HCO_{3}^{-}+O{{H}^{-}}$ …..(iv)

$\[\underset{\text{Acid}}{\mathop{HCl}}\,\,+\,\,\underset{\text{Base}}{\mathop{N{{H}_{3}}}}\,\rightleftharpoons NH_{4}^{+}+C{{l}^{-}}$ …..(v)

The following important results may be derived from above equations,

(i) HCl and CH₃COOH are acids because they donate a proton to H₂O.

(ii) NH₃and $\[CO_{3}^{2-}$ are bases because they accept a proton from water.

(iii) Not only molecules but even the ions can act as acids or bases e.g., $\[CO_{3}^{2-}$ ion in the above case is acting as a base.

(iv) In the first two reactions, water is accepting a proton and hence is base. In the next two reactions, water is donating a proton and hence is acting as an acid. Thus water acts both as an acid as well as a base and hence is called amphoteric or amphiprotic. Other examples are

$\[HCO_{3}^{-},\,\,{{H}_{2}}PO_{4}^{-},\,\,NH_{2}^{-},\,\,HSO_{4}^{-}$

(v) The reaction (v) indicates that Bronsted – Lowry definitions of acids and bases are not restricted to aqueous solutions. In this reaction, HCl is acid because it gives a proton and NH₃ is a base because it accepts the proton.

(vi) The reverse reactions are also acid-base reactions. For example, in reaction (i), in the reverse process, H₃O⁺ can give a proton and hence is an acid while Cl⁻− can accept the proton and hence is a base. Thus there are two acid-base pairs in reaction (i). These are HCl – Cl⁻ and H₃O⁺ − H₂O. These acid-base pairs are called conjugate acid-base pairs. Obviously.

A conjugate pair of acid and a base differs by a proton only i.e., Conjugate acid ⇌ Conjugate base + H⁺

The conjugate acid-base pairs in reactions (i) to (v) may be represented as follows,

Conjugate Pair I Conjugate Pair II



Acid₁		*Base₂*		Acid₂		*Base₁*
HCl	+	H₂O	⇌	H₃O⁺	+	Cl⁻
CH₃COOH	+	H₂O	⇌	H₃O⁺	+	CH₃COO⁻
H₂O	+	NH₃	⇌	$\[NH_{4}^{+}$	+	OH⁻
H₂O	+	$\[CO_{3}^{2-}$	⇌	$\[HCO_{3}^{-}$	+	OH⁻
HCl	+	NH₃	⇌	$\[NH_{4}^{+}$	+	Cl⁻

Acid-base chart containing some common conjugate acid-base pairs

Acid			Conjugate base
HClO₄	(Perchloric acid)	Increasing order of acidic strength	$\[ClO_{4}^{-}$	(Perchlorate ion)	Increasing order of basic strength
H₂SO₄	(Sulphuric acid)		$\[HSO_{4}^{-}$	(Hydrogen sulphate ion)
HCl	(Hydrogen chloride)		Cl⁻	(Chloride ion)
HNO₃	(Nitric acid)		$\[NO_{3}^{-}$	(Nitrate ion)
H₃O⁺	(Hydronium ion)		H₂O	(Water)
	(Hydrogen sulphate ion)		$\[SO_{4}^{2-}$	(Sulphate ion)
H₃PO₄	(Ortho phosphoric acid)		$\[{{H}_{2}}PO_{4}^{-}$	(Dihydrogen phosphate ion)
CH₃COOH	(Acetic acid)		CH₃COO⁻	(Acetate ion)
H₂CO₃	(Carbonic acid)		$\[HCO_{3}^{-}$	(Hydrogen carbonate ion)
H₂S	(Hydrogen sulphide)		HS⁻	(Hydrogen sulphide ion)
$\[NH_{4}^{+}$	(Ammonium ion)		NH₃	(Ammonia)
HCN	(Hydrogen cyanide)		CN⁻	(Cyanide ion)
C₆H₅OH	(Phenol)		C₆H₅O⁻	(Phenoxide ion)
H₂O	(Water)		OH⁻	(Hydroxide ion)
C₂H₅OH	(Ethyl alcohol)		C₂H₅O⁻	(Ethoxide ion)
NH₃	(Ammonia)		$\[NH_{2}^{-}$	(Amide ion)
CH₄	(Methane)		$\[CH_{3}^{-}$	(Methyl carbanion)

Levelling effect and Classification of solvents : In acid-base strength series, all acids above H₃O⁺ in aqueous solution fall to the strength of H₃O⁺. Similarly the basic strength of bases above OH⁻ fall to the strength of OH⁻ in aqueous solution. This is known as levelling effect. Levelling effect of water is due to its high dielectric constant and strong proton accepting tendency.

