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Showing posts with label ACIDS AND BASES. Show all posts
Showing posts with label ACIDS AND BASES. Show all posts

Jan 8, 2019

Soft and Hard Acids and Bases (SHAB) Principle

SHAB Principle and Stability of the Complex:

According to Lewis Concept an Acid-Base reaction involves an interaction of a vacant orbital of an acid (A) and a filled or unshared orbital of a base (B).
A + :B
A : B 
Lewis Acid
(Acceptor)

Lewis Base
(Donor)

Adduct or Complex
The species A is called Lewis Acid or a generalized acid and B is called Lewis Base or a generalized base. A strong acid A and a strong base B, will form the stable complex A : B.
A concept known as Principle of Soft and Hard Acids and Bases(SHAB) Principle is very helpful in making a stability of the complex A : B.
According to this principle the complex A : B is most stable when A and B are either both soft or both hard.
The complex is least stable when one of the reactants (namely A and B) is very hard and the other one is very soft. 
In order to arrive at a comparative estimate of the donor properties of different bases, the preferences of a particular base to bind a proton H+ and methyl mercury (II) ion, [CH3HgB]+ was determined. Both the proton and methyl mercury cation can accommodate only one coordinate bond but the two cation vary widely in their preferences to bases. This preferences was estimated from the experimental determination of equilibrium constants for the exchange reactions:
BH+ + [CH3Hg(H2O)]+
[CH3HgB]+ 
+ H3O+
The results indicate that bases in which the donor atom is N, O or F prefer to coordinate to the proton. Bases in which the donor atoms is P, S, I, Br, Cl or C prefer to coordinate to mercury.

Hard and Soft Bases:

The donor atoms in the first group have high Electronegativity, low Polaris ability and hard to oxidize. Such donors have been named ‘hard bases by Pearson, since they hold on to their electrons strongly. 
The donor atoms of second category are of low electronegativity, high polarisability, and are easy to oxidize. Such donors have been called ‘soft’ bases since they are hold on to their valence electrons rather loosely.
In simple terms hardness is associated with a tightly held electron shell with little tendency to polarise. On the other hand softness is associated with a loosely bound polarisable electron shell.
It will be seen that within a group of the periodic table softness of the Lewis bases increases with the increase in size of the donor atoms. 
Thus, among the halide ions softness increases in the order:
-Cl -Br -I -
Thus F - is the hardest and I - is the Softest base.
Classification of Bases
(R = alkyl or aryl group)
Hard Borderline
Soft
H₂O, OH⁻, F⁻, CH₃COO⁻, PO₄⁻³, SO₄⁻², Cl⁻, CO₃⁻², ClO₄⁻, NO₃⁻, ROH, RO⁻, R₂O, NH₃, RNH₂, N₂H₄
C₆H₅NH₂, C₅H₅N, N₃⁻, Br⁻, NO₂⁻, SO₃⁻²,N₂
R₂S, RSH, RS⁻, I⁻, SCN⁻, S₂O₃⁻², R₃P, R₃As, (RO)₃P, CN⁻, RNC, CO, C₂H₄, C₆H₆, H⁻

Hard and Soft Acids:

After having gone through a classification of bases, a classification of Lewis acids is necessary. The preferences of a given Lewis acid towards ligands of different donor atoms is usually determined from the stability constant values of the respective complexes or from some other useful equilibrium constant measurements. When this is done, metal complexes with different donor atoms can be classified into two sets based on the sequences of their stabilities.
Hard acids have small acceptor atoms, are of high positive charge and do not contain unshared pair of electrons in their valence shell, although all these properties may not appear in one and the same acid. These properties lead to high electronegativity and low polarisability. In keeping with the naming of the bases, such acids are termed as 'Hard'Acids.
Hard Acids:
NPOSFClBrI
Soft acids have large acceptor atoms, are of low positive charge and contain ushered pairs of electrons in their valence shell. These properties lead to high polarisability and low electronegativity. Again in keeping with the naming of the bases, such acids are termed 'Soft' Acids
Soft Acids:
NPOSFClBrI
Classification of Lewis Acids as Hard, Intermediate and Soft Acids.
Hard Borderline
Soft
H⁺, Li⁺, Na⁺, K⁺, Be⁺², Mg⁺², Ca⁺², Sr⁺², Mn⁺², Al⁺³, Ga⁺³, In⁺³, La⁺³, Lu⁺³, Cr⁺³, Co⁺³, Fe⁺³, As⁺³, Si⁺⁴, Ti⁺⁴, Zr⁺⁴, Th⁺⁴, U⁺⁴, Ce⁺³, Sn⁺⁴, VO⁺², UO₂⁺², MoO⁺³, BF₃, AlCl₃, SO₃, Cr⁺⁶
Fe⁺, Co⁺², Ni⁺², Cu⁺², Zn⁺², Pb⁺², Sn⁺², Sb⁺³, Bi⁺³, Rh⁺³, B(CH₃)₃, SO₂, NO⁺, GaH₃
Cu⁺, Ag⁺, Au⁺, Tl⁺, Hg⁺, Pd⁺², Cd⁺², Pt⁺², Hg⁺², CH₃Hg⁺, Tl⁺³, BH₃, GaCl₃, InCl₃, I⁺, Br⁺, I₂, Br₂, Zerovalent Metal atoms

