Acid-Bases

Introduction:

In general there are three theories as below :

1. Arrhenius Acid Base Theory

2. Bronsted Lowry Theory

3. Lewis Acid Base Theory

1. Arrhenius Acid Base Theory:

Introduction:

In 1884, the Swedish chemist Svante Arrhenius proposed two specific classifications of compounds, termed acids and bases.

When dissolved in an aqueous solution, certain ions were released into the solution.

As defined by Arrhenius, acid-base reactions are characterized by acids, which dissociate in aqueous solution to form hydrogen ions (H+) and bases, which form hydroxide (OH-) ions.

a) Acids:

Acids are defined as a compound or element that releases hydrogen (H+) ions into the solution.

In this reaction nitric acid (HNO3) disassociates into hydrogen (H+) and nitrate (NO3-) ions when dissolved in water.

b) Bases:

Bases are defined as a compound or element that releases hydroxide (OH-) ions into the solution

In this reaction lithium hydroxide (LiOH) dissociates into lithium (Li+) and hydroxide (OH-) ions when dissolved in water.

- Arrhenius acids and bases are divided strong and weak based on its ionisation in water.

Application of Arrhenius Concept

With the help of this concept, we can explain the following:

(i) Aqueous solution of non-metallic oxides (e.g., CO2, SO2, SO3, N2O3, N2O5, P4O6, P4O10 etc.) is acidic, since itgives H+ ions in water.

(ii) Aqueous solution of metallic oxides (e.g., CaO, Na2O etc.,) and the compounds like NH3, N2H4, NH2OH etc are basic, since these substances give OH- ions in water.

CaO + H2 ⇌ Ca(OH)2 ⇌ Ca2+ + 2H-

NH3 + H2O ⇌ NH4OH ⇌ NH4+ + OH-

N2H4 + H2O ⇌ [N2H5]+[OH]- ⇌ N2H5+ + OH-

NH2OH + H2O ⇌ [NH3OH]+[OH]- ⇌ NH3OH+ + OH-

(iii) The strength of an acid (HA) and a base (BOH)can be expressed quantitatively in terms of the ionisation ( ordissociation) constant of the acid and base in aqueous solution.

c) Limitations:

The Arrhenius theory, which is the simplest and least general description of acids and bases, includes acids such as HClO4 and bases such as NaOH or Mg(OH)2.

This theory successfully describes how acids and bases react with each other to make water and salts. However, it does not explain why some substances that do not contain hydroxide ions, for example F- and NO2-, can make basic solutions in water.

The Bronsted-Lowry definition of acids and bases addresses this problem :

In this theory an acid is a substance that can release a proton (like in the Arrhenius theory) and a base is a substance that can accept a proton. A basic salt such as Na+ F- generates OH- ions in water by taking protons from water itself (to make HF):

F-(aq) + H2O(l) ⇌ HF(aq) + OH-

2. Bronsted Concept of Acids and Bases:

Introduction

In 1923, chemists Johannes Nicolaus Bronsted and Thomas Martin Lowry independently developed definitions of acids and bases based on the compounds' abilities to either donate or accept protons (H+ ions).

In this theory, acids are defined as proton donors; whereas bases are defined as proton acceptors.

A compound that acts as both a Bronsted-Lowry acid and base together is called amphoteric.


Advantages of Bronsted-Lowery Concept over Arrhenius Concept:

a) It can explain the basic character of substances like Na2CO3

i.e. which do not contain OH- group and hence were not bases according to Arrhenius concept on the basis that they accept protons.

b) This concept is not limited to molecules but also covers even the ionic species to act as acids or bases.

c) It can also explain the acid-base reactions in the non-aqueous medium.

2.1) Concept of Acid Base:

Bronsted-Lowry theory of acid and bases took the Arrhenius definition one step further, as a substance no longer needed to be composed of hydrogen (H+) or hydroxide (OH-) ions in order to be classified as an acid or base.