The strength of an acid also depends upon the solvent. the acids HClO₄, H₂SO₄, HCl and HNO₃ which have nearly the same strength in water will be in the order of HClO₄ > HCl > HNO₃ in acetic acid, since the proton accepting tendency of acetic acid is much weaker than water. So the real strength of acids can be judged by solvents. On the basis of proton interaction, solvents can be classified into four types,

(i) Protophilic solvents : Solvents which have greater tendency to accept protons, i.e., water, alcohol, liquid ammonia, etc.

(ii) Protogenic solvents : Solvents which have the tendency to produce protons, i.e., water, liquid hydrogen chloride, glacial acetic acid, etc.

(iii) Amphiprotic solvents : Solvents which act both as protophilic or protogenic, e.g., water, ammonia, ethyl alcohol, etc.

(iv) Aprotic solvents : Solvents which neither donate nor accept protons, e.g., benzene, carbon tetrachloride, carbon disulphide, etc.

acts as acid in H₂O, stronger acid in NH₃, weak acid in CH₃COOH, neutral in C₆H₆ and a weak base in HF.

$\[\underset{Base}{\mathop{HCl}}\,\,\,\,+\,\,\underset{Acid}{\mathop{HF}}\,\,\,\,\to \,\,\,\underset{Acid}{\mathop{{{H}_{2}}C{{l}^{+}}}}\,\,\,+\,\,\,\underset{Base}{\mathop{{{F}^{-}}}}\,$

Utility of Bronsted – Lowry concept

(i) Bronsted – Lowry concept is not limited to molecules but includes even the ionic species to act as acids or bases.

(ii) It can explain the basic character of the substances like Na₂CO₃, NH₃ etc. (which do not contain OH⁻ group and hence were not bases according to Arrhenius definition) on the basis that they are proton acceptors.

(iii) It can explain the acid-base reactions in the non-aqueous medium or even in the absence of a solvent (e.g., between HCl and NH₃).

Limitations

(i) There are a number of acid-base reactions in which no proton transfer takes place, e.g.,

$\[\underset{Aci{{d}_{1}}}{\mathop{S{{O}_{2}}}}\,+\underset{Bas{{e}_{2}}}{\mathop{S{{O}_{2}}}}\,\rightleftharpoons \underset{Aci{{d}_{2}}}{\mathop{S{{O}^{2+}}}}\,+\underset{Bas{{e}_{1}}}{\mathop{SO_{3}^{2-}}}\,$

Thus, the protonic definition cannot be used to explain the reactions occuring in non-protonic solvents such as COCl₂, SO₂, N₂O₄, etc.

(ii) It cannot explain the reactions between acidic oxides like CO₂, SO₂, SO₃ etc and the basic oxides like CaO, BaO, MgO etc which take place even in the absence of the solvent e.g., CaO + SO₃ → CaSO₄

There is no proton transfer in the above example.

(iii) Substances like BF₃, AlCl₃ etc, do not have any hydrogen and hence cannot give a proton but are known to behave as acids.

(3) Lewis concept : This concept was proposed by G.N. Lewis, in 1939. According to this concept, “a base is defined as a substance which can furnish a pair of electrons to form a coordinate bond whereas an acid is a substance which can accept a pair of electrons.” The acid is also known as electron pair acceptor or electrophile while the base is electron pair donor or nucleophile.

