Theories of acids and Bases

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Classical or Functional Definitions of Acid and Base:

  • Acid: An acid is defined as a substance whose water solution has sour taste, turns blue litmus to red, can neutralize base and evolves hydrogen gas when treated with active metals like Zn, Mg, Na etc.  e.g. HCl, HNO3, H2SO4, CH3COOH
  • Base: A base is defined as a substance whose water solution has a bitter taste, has soapy touch, turns red litmus to blue and can neutralize an acid. e.g. NaOH, KOH, NH4OH
  • The word alkali is used to water-soluble bases.

Need of Conceptual Definitions of Acids and Bases:

  • The classical definitions of acids and bases were based on some observed properties of acids and bases.
  • These definitions were unable to explain the structure responsible for their properties. Hence there was need of conceptual definitions of acids and bases.

Chemists Acids and Bases

Arrhenius concept of Acids and Bases:

  • In 1887,  Arrhenius, the Swedish chemist, proposed theory of ionization to account for the properties of the aqueous solution of electrolytes.
  • According to this concept,
  • Acid: An acid is defined as a hydrogen-containing compound which gives hydrogen ions  (H+ ) in its aqueous solution.
    • Examples:

 HCl(aq)     ⇌   H+(aq)  +      Cl(aq)

CH3COOH(aq)   ⇌  CH3COO(aq)    + H+(aq)

In general, an equilibrium for all acids exists as,



HA(aq)   ⇌     H+(aq)  +    A(aq)

  • Base: A base is defined as a hydroxide compound which gives hydroxyl (OH) ions in its aqueous solution.
    • Examples :

NaOH(aq)     ⇌   Na+(aq)  +      OH(aq)

NH4OH(aq)     ⇌   NH4+(aq)  +      OH(aq)

  • In general, an equilibrium for all bases exist as,

BOH(aq)   ⇌     B+(aq)  +    OH(aq)



  • Neutralisation: Neutralisation reaction is the reaction  in which the acid and base react together to produce salt and water

Consider a reaction between strong acid like HCI and strong base like NaOH.

HCl   +    NaOH    → NaCl  +  H 2 O

(Acid)    (base)           (salt)        (water)

  • By ionic (Arrhenius) theory, HCI, NaOH and NaCI dissociate into their ions in an aqueous medium.

H+(aq) + Cl(aq)  + Na+(aq)  + OH(aq)   →  Na+(aq) +  Cl(aq) +   H2O

By canceling the common ions of both sides, net equation is,



H+(aq)   +    OH(aq)    →        H2O

  • Thus in neutralization, H+ ions of acid combine with OH ions of the base forming an unionized water molecule. Thus by Arrhenius theory, A process in which H+ ions of an acid combine with OH- ions of an alkali to form unionized water molecule is called as neutralization.


 Notes:

  • Properties of acid are due to properties of H+ ions present in the solution.
  • Strong acids are highly ionized aqueous solution producing a large number of H+ ions or protons.
  • Weak acids are very little ionized and produce a very small number of protons or H+ ions.
  • Properties of bases are due to the presence of  OH ions present in the solution.
  • A strong base is highly ionized and gives a large number of OH ions.
  • A weak base is very little ionized and gives very few OH ions.
  • Bases which are highly soluble in water are known as alkalies.

Advantages of Arrhenius Theory:

  • Arrhenius concept is used to explain,
    • acid-base properties of substances in an aqueous medium
    • neutralization, hydrolysis and
    • the strength of acids and bases.

Limitations of Arrhenius Theory:

  • Acids and bases are defined in terms of their aqueous solution and not in terms of the substances themselves. Hence this theory is applicable to aqueous solutions only and not applicable to non-aqueous and gaseous reactions.
  • It is applicable only to compounds having formula HA for acids or BOH for bases. Thus the theory is unable to explain acidic properties of CuSO4, AlCl3, CO2, SO2 as they cannot be represented by the formula HA. Similarly, the theory is unable to explain the basic properties of Na2CO3, amines, pyridine, NH3 as they cannot be represented by the formula BOH.
  • The theory does not consider the role of solvent in deciding the nature of acid and base. Thus HCl is strong acid when dissolved in water but it is weak acid when dissolved in benzene.
  • This theory doesn’t explain acidic property of HCl and basic property of NH3 in a nonaqueous medium like benzene, acetone or in the gaseous state.
  • According to  Arrhenius theory, proton (H+) exist free in aqueous solution.  However, in aqueous solution, H+ ion is always hydrated and exist as hydronium ion (H3O+).
  • By Arrhenius theory neutralization process in which H+ ions of an acid combine with OH ions of an alkali to form unionized water. Thus the theory is unable to explain the neutralization reaction between HCl(g) and NH3(g)  not involving combination of H+ and OH ions

HCl(g) +    NH3(g)   →     NH4Cl(g)