For exmaple, consider the below chemical equation:

HCl(aq)+NH3(aq)⟶ NH+4(aq)+Cl-(aq)

- Here, hydrochloric acid (HCl) "donates" a proton (H+) to ammonia (NH3) which "accepts" it , forming a positively charged ammonium ion (NH4+) and a negatively charged chloride ion (Cl-).

- Therefore, HCl is a Bronsted-Lowry acid (donates a proton) while the ammonia is a Bronsted-Lowry base (accepts a proton).

Also, Cl- is called the conjugate base of the acid HCl and NH4+ is called the conjugate acid of the base NH3.

- A Bronsted-Lowry acid is a proton (hydrogen ion) donor.

- A Bronsted-Lowry base is a proton (hydrogen ion) acceptor.

In this theory, an acid is a substance that can release a proton (like in the Arrhenius theory) and a base is a substance that can accept a proton.

A basic salt, such as Na+F-, generates OH- ions in water by taking protons from water itself (to make HF):

F-(aq)+ H2O(l)⟶ HF(aq)+ OH-

When a Bronsted acid dissociates, it increases the concentration of hydrogen ions in the solution, [H+]; conversely, Bronsted bases dissociate by taking a proton from the solvent (water) to generate [OH-].

2.2) Amphoteric compounds:

Some substances can act both as an acid and as a base.

An example is water.

H2O molecules may either donate a hydrogen ion or accept one. This property makes water an Amphoteric Solvent .

In the situation where an acid dissociates in solution, water is acting as a base. Conversely, water acts as an acid when bases dissociate.

The strongest acid we can make in H2O is H+ (aq), and the strongest base we can make in H2O is OH- (aq).

Periodic table showing basic (blue),amphoteric (green) and acidic (red) oxides. The metal-nonmetal boundary is indicated by the gray staircase line.

2.3) Solvent leveling:

Solvent leveling is an effect that occurs when a strong acid is placed in a solvent such as (but not limited to) H2O.

Because strong acids donate their protons to the solvent, the strongest possible acid that can exist is the conjugate acid of the solvent.

In aqueous solution, this is H3O+.

This means that the strength of acids such as HCl and HBr cannot be differentiated in water as they both are dissociated 100% to H3O+.

In the context of our discussion of conjugate bases above, we would say that both Cl- and Br- are spectator ions in water: neither one is a strong enough base to accept a proton from H3O+.

In order to differentiate the acidities of strong acids such as HClO4 and HCl, or the basicities of strong bases such as CH3O- and NH2-, we must typically work in non-aqueous solvents, as explained below.

2.4) Conjugate Acids and Bases

For a reaction to be in equilibrium a transfer of electrons needs to occur. The acid will give an electron away and the base will receive the electron.

Acids and Bases that work together in this fashion are called a ConjugatePair made up of Conjugate Acids and Conjugate Bases.

HA + Z ⇌ A- + HZ+

A stands for an Acidic Compound and Z stands for a Basic Compound

- A Donates H to form HZ+.

- Z Accepts H from A which forms HZ+

- A- becomes conjugate base of HA and in the reverse reaction it accepts a H from HZ to recreate HA in order to remain in equilibrium.

- HZ+ becomes a conjugate acid of Z and in the reverse reaction it donates a H to A- recreating Z in order to remain in equilibrium.

Strength of Bronsted acid and Bronsted bases in aqueous medium

Relative strength of Bronsted acids and Bronsted bases can be determined by treating these substances with water.

When an acid reacts with H2O, water acts as a base in this reaction.

In the reaction between a base and water, water acts as an acid.

Here we shall consider the reaction of HCI (strong acid) and CH3COOH (weak acid) with H2O.

The strength of an acid depends on its ability to transfer its proton (H+) to H2O (which acts as a Bronsted'base) to form its conjugate base.

When a monoprotonic acid like HCI, CH3COOH etc., reacts with H2O, it transfers it proton to H2O to form H3O+ ion and a conjugate base of the monoproton acid.

(i) HCI (strong acid) with H2O (base):

Since HCI is a strong acid, it is nearly 100% ionised in water and hence this reaction proceeds almost completely towards right. This state has been shown by a longer half arrow, pointing towards right.