A simple example of an acid-base is the reaction of a proton with hydroxyl ion,

$\[\underset{Acid}{\mathop{{{H}^{+}}}}\,+\underset{Base}{\mathop{:\underset{\centerdot \,\,\centerdot }{\overset{\centerdot \,\,\centerdot }{\mathop{O}}}\,:}}\,\to H:\underset{\centerdot \,\,\centerdot }{\overset{\centerdot \,\,\centerdot }{\mathop{O}}}\,:H$

Some other examples are,

$\[\underset{Base}{\mathop{{{H}_{3}}N:\,}}\,\,\,+\,\,\,\underset{Acid}{\mathop{B{{F}_{3}}}}\,={{H}_{3}}N\to B{{F}_{3}};\,\,\,\underset{Acid}{\mathop{{{H}^{+}}\,}}\,\,+\,\,:\underset{Base}{\mathop{N{{H}_{3}}}}\,={{[H\leftarrow N{{H}_{3}}]}^{+}}\,\,;$ $\[\underset{Acid}{\mathop{B{{F}_{3}}}}\,+\underset{Base}{\mathop{{{[F]}^{-}}}}\,={{[F\to B{{F}_{3}}]}^{-}}$

Lewis concept is more general than the Bronsted Lowry concept.

(i) Types of Lewis acids : According to Lewis concept, the following species can act as Lewis acids.

(a) Molecules in which the central atom has incomplete octet : All compounds having central atom with less than 8 electrons are Lewis acids, e.g., BF₃, BCl₃, AlCl₃, BeCl₂, etc. (such compounds are electron deficient compounds)

(b) Simple cations : All cations are expected to act as Lewis acids since they are deficient in electrons. However, cations such as Na⁺, K⁺, Ca²⁺, etc., have a very little tendency to accept electrons, while the cations like H⁺, Ag⁺, etc., have greater tendency to accept electrons and therefore, act as Lewis acids.

(c) Molecules in which the central atom has empty d-orbitals : The central atom of the halides such as SiX₄, GeX₄, TiCl₄, SnX₄, PX₃, PF₅, SF₄, SeF₄, TeCl₄, etc., has vacant d– orbitals. These can, therefore, accept an electron pair and act as Lewis acids.

(d) Molecules having a multiple bond between atoms of dissimilar electronegativity : Typical examples of molecules falling in this class of Lewis acids are CO₂, SO₂ and SO₃. Under the influence of attacking Lewis base, one π-electron pair will be shifted towards the more electro-negative atom.

(ii) Types of Lewis bases : The following species can act as Lewis bases.

(a) Neutral species having at least one lone pair of electrons : For example ammonia, amines, alcohols, etc., act as Lewis bases because they contain a pair of electrons.

: NH₃, R – NH₃, R – O − O

(b) Negatively charged species or anions : For example; chloride, cyanide, hydroxide ions, etc. act as Lewis base. CN−, Cl⁻, OH⁻

Note : All Bronsted bases are also Lewis bases but all Bronsted acids are not Lewis acids.

(iii) Hard and Soft principle of acids and bases (HSAB principle) : Lewis acids and bases are classified as hard and soft acids and bases. Hardness is defined as the property of retaining valence electrons very strongly. Thus a hard acid is that in which electron-accepting atom is small, has a high positive charge and has no electron which are easily polarised or removed. On the contrary, a soft acid is that in which the acceptor atom is large, carries a low positive charge or it has electrons in orbitals which are easily polarised or distorted.

A Lewis base which holds its electrons strongly is called hard base, e.g., OH⁻, F⁻, H₂O, NH₃, CH₃OCH₃, etc. on the other hand, a Lewis base in which the position of electrons is easily polarised or removed is called a soft base e.g., I⁻, CO, CH₃S⁻, (CH₃)₃P, etc.

In general, hard acids prefer to bind to hard bases and soft acids prefer to bind to soft bases. The bonding between hard acids and hard bases is chiefly ionic and that between soft bases and soft acids is mainly covalent.

Some Hard and Soft Acids and Bases

Hard acids	H⁺, Li⁺, Na⁺, K⁺, Be²⁺, Mg²⁺, Ca²⁺, Sr²⁺, Mn²⁺, Al³⁺, Cr³⁺ Co³⁺, Fe²⁺, La³⁺, Ce³⁺, As³⁺, BF₃, AlCl₃, AIMe₃, SO₃
Soft acids	Pb²⁺, Cd²⁺, Pt²⁺, Hg²⁺, Cu⁺, Ag⁺, Ti⁺, Hg₂²⁺ GaCl₃, B₂H₆, RO⁺, RS⁺, I₂, Br₂ metal atom
Hard bases	H₂O, NH₃, OH⁻, F⁻, Cl⁻, CO₃²⁻ClO₄⁻,NO₃⁻
Soft bases	R₂S, RSH, RS⁻, I⁻, SCN⁻, R₃P, R₃P, R₃As_,CN⁻, RNC, CO, C₂H₄, H⁻, R⁻

(iv) Utility of Lewis concept : Lewis concept is the most general out of all the concepts and can explain the acidic and basic nature of all those substances which could not be explained by the earlier concepts. Similarly, it could explain even those acid-base reactions which could not be explained by the other concepts.