Bronsted- Lowry Concept of Acid and Base:

  • In 1923, scientists Bronsted and Lowry proposed more general definitions of acids and bases to overcome the limitations of Arrhenius theory.  This concept is independent of solvent and is also called as  Protonic Concept. According to Bronsted Lowry concept
  • Acid:  An acid is defined as a substance (molecule or ion) which has a tendency to donate one or more protons  ( H+) to other substance.  Thus acid is proton donor species. e.g.  Molecules like HCl, HNO3, H2SO4, H2O ions like HSO4, H3O+, HCO3, NH4+  etc.
  • Base: A base is defined as a substance (molecule or ion) which has a tendency to accept one or more protons  (H+) from other substance. Thus base is proton acceptor species. e.g. Molecules like  NH3,  RNH2, H2O ions like CH3COO, OH, HSO4, Cl–  etc.
  • Acid Base Reaction (Neutralisation) :

HCl(g) +    NH3(g)   →     NH4+ +    Cl

(acid)      (base)             (acid )     (base)



 

In the above reaction, HCl and NH4+  are the acids as both donate a proton.

While NH3  and  Cl are the bases as they accept a proton.

Concept of Conjugate Acid-Base Pair:

  • When an acid donates a proton , the remaining part of it has a tendency to regain proton.  Therefore it acts as a base which is called as linked or conjugate base.

Acid1    ⇌  Base1    +      H+

Similarly



Base2   +   H+  ⇌      Acid2

Acid1 and Base1 as well as Acid2 and Base2 differ by a proton and are called as conjugate acid-base pairs.

  • In an acid-base reaction, pairs of substances which differ by a proton and which can be formed from one another by the mutual gain or loss of a proton are called as conjugate acid-base pairs. e.g.

HCl(g) +    NH3(g)  ⇌     NH4+ +    Cl

(acid1)      (base2)             (acid2)     (base1)



Strength of Acid and Base on the Basis of Bronsted- Lowry Concept:

  • The strength of a base is measured in terms of its ability to capture proton while the strength of an acid is measured in terms of ability to donate a proton. It is obvious that stronger the acid weaker its conjugate base and stronger the base weaker is its conjugate acid.

HCl     ⇌    H+    +         Cl

(strong acid )           ( weak conjugate base)

Neutralization Reaction on the Basis of Lowry and Bronsted Concept:

  • According to this theory, a neutralization reaction is a reaction in which conjugate base and conjugate acid are formed from reacting base and acid respectively.

H2SO4    + H2O   ⇌  H3O+ +  HSO4

(acid1)      (base2)             (acid2)     (base1)

 

Amphoteric Nature of Water:

  • A substance which can act as an acid, as well as a base, is called an amphoteric substance. Thus by Bronsted -Lowry concept, a substance which has a capacity both to accept or donate protons is called amphoteric substance or amphiprotic substance.
  • Water is an amphoteric substance.  Water can accept a proton and can act as a base, as well as it can donate a proton and can act as an acid.  Dual nature of water depends upon the nature of other substance with which it is treated.
  • Water acts as a base when treated with the strong acid like HCl.

HCl     +   H2O     ⇌     H3O+    +    Cl



(acid1)      (base2)             (acid2)     (base1)

  • Water acts as an acid when treated with base stronger than itself like NH3.

H2O       +      NH3     ⇌      NH4+    +     OH

(acid1)      (base2)             (acid2)     (base1)

Advantages of Bronsted – Lowry Concept:

  • This theory made more general definitions of acids and bases.
  • This concept is applicable to aqueous as well as nonaqueous solutions.
  • This concept is independent of solvent.

Limitations of Bronsted -Lowry Concept: 

  • Bronsted – Lowry concept can’t explain certain acid-base reactions which do not involve proton.


Lewis Concept of Acids and Bases:

  • In 1923, scientist Gilbert Lewis proposed a more general concept of acids and bases. This concept is based on the electronic theory of valency.
  • Acid: Lewis acid is defined as any species (molecule, atom or ion) which can accept a lone pair of electrons to form a coordinate bond.
    • Thus Lewis bases are electrophiles.
    • Examples: All cations like Al3+ , Mg2+ , Cu2+ , H+ , Zn2+, Fe2+, Ag+ etc. Neutral molecules like AlCl3, BF3, BeCl2, SO3, SiF4, SnCl2 etc. In these neutral molecules, the central atom has only six electrons around it and they have empty d orbitals. Atoms like S and O. In these atoms the atom has only six electrons in its valence orbit.
  • Base: Lewis base is defined as any species (molecule, atom or ion) which can donate a lone pair of electrons to form a coordinate bond.
    • Thus Lewis bases are nucleophiles.
    • Examples: Molecule like NH3, H2O, Amines etc, All anions like SO4,  Cl , Br, O2- etc.
  • Neutralisation: By Lewis acid-base theory the process of neutralisation is simply the formation of a coordinate bond between the electron donor and electron acceptor.
    • Consider the reaction between boron trifluoride BF3 and ammonia NH3. In BF3 boron atom contains six electrons in its final orbit thus boron has an incomplete octet. Thus boron is capable of accepting a lone pair of electrons, while nitrogen in NH3  has a lone pair of electrons.
    • During reaction between them ‘N’ atom in NH3 donate a lone pair of electrons to Boron and acts as Lewis base, while  BF3 accepts the lone pair of electrons from ammonia and BF3  behaves like Lewis acid.