Thus the reaction of HCl with H2O can be depicted as:

Above reaction the two acids viz., HCI and H3O+ are competing for donating a proton to the base

Since the equilibrium of this reaction lies towards right, HCI donates the proton more strongly than does H3O+ ion.

Thus HCI is a stronger acid than H3O+ ion.

In this reaction the two bases namely H2O and CI- are competing for the gain of a proton from :

Since the reaction proceeds towards right, H2O molecule gains the proton more strongly than does CI- ion.

Thus H2O is stronger base than CI- ion, i.e., H2O is a strong base and CI- ion is a weak base.

(ii) CH3COOH (weak acid) with H2O (base) :

Since CH3COOH is a weak acid, it is less than 1 % ionised in H2O and hence this reaction proceeds mostly towards left.

This state of affairs is represented using a longer half arrow, pointing towards left.

Thus, the reaction of CH3COOH with H2O can be shown as:

In the above reaction, the two acid viz. CH3COOH and H3O+ are competing for donating a proton to the base:

Since the equilibrium of the reaction lies towards the left, H3O+ ion donates the proton more strongly than does CH3COOH molecule.

Thus H3O+ ion is stronger acid than CH3COOH, i.e., H3O+ ion is strong acid and CH3COOH is weak acid.

In the above reaction, the two bases, namely H2O molecule and CH3COO- ion, are competing for the gain of a prolon from the acid :

Since the reaction proceeds towards left, CH3COO- ion gains the proton more strongly than does H2O.

Thus CH3COO- ion is stronger base than H2O i.e. CH3COO- ion is a strong base and H2O is a weak base.


Hydrated Cations Acids

The species which, in aqueous solution, furnish a proton by dissociation of water molecule bound to a metal ion are called hydrated cation acids.

Hydrated iron (H) ion, for example, will release a proton in accordance with the following reaction:

[Fe (H2O)6]2+ + H2O ⇌ [Fe(H2O)5(OH)+] + H3O+

The proton is lost from the hydrated cation, since the iron (Il) ion tends to displace the lone pairs on oxygen of attached water towards it self.

This induces the oxygen to withdraw the electron pair that it shares with hydrogen towards itself.

As a result, the oxy en-hydrogen bond gets weakened and the proton dissociates.

Smaller the size of and/or high the charge on the metal ion in the hydrated cation, more strong is the acid.

A smaller cation, with a higher charge can displace the lone pairs on oxygen of water more towards itself and, in consequence, facilitate the release of proton more efficiently than a larger cation with a lower charge.

Therefore, the hydrated Iron(Ill) ion is a stronger acid than hydrated iron(Il) ion.

Hydrated aluminium(Ill) cation is a more strong acid than hydrated gallium(III) cation

[Fe(H2O)6]3+ > [Fe(H2O)6]2+ ;

[Co(H2O)6]3+ > [Co(H2O)6]2+ ;

[A](H2O)6]3+ > [Ga(H2O)6]3+

2.5) Limitations of the Bronsted-Lowry concept:

a) It cannot explain the reactions between acidic oxides like CO2, SO2, SO3 etc. and basic oxides like CaO, BaO, MgO, etc. which take place even in the absence of the solvent.

b) Substances like BF3, AlCl3, etc. behave as acids but they do not have protons to lose or donate.

Solvent system concept or auto-ionisation :

Notice the similarity with autoionisation (ie. autoprotolysis) of water.

In the solvent-system definition, a solute which increases the concentration of the cation generated by autoionisation of the solvent is called an acid, and a solute which increases the concentration of the anion generated by autoionisation of the solvent is called a base.

Auto-ionisation of water:

Water (H2O) under-goes self-ionisation of the following three ways:

Auto-ionisation of liq. NH3 :

Different mode of auto-ionisation of liq. ionisation of liq NH3, given above, indicate that, according to solvent system concept, in liq. NH3, any substance that gives NH2- ions (which are solvent cations) will act as an acid in liq. NH3, while that which produces NH-2 or NH2- or NH3- ions ( which are solvent cations) will behave as a base in liq. NH3.