(v) Limitations of Lewis concept

(a) Lewis concept is so general that it considers every reaction forming a co-ordinate bond to be acid-base reaction. But it is not always true for example, according to this concept, even some metals are acids e.g., nickel is acid because it forms the co-ordination compound with Co,[Ni(CO)₄] called tetracarbonyl nickel (0).

(b) The necessary requirement is the formation of a co-ordinate bond between the acid and base. However, the well known acids like HCl and H₂SO₄ do not form any co-ordinate bond and therefore it should not be acids according to this concept.

(c) Arrhenius concept and Bronsted Lowry concept can explain the strengths of acids and bases but Lewis concept cannot.

(d) Acid-base reactions are usually fast but formation of co-ordination compound is slow. Hence it does not fit in the acid-base concept.

(e) The catalytic activity of an acid is due to H⁺(aq) ion. Since the presence of hydrogen is not an essential requirement for a Lewis acid, many Lewis acids will not have this property.

Relative strength of Acids and Bases

In practice K_a is used to define the strength only of those acids that are weaker than H₃O⁺ and K_b is used to define the strength of only those bases that are weaker than OH⁻. For two weak acids HA₁ and HA₂ of ionisation constant K_a1 and K_a2 respectively at the same concentration , we have,

$\[\frac{\text{Acid}\,\,\text{strength}\,\,\text{of}\,\,H{{A}_{\text{1}}}}{\text{Acid}\,\,\text{strength}\,\,\text{of}\,\,H{{A}_{2}}}=\sqrt{\frac{{{K}_{{{a}_{1}}}}}{{{K}_{{{a}_{2}}}}}}$

Similarly, relative strengths of any two weak bases at the same concentration are given by the ratio of the square-roots of their dissociation constants. i.e.,

$\[\frac{\text{Basic}\,\,\text{strength}\,\,\text{of}\,\,B\text{O}{{\text{H}}_{\text{1}}}}{\text{Basic}\,\,\text{strength}\,\,\text{of}\,\,B\text{O}{{H}_{2}}}=\sqrt{\frac{{{K}_{{{b}_{1}}}}}{{{K}_{{{b}_{2}}}}}}$

(1) Relative strength of Inorganic acids

(i) Hydrides

(a) The acidic strength increases with the increase in the electronegativity of the element directly attached with the hydrogen.

H – F > H – OH > H – NH₂ > H – CH₃ ; HCl > H₂S > PH₃ > SiH₄

(b) The acidic strength increases with the increase in atomic size

HF < HCl < HBr ; H₂O < H₂S < H₂Se < H₂Te

(ii) Oxyacids

(a) Among oxyacids of the same type formed by different elements, acidic nature increases with increasing electronegativity,

HOI < HOBr < HOCl ; HIO₃ < HBrO₄ < HClO₄

(b) In oxyacids of the same element, acidic nature increases with its oxidation number

$\[\underset{+1}{\mathop{HOCl}}\,<\underset{+3}{\mathop{HCl{{O}_{2}}}}\,<\underset{+5}{\mathop{HCl{{O}_{3}}}}\,<\underset{+7}{\mathop{HCl{{O}_{4}}}}\,$ ; H₂SO₃ > H₂SO₄ ; HNO₂ < HNO₃

(c) The strength of oxyacids increases from left to right across a period H₄SiO₄ < H₃PO₄ < H₂SO₄ < HClO₄

(d) For the same oxidation state and configuration of the elements, acid strength decreases with increase in size of the atom.