Classify the following as Hard and Soft Acids and Bases. 
(i) H⁻ (ii) Ni⁺⁴ (iii) I⁺ (iv) H⁺

(i) The hydride ion has a negative charge and is far too large in size compared to the hydrogen atom. Its electronegativity is quite low and it will be highly polarisable by virtue of its large size. Hence it is Soft Base.
(ii) Quadrivalent nickel has quite a high positive charge. Compared to bivalent nickel its size will be much smaller. Its electronegativity will be very high and polarisability will be low. Hence it is Hard Acid.
(iii) Mono-positive iodine has a low positive charge and has a large size. It has a low electronegativity and a high polarisability. Hence it is a Soft Acids.
(iv) H⁺ has the smallest size with a high positive charge density. It has no unshared pair of electrons in its valence shell. All these will give a high electronegativity and very low polarisability. Hence H⁺ is a Hard Acid.

Soft and Hard Acids and Bases (SHAB) Principle:

This principle also means that if there is a choice of reaction between an Acid and two Bases and two Acids and a Base,
A Hard Acid will prefer to combine with a Hard Base and a Soft Acid will prefer to combine with Soft Base and thus a more stable product will be obtained.
Hard Acid - Hard Base may interact by strong ionic forces. Hard Acids have small acceptor atoms and positive charge while the Hard Bases have small donor atoms but often with negative charge. Hence a strong ionic interaction will lead to Hard Acid - Base combination.
On other hand a Soft Acid - Soft Base combination mainly a covalent interaction. Soft Acids have large acceptor atoms, are of low positive charge and contain ushered pair of electrons in their valence shell.

Application of Soft and Hard Acids and Bases (SHAB) Principle:

(i) [CoF₆]⁻³ is more stable than [CoI₆]⁻³
It will be seen that Co⁺³ is a hard acid
F⁻ is a hard base and I⁻ is a soft base
Hence [CoF₆]⁻³ (Hard Acid + Hard Base) is more stable 
than [CoI₆]⁻³ (Hard Acid + Soft Base).

AgI₂⁻ is stable, but AgF₂⁻ does not exist. Explain.

We know that Ag⁺ is a soft acid, F⁻ is hard base and I⁻ is soft base
Hence AgI₂⁻ (Soft Acid + Soft Base) is a stable complex and AgF₂⁻ (Soft Acid + Hard Base) does not exist.
(ii) The existence of certain metal ores can also be rationalised by applying SHAB principle. Thus hard acids such as Mg⁺², Ca⁺² and Al⁺³ occur in nature as MgCO₃, CaCO₃ and Al₂O₃ and not as sulphides (MgS, CaS and Al₂S₃), since the anion CO₃⁻² and O⁻² are hard bases and S⁻² is a soft base.
Soft acids such as Cu⁺, Ag⁺ and Hg⁺² on the other hand occurs in nature as sulphides.
The borderline acids such as Ni⁺², Cu⁺² and Pb⁺² occur in nature both as carbonates and sulphides. 
The combination of hard acids and hard bases occurs mainly through ionic bonding as in Mg(OH)₂ and that of soft acids and soft bases occurs mainly by covalent bonding as in HgI₂.

Explains why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not?

In the case of Hg(OH)₂ and HgS, Hg is a soft acid and OH⁻ and S⁻² is hard base and soft base respectively. Evidently HgS (Soft acid + Soft base) will be more stable than Hg(OH)₂ (Soft Acid + Hard base). More stability of HgS than that of Hg(OH)₂ explain why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not. 
If we arrange the donor atoms of most common Lewis bases in an increasing order of their electronegativity values we get,
Soft and Hard Acids and Bases and Hard Soft Acid Base Principle
Donor Atoms of Most Common Lewis Bases





Jan 3, 2019

Acids and Bases Questions and Answers


What is the acidic order of oxoacids of phosphorus?