The compounds which give ions NH4+ in liq. NH3 are called ammono acid, while those which give NH-2, NH2- or NH3- ions in this solvent are called ammono base.

Example of ammono acids:

Example of amino bases:

Auto-ionisation of liq. SO2

Self-ionisation of liquid. SO2 takes place as follows:

Relation between the ion given by the auto-ionisation of H2O, liq. NH3 and Iiq. SO2

The equations representing the auto ionisation of H2O, liq. NH3 and liq. SO2 indicate that NH4+ and SO2+ ions which are obtained by the self-ionisation liq. NH3 and liq. SO2 respectively, are analogous to H3O+ (or H+) ions, which are obtained by the ionisation of H2O, similarly NH2- and NH2- ion, obtained by the ionisation of liq. NH3 and SO32- ions (obtained from liq SO2). Are analogous to OH- ions obtained from H2O, N3- ions obtained from liq. NH3 are equivalent to O2- ions obtained from H2O.

Auto-ionisation of liq. BrF3:

Liq. BrF3 has high specific conductance and hence undergoes auto-ionisation to produce BrF+2 and Br-4 ions.

2BrF3 ⇌ BrF+2 + BrF-4

Auto-ionisation of liq. HF:

The high specific conductance of liq. HF suggests a relatively high degree of auto - ionisation of liq. HF shown below:

Thus any substances that can give H2F+ ions would beahave as an acid liq. HF and any substance that can furnish F- or HF-2 ions would act as a base in this solvent.

Example :

(i) CH3COOH (Which acts as a weak acid in water) acts as a base in liq. HF, because its gives F- ions ( solvent anions-base ions) when dissolve in liq. HF

CH3COOH ( Base ) + HF ( Solvent ) ⟶ CH3COOH+2 + F- [Solvent anions ( base ions )]

(ii) H2SO4 and HNO3 both acts as strong acids and aqueous medium, but show basic character in liq. HF, due to the production of F- ions which are solvent anions (base ions)

HNO3( Base ) + HF ( solvent ) ⟶ H2NO+3 [ Solvent anions ( base ions ) ]

NHO3( Base ) + HF ( Solvent ) ⟶ H2NO+3 + F- [ Solvent anions ( base ions ) ]

(iii) HCIO4 behaves as the strongest acid in water, but in liq. HF it acts as an amphoteric substance, since it produces both H2F+ ( solvent cation – acid ions ) and F- ( solvent anions-base ions) in this solvent.

HCIO4 + HF ( Solvent ) ⟶ H2NO+3 + F- [ Solvent anions ( base ions ) ]

HCIO4 + HF ( Solvent ) ⟶ H2F+ + CIO-4 [ Solvent anions ( acid ions ) ]

Self-ionisation of some other non-aqueous solvent:

The self-ionisation of liq. N2O4, CH3 COOH, CHN and H2SO4 in shown below:

(a) 2N2O4 (solvent) ⇌ 2NO+ (nitrosyl ions : solvent cation – acid ions ) + 2NO-3 (nitrate ions : solvent anions – base ions)

(b) 2CH3COOH (solvent) ⇌ CH3COOH2+ (solvent cation : acid ions) +CH3COO- (solvent anion : base ions)

(c) 2HCN (solvent) ⇌ H2CN+ (solvent cation : acid ions) + CN- (solvent anion : acid ions)

(d) 2H2SO4 (solvent) ⇌ H3SO4+ (solvent cation : acid ions) +HSO-4 (solvent anion : base ions)

Applications of solvent system concept

To explain the acidic/basic nature of a given substance in a given solvent.

With the help of solvent-system concept of acids and bases, we can predict whether a given substance will behave as an acid, as a base, as an amphoteric substance or as a neutral substance in a given solvent.

The following examples illustrate this application:

Behaviour of CH3COOH in H2O, liq. NH3, liq, HF and H2SO4

(a) In water CH3COOH ionises to a small extent to produce H3O+ ions (solvent cations-acid ions).