HNO₃ > HPO₃ ; H₃PO₄ > H₃AsO₄ ; HClO₄ > HBrO₄ > HIO₄

(2) Relative strength of organic acids,

(i) A compound is acidic in nature, if its conjugate base can stabilize through resonance. Thus phenol is acidic while ethanol is neutral because the conjugate base of phenol (C₆H₅O⁻) can stabilize through resonance while that of alcohol (C₂H₅O⁻) cannot.

(ii) Hydrogen atom attached to sp-hybridized carbon is more acidic than that on sp² hybridized carbon which in turn is more acidic than that on sp³ hybridized carbon. Thus,

$\[HC\equiv \underset{sp}{\mathop{CH}}\,>C{{H}_{2}}=\underset{s{{p}^{2}}}{\mathop{C{{H}_{2}}}}\,>C{{H}_{3}}-\underset{s{{p}^{3}}}{\mathop{C{{H}_{3}}}}\,$

(3) Relative strength of Inorganic bases

(i) The basicity of a compound decreases with the increase in the electronegativity of the atom holding the electron pair,

$\[\overset{\centerdot \,\,\centerdot }{\mathop{N}}\,{{H}_{3}}>{{H}_{2}}\overset{\centerdot \,\,\centerdot }{\mathop{O}}\,:\,\,>\,HF:$

(ii) The larger size of the atom holding the unshared electrons, the lesser is the availability of electrons.

$\[{{F}^{-}}>C{{l}^{-}}>B{{r}^{-}}>{{I}^{-}}\,\,;\,\,\,{{O}^{2-}}>{{S}^{2-}}$

(iii) Presence of negative charge on the atom holding the electron pair increases the basicity, while the presence of positive charge on the atom holding the electron pair decreases the basicity.

OH^– > H₂O > H₃O⁺

(iv) Among alkali and alkaline earth hydroxides (oxides) the basic nature increases with electropositivity

LiOH < NaOH < KOH < RbOH < CsOH ;

Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂

CsOH is the strongest known base

(v) On going down the group; basic nature decreases with size of the central atom due to decrease in the ability to donate the lone pair.

NH₃ < PH₃ < AsH₃ > SbH₃ > BiH₃

(4) Relative strength of Organic bases

(i) Higher the electron density on nitrogen, more is the basic character of amine.

(ii) A compound is basic in nature, if its conjugate acid can stabilize through resonance. Thus guanidine $\[(N{{H}_{3}}-\overset{N{{H}_{2}}}{\mathop{\overset{|}{\mathop{C=}}\,}}\,NH)$ is as strong alkali as metal hydroxides because its conjugate acid $\[({{H}_{3}}{{N}^{+}}-\overset{N{{H}_{2}}}{\mathop{\overset{|}{\mathop{C=}}\,}}\,NH)$ is very much stabilised through resonance.

9.10 Acid-Base neutralisation & Salts

The reaction between an acid and a base to form salt and water is termed neutralization

$\[HCl(aq)+NaOH(aq)\rightleftharpoons \underset{\begin{smallmatrix} Sodium\,\,Chloride \\ \,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,Salt\end{smallmatrix}}{\mathop{NaCl(aq)}}\,+{{H}_{2}}O(l)\$

$\[{{H}^{+}}(aq)+O{{H}^{+}}(aq)\rightleftharpoons {{H}_{2}}O(l)$

The process of neutralisation does not produce the resulting solution always neutral; no doubt it involves the interaction of H+⁺ and OH⁻ ions. The nature of the resulting solution depends on the particular acid and a particular base involved in the reaction. The following examples illustrate this point when equivalent amounts of acids and bases are reacted in aqueous solution,

A strong acid plus a strong base gives a neutral solution because both are completely ionised and the reaction goes to completion. H⁺ + Cl⁻ + Na + OH⁻ ⇌ H₂O + Na⁺ + Cl⁻
A strong acid plus a weak base gives an acidic solution as the weak base is not completely ionised. The reaction does not go to completion and there is an excess of hydrogen ions in solution.

$\[{{H}^{+}}+C{{l}^{-}}+N{{H}_{4}}OH\rightleftharpoons {{H}_{2}}O+NH_{4}^{+}+C{{l}^{-}}$

A weak acid plus a strong base gives a basic solution as the weak acid is not completely ionised. The reaction does not go to completion and there is an excess of hydroxyl ions in solution.