 



Acids and Bases

Note:

  • All Bronsted bases are also Lewis bases but all Bronsted acids are Lewis acids but reverse is not true.
  • Lewis base is defined as any species (molecule, atom or ion) which can donate a lone pair of electrons to form a coordinate bond. while according to Bronsted Lowry theory a base is anything that donates a pair of electrons to an acidic hydrogen. A Lewis base is anything that donates a pair of electrons, while a Bronsted base is anything that donates a pair of electrons to an acidic hydrogen. Thus the Lewis and Bronsted  Lowry definitions of bases are identical. Thus All Bronsted bases are also Lewis bases.
  • Bronstead Lowry concept defines acid as a substance (molecule or ion) which has a tendency to donate one or more protons  ( H+) to other substance. Thus according Bronstead Lowry concept an acid is hydrogen containing compound. A Lewis acid is anything that accepts a pair of electrons, while a Bronsted acid accepts pairs of electrons at an acidic hydrogen. Lewis acid may or may not contain hydrogen. Hence  all Bronsted acids are Lewis acids but all Lewis acids are not Bronsted Lowry acids.

Acidic Nature of Boron Trifluoride:

Acids and Bases 02

  • In BF3 boron atom contains six electrons in its final orbit thus boron has an incomplete octet. Thus boron is capable of accepting a lone pair of electrons, During the reaction, BF 3  accepts the lone pair of electrons and behaves like Lewis acid.

Basic Nature of Ammonia:

Acids and Bases 03

  • Nitrogen in NH3  has a lone pair of electrons. During reaction  ‘N’ atom in NH3 donate a lone pair of electrons and acts as Lewis base. Hence NH3  is Lewis base.

Notes:

  • Compounds in which central atom has expanded octet act as Lewis acid using vacant d orbitals.

SnCl4  +   2 Cl  →   [SnCl6]2-

Lewis acid                   Lewis Base

SiF4   +      2 F    →   [SiF6]2-



Lewis  acid                    Lewis base

  • All simple cations which have vacant valency orbitals act as a Lewis acid.

Cu2+  +     4 : NH3     →      [Cu(NH3)4]2+

Lewis acid        Lewis base            Complex

Zn2+ +     4: OH     →      [ Zn(H2O) 4 ]2+

Lewis acid       Lewis base             Complex

Ag +     +    2 : NH3       →    [Ag (NH3)2]+

Lewis acid       Lewis base               Complex

Fe3+    +    6 CN       →  [Fe(CN)6}3-

Lewis acid       Lewis base             Complex

  • Elements with electron sextet that is having six electrons in their valence shell, acts as Lewis acid.

O    +         SO32-  →    [O ¬ SO3]2-

Lewis acid       Lewis base

Sulphite ion       Sulphate ion

  • Electron-deficient molecules act as Lewis acid.  BF3, AlCl3, SO3 are electron deficient molecules containing central B, Al and S atom respectively having only six electrons. They complete their octet by forming coordinate bond..

O    +         SO32-  →    [O ¬ SO3]2-

Lewis acid         Lewis base

Sulphite ion       Sulphate ion

Limitations of Lewis Concept of Acids and Bases:

  • All acid-base reactions do not involve co.ordinate bond formation.
  • Lewis concept no don explain the behavior of well known protonic acids like HCl, H2SO4, etc. which do not form coordination bonds with bases. Therefore, according to Lewis, these are not regarded as acids.
  • This theory fails to explain relative strength of acids and bases.
  • Actually, the formation of a coordination compound is a slow process, but the acid-base reaction is a fast process. This behaviour cannot be explained on the basis of the Lewis concept.
  • The catalytic activity of many acids is due to proton (H+). Lewis acids do not have proton hence Lewis acids do not possess catalytic property.
Science > Chemistry > Ionic EquilibriaYou are Here
Physics Chemistry  Biology  Mathematics

3 Comments

  1. I am very thankful of your act to make it

  2. I'm very grateful for this write-up: is so comprehensive. Nonetheless, the limitations of the Lewis Concept wasn't exemplified here. And the stipulation that "not all Bronsted-Lowry acids were Lewis acids wasn't elucidated.

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