Due to the feeble ionisation of CH3COOH in water, the concentration of H3O+ ions obtained is very low and hence, according to solvent system concept, CH3COOH acts as a weak acid in water

The longer half arrow pointing towards left indicates that CH3COOH undergoes partial ionisation in water or in other words the above equilibrium lies mosly towards the left

(b) When dissolved in liq. NH3, CH3COOH is completely converted into NH4+ ions which are solvent cations or acid ions.

CH3COOH, therefore, behaves as a strong acid in liq NH3

(c) When CH3COOH is dissolved in liq. HF, it ionises to produce F- ions (solvent anions-base ions) and hence behaves as a base in liq. HF

(d) When CH3COOH is dissolved in H2SO4, the concentration of ions HSO4- (solvent anions base ions) is increased and hence CH3COOH behaves as a base in H2SO4

It follows from the above discussion that according to solvent system concept CH3COOH acts as a weak acid in water, in liq. NH3 it behaves as a strong acid but in liq. HF and H2SO4 both it shows basic character.

Behaviour of H2SO4 in water and liq. HF

H2SO4 (Strong acid) + 2H2O (Solvent) ⟶ 2H2O (Solvent) 2H3O+ + SO2-4 [Solvent (base ions) ]

H2SO4 (Base) + HF (Solvent) ⟶ H3SO+4 [Solvent anions (base ions) ]

HClO4 (Strong acid) + H2O (Solvent) ⟶ H3O+ + ClO-4 [Solvent cations (base ions)]

HClO4 + HF (Solvent) ⟶ H3OClO+4 + HF [Solvent anions (base ions)]

(iii) Behaviour of HCIO4 in water and liq. HF :

(a) HCIO4 ionises completely in H2O and produces H3O+ ions (solvent cations-acid ions) in high concentration HCIO4, therefore, acts as a strong acid in aqueous solution

HCIO4(strong acid) + H2O(solvent) ⟶ H3O+ + CIO4- [solvent cations (acid ions)]

When HCIO4 is put into liq. HF, it is ionised to produce both H2F+ (which are solvent cations -acid ions) and F- (which are solvent anions-base ions) and hence, according to solvent system concept, HCIO4 behaves as an amphoteric substance in liq. HF

HCIO4 + HF (solvent) ⟶ H3CIO4+ + F- [Solvent anions (base ions)]

HCIO4 + HF (solvent) ⟶ H2F+ + CIO4- [Solvent anions (acid ions)]

BEHAVIOUR OF DIFFERENT SUBSTANCE AND DIFFERENT SOLVENTS

3) The Lewis Theory of Acids and Bases

Introduction

The Bronsted Theory defines acids and bases as proton donors and acceptors.

Around the same time that Johannes N. Bronsted and Thomas M. Lowry came up with their theory of acids and bases, Gilbert N. Lewis proposed his own theory.

While the Bronsted-Lowry theory is based on the transfer of protons, Lewis' theory is based on the transfer of electrons.

3.1) Concept:

A Lewis Acid is a substance that can accept a pair of electrons to form a new bond.

They are sometimes referred to as electrophiles, or seekers of an additional electron pair.

A Lewis Base is a substance that can donate a pair of electrons to form a new bond.

They are sometimes referred to as nucleophiles, or seekers of a positive nucleus.

Neutralization is the sharing of an electron pair between an acid and base.

The product formed in a neutralization reaction is sometimes called as adduct or complex.

3.2) Lewis Acid and Bases

According to the Lewis theory, an acid is an electron pair acceptor, and a base is an electron pair donor.

Lewis bases are also Bronsted bases; however, many Lewis acids, such as BF3, AlCl3 and Mg2+, are not Bronsted acids.

The product of a Lewis acid-base reaction, is a neutral, dipolar or charged complex, which may be a stable covalent molecule.

As shown at the top of the following drawing, coordinate covalent bonding of a phosphorous Lewis base to a boron Lewis acid creates a complex in which the formal charge of boron is negative and that of phosphorous is positive.

In this complex, boron acquires a neon valence shell configuration and phosphorous an argon configuration.

If the substituents (R) on these atoms are not large, the complex will be favoured at equilibrium.

However, steric hindrance of bulky substituents may prohibit complex formation.