CH₃COOH + Na⁺ + OH⁻ ⇌ H₂O + CH₃CH₃COO⁻ + Na⁺

A weak acid plus a weak base gives an acidic or a basic or a neutral solution depending in on the relative strengths of acid and base. In case both have equal strength, the resulting solution is neutral in nature.

$\[C{{H}_{3}}COOH+N{{H}_{4}}OH\rightleftharpoons {{H}_{2}}O+NH_{4}^{+}+C{{H}_{3}}CO{{O}^{-}}$

Salts : Salts are regarded as compounds made up of positive and negative ions. The positive part comes from a base while negative part from an acid. Salts are ionic compounds.The salts can be classified into the following classes,

(1) Simple salts : The salt formed by the interaction between acid and base, is termed as simple salt. These are of three types,

(i) Normal salts : the salts formed by the loss of all possible protons (replaceable hydrogen atoms as ) are called normal salts. Such a salt does not contain either replacable hydrogen or a hydroxyl group.

Examples : NaCl, NaNO₃, K₂SO₄, Ca₃(PO₄)₂, Na₃BO₃, Na₂HPO₃ (one H atom is not replaceable as H₃PO₂ is a dibasic acid) NaH₂PO₂ (both H atoms are not replaceable as H₃PO₂ is a monobasic acid), etc.

(ii) Acidic salts : Salts formed by incomplete neutralisation of poly-basic acids are called acidic salts. Such salts still contain one or more replaceable hydrogen atoms. These salts when neutralised by bases form normal salts.

Examples : NaHCO₃, NaHSO₄, NaH₂PO₄, Na₂HPO₄, etc.

(iii) Basic salts : Salts formed by incomplete neutralisation of poly acidic bases are called basic salts. Such salts still contain one or more hydroxyl groups. These salts when neutralised by acids form normal salts.

Examples : Zn(OH)Cl, Mg(OH)Cl, Fe(OH)₂Cl, Bi(OH)₂Cl, etc

(2) Double salts : The addition compounds formed by the combination of two simple salts are termed double salts. Such salts are stable in solid state only.

Examples: Ferrous ammonium sulphate, FeSO₄(NH₄)₂SO₄, 6H₂O

Potash alum, K₂SO₄, Al₂(SO₄)₃, 24H₂O and other alums.

(3) Complex salts : These are formed by combination of simple salts or molecular compounds. These are stable in solid state as well as in solutions.

$\[\underset{Simple\,\,salt}{\mathop{FeS{{O}_{4}}+6KCN}}\,\to \underset{Complex\,\,salt}{\mathop{{{K}_{4}}[Fe{{(CN)}_{6}}]}}\,+{{K}_{2}}S{{O}_{4}}\,\,;$

$\[\underset{Simple\,\,salt}{\mathop{CoS{{O}_{4}}}}\,\,\,+\underset{Molecular\,\,compound}{\mathop{6N{{H}_{3}}}}\,\to \underset{Complex\,\,salt}{\mathop{[Co{{(N{{H}_{3}})}_{6}}]}}\,S{{O}_{4}}$

(4) Mixed salts : The salt which furnishes more than one cation or more than one anion when dissolved in water is called a mixed salt.

Examples :

(1) $\[\begin{align*} & \,\,\,\,\,\,\,OCl \\ & \,\,\,\,\,\,/ \\ & Ca \\ & \,\,\,\,\,\,\backslash \\ & \,\,\,\,\,\,\,\,\,\,Cl \\\end{align*}$

(2) $\[\begin{align*} & Na \\ & \,\,\,\,\,\,\,\backslash \\ & \,\,\,\,\,\,\,\,\,\,S{{O}_{4}} \\ & \,\,\,\,\,\,\,\,/ \\ & \,\,\,\,K \\\end{align*}$

(3) $\[\begin{align*} & \,\,\,\,\,\,Na \\ & \,\,\,\,\,\,\,\,\,\,\,\,\,\backslash \\ & N{{H}_{4}}-P{{O}_{4}} \\ & \,\,\,\,\,\,\,\,\,\,\,\,\,\,\,/ \\ & \,\,\,\,\,\,\,\,\,\,\,H \\\end{align*}$

Ionic Equilibrium : Acids and Bases

Ionic Equilibrium : Acids and Bases

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