When the oxidation state rule applied to the oxoacids of phosphorus (H₃PO₂, H₃PO₃ and H₃PO₄) it is predicted that the acidic Character of these acids should be in the order:
H₃PO₂ㄑH₃PO₃ㄑH₃PO₄, but the experimental observation suggested the reverse order:
H₃PO₂ ≥ H₃PO₃〉H₃PO₄.
The experimental order can be explained when we consider the structures of these acids given as,
Acids and Bases Questions and Answers
Oxoacids-of-Phosphorus
H₃PO₂ is a mono-basic acid. The proton attached to an oxygen has a far greater chance of dissociation than any directly bonded hydrogen. the structure of H₃PO₂ involves one protonated oxygen and another unprotonated oxygen. H₃PO₃ is dibasic and hence has two protonated oxygen's and one unprotonated oxygen. H₃PO₄ is tribasic, has three protonated oxygen's and one unprotonated oxygen. 
In this series therefore the number of unprotonated oxygen's, which are the vehicles for the enhancement of acidity, is the same for all the three acids. But dissociable protons increase from one in H₃PO₂ to three in H₃PO₄. Therefore the overall inductive effect of the unprotonated oxygen decreases from H₃PO₂ to H₃PO₄. Hence the acidity slightly falls off in the order:
H₃PO₂ ≥ H₃PO₃H₃PO₄

pH of a solution is 4.5. Calculate the concentration of H⁺ ion.

We have from the definition,
pH = -log[H⁺] =4.5
or, log[H⁺] = -4.5
∴ [H⁺] = 3.16×10⁻⁵

Calculate the [H⁺], [OH⁻] and pH of a solution prepared by diluting 20 ml of 0.1M HCl to one lit. 

[H⁺] = (20×0.1)/1000 
=0.002 
= 2×10⁻³
[OH⁻] = (1×10⁻¹⁴)/[H⁺]
=(1×10⁻¹⁴)/(2×10⁻³) 
= 0.5×10⁻¹¹
pH = -log[H⁺] 
= -log2×10⁻³ 
= 3-log2 
= 2.7

Calculate the [H⁺], [OH⁻] and pH of a solution obtained by dissolving 28 gm KOH to make 200 ml of a solution.

[OH⁻] = (28×1000)/(56×200) 
= 2.5M 
(Molecular Weight of KOH = 56gm)
[H⁺] = (1×10⁻¹⁴)/[OH⁻] 
= (1×10⁻¹⁴)/(2.5) 
= 4×10⁻¹⁵ 
pH = -log[H⁺] 
= - log4 × 10⁻¹⁵ 
= (15 - log4) 

What will be the effect of adding KNH₂ to liquid ammonia with respect to acidity?

2NH₃ NH₄⁺ + NH₂⁻
KNH₂ K⁺ + NH₂⁻
Due to common ion effect the equilibrium shift to left and decrees the acidity of NH₃.

Arrange in order of Lewis acidity :
(i) BF₃ (ii) BCl₃ (iii) BI₃ (iv) BBr₃

These boron halides have pi-interaction between filled p- orbitals of the halogen and empty p-orbital of boron. The effectiveness of the pi-interaction falls off with increasing size of halogens so that it is the strongest in BI₃. When boron halides receives an electron pair the pi-bond between boron and halogen has to be ruptured in order to make room for a coordinate bond.
H₃NBF₃
Thus with BF₃ it will be hardest to rupture the pi-bond. Therefore the order of the Lewis acidity is:
BF₃BCl₃BBr₃BI₃

Why all alkali are bases but all bases are not alkali?

All the bases are dissolved in water to produce alkali whereas all the bases are not dissolved in water but all the alkali are dissolved in water. Thus all alkali are bases but all bases are not alkali.
Na₂O is a base because it dissociate in water to produce NaOH. NaOH, KOH and Ca(OH)₂ dissolved in water to produce OH⁻ ion. Thus all these hydroxide are Alkali.
NaOH Na⁺ + OH⁻
KOH K⁺ + OH⁻
Ca(OH)₂ Ca⁺² + 2OH⁻
Al(OH)₃, Fe(OH)₃, Zn(OH)₂ etc does not dissolved in water but reacts with acids to produce Salt and water. Thus These are bases but but not alkali.





Explain - NH₃ behaves as a base but BF₃ as an Acid.

In NH₃, the central N atom have lone pair of electrons This lone pair coordinate to empty orbital, is termed as a base according to Lewis concept. 
The compounds with less than an octet for the central atom are Lewis acids. In BF₃B in BF₃, contain six electrons in the central atom thus it behaves as an acid.
F₃B + :NH₃ F₃BNH₃
Acid Base Adduct


Explain why tri-covalent phosphorus compounds can serve both as Lewis acids and also as bases?

Tri-covalent phosphorus compounds like PCl₃ have a lone pair of electrons in phosphorus. This lone pair may coordinate to a metal ion thus allowing the compound to serve as a Lewis base. 
Ni + 4PCl₃   [Ni(PCl₃)₄]⁰
Again the quantum shell of phosphorus has provision for d-orbitals which can receive back donated electrons from electron reach low oxidation state of a metal ion. In this latter case tri-covalent phosphorus compound serves as a Lewis acid.

Bi-positive tin can function both as a Lewis acid and a Lewis base. Explain?

The Lewis representation of SnCl₂ shows a lone pair on tin through it is as yet short of an octet. Ligands, particularly donor solvents with lone pairs, may coordinate to tin giving complexes. Here SnCl₂ behaves as a Lewis acid.
Again interaction of platinum group metal compounds with SnCl₂ (SnCl₃⁻) as donor leads to the formation of coordinate complexes. As for examples,
[(Ph₃P)₂PtCl(SnCl₃)], [RuCl₂(SnCl₃)₂]⁻². These are examples of Lewis base behavior of SnCl₂.

Can SiCl₄ and SnCl₄ function as Lewis acids?

Both silicon and tin are members of Group IV of the periodic table and their quantum shells admit of d-orbitals. As a result they can expand their valence shell through SP³d² hybridisation and can give rise to six coordinate complexes. In fact complex like [SiCl₄{N(CH₃)₃}₂], [SiCl₂(Py)₄]Cl₂, [SnCl₄(Bpy)] are known. Thus SiCl₄ and SnCl₄ can function as Lewis acids.

Name the conjugate acids and the conjugate bases of HX⁻ and X⁻².

Conjugate acid of a species is the one that is obtained on the addition of a proton and conjugate base of a species is one that is obtained on the release of a proton.
H₂O + HX⁻   OH⁻ + H₂X
Acid₁ Base₂ Base₁ Acid₂
In the above reaction HX⁻ acts as a base and its conjugated acid is H₂X.
HX⁻ + H₂O   X⁻² + H₃O⁺
Acid₁ Base₂ Base₁ Acid₂
In this reaction HX⁻ acts as an acid thus its conjugated base is X⁻².
In the same way conjugate acid of X⁻² is HX⁻ but X⁻² cannot have any conjugated base because there is no proton that can release.

Arrange these oxides in order of their acidic nature : N₂O₅, As₂O₃, Na₂O, MgO.

Acidic oxides react with water to give oxoacids. The higher the oxidation number and the higher the electronegativity the grater the central element will force to reaction with water to give the oxoacid. Of all these oxoacids nitrogen has the highest oxidation number and electronegativity. Thus the order of acidic nature and oxidation number are-
Acidic nature :
Na₂OMgOAs₂O₃N₂O₅
Oxidation number of Oxide forming element: +1  +2  +3  +5. 

Bisulphate ion can be viewed both as an acid and a base. Explain.

Bisulphate ion is HSO₃⁻. It may lose a proton to give the conjugate base SO₃⁻², thus behaving as an acid. Again it may add on a proton to give the conjugate acid, thus showing its base character.

SO₃ Behaves as an acid and H₂O as a base. Explain.

SO₃ like BF₃ has less than an octet,and will be termed as an acid according to Lewis concept.
But in H₂O oxygen atom contain lone pairs to donate the Lewis acid.
Oxygen and sulphur contain six electrons in their valence shell and therefore regarded as Lewis Acids. The oxidation of SO₃⁻² to SO₄⁻² ion by oxygen and S₂O₃⁻² ion by sulphur are the acid-base reactions.
SO₃⁻²  [O]   ⇆   [OSO₃]⁻²
SO₃⁻² +  [S]       [SSO₃]⁻²

Explains why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not?

In the case of Hg(OH)₂ and HgS, Hg is a soft acid and OH⁻ and S⁻² is hard base and soft base respectively. Evidently HgS (Soft acid + Soft base) will be more stable than Hg(OH)₂ (Soft Acid + Hard base). More stability of HgS than that of Hg(OH)₂ explain why Hg(OH)₂ dissolved readily in acidic aqueous solution but HgS does not. 

Classify the following as Hard and Soft Acids and Bases.
(i) H⁻ (ii) Ni⁺⁴ (iii) I