The resulting mixture of non-bonded Lewis acid/base pairs has been termed "frustrated", and exhibits unusual chemical behaviour.

In the first example, an electron deficient aluminium atom bonds to a covalent chlorine atom by sharing one of its non-bonding valence electron pairs, and thus achieves an argon-like valence shell octet.

Because this sharing is unilateral (chlorine contributes both electrons), both the aluminium and the chlorine have formal charges, as shown.

If the carbon chlorine bond in this complex breaks with both the bonding electrons remaining with the more electronegative atom (chlorine), the carbon assumes a positive charge. We refer to such carbon species as carbocations.

Carbocations are also Lewis acids, as the reverse reaction demonstrates.

Example: The interaction between a magnesium cation (Mg2+) and a carbonyl oxygen is a common example of a Lewis acid-base reaction.

The carbonyl oxygen (the Lewis base) donates a pair of electrons to the magnesium cation (the Lewis acid).

Relative order of Lewis acid character of BF3, BCl3, BBr3 and BI3 molecules

The measurement of the molar heats of formation of the solutions of BX3 molecules in nitrobenzene and the dipole moments of their adducts like H3N ⟶ BX3 have shown that the Lewis acid character of the given molecules is in the order

BF3 < BC13 < BBr3 < BI3

This order shows that BBr3 will give more stable adducts than BCl3 and BC13 will yield more stable adducts than BF3.

The above order can be explain on the basis of boron halogen π-back bonding.

We know that each of the given molecules has a trigonal planar geometry which arises due to sp2 hybridisation of B-atom in its excited state.

In sp2 hybridistion, one 2p orbital, say 2pz of B-atom, remains unhybridised and vacant.

The halogen atom has valence-shell configuration as ns2 npx2 npy1 npz2.

The singly filled npy orbital of halogen atom overlaps with the singly filled sp2 hybrid orbital of B-atom and forms B-X σ-bond, while each of the remaining three orbitals of halogen atom contains one lone pair of electrons.

Thus the structure of BF3 molecule can be shown following figure.

Now the filed 2pz orbital of B-atom makes a lateral overlap with the vacant 2Pz orbital of B-atom and gives rise to the formation of an additional F ⟶ B π-bond, called Pπ - Pπ back bond

The formation of Pπ - Pπ back bonding results in the following:

(a) Due to the formation of extra F ⟶ B π-bond, B - F bond aquires some double bond character .

Thus B - F bond should be shown as :

Since any one of the three F -atoms can take part in the formation of B - F bond, the structure of BF3 can be supposed to be a resonance hybrid of the above three equivalent resonating structures.

(b) In case of BF3 molecule F ⟶ B π-bond results by 2p (F) - 2p (B) overlap while in case of BCL3, BBr3 and Bl3 molecules this bond is produced by 3p (Cl) -2p (B), 4p (Br) - 2p (B) and, 5p (l) -2p (B) overlaps respectively.

Since in case of BF3 molecule the overlap takes place between the orbitals which have the same energy and shape their overlap occurs with great effectiveness.

As we move from F to I, since the difference in energy between 3p (Cl)-2p (B); 4p (Br) - 2p (B) and 5p (I) - 2p (B) pairs of orbitals increases, the effectiveness with which these pair of orbitals overlap decreases.

Thus we see that the tendency of back bonding is maximum in BF3 molecule.

This tendency falls rapidly on passing from BCL3 to BI3.

This means that the tendency of BF3 molecule to accept electron pair given by Lewis base

(e.g, NH3, PH3, F- etc.) is minimum and this tendency increases as we move from BF3 to BI3 i.e., the Lewis acid character of BX3 molecules is in the order

BF3 < BCl3 < BBr3 < Bl3

3.3) Limitations of Lewis theory:

a) This theory cannot explain the strength of acids and bases.

b) Generally neutralisation reactions are instantaneous (very fast) but Lewis acid base reaction go slowly.

c) All acid base reactions do not involve co-ordinate bond formation.

d) H+ ion as a catalyst (in some reactions) cannot be explained by this theory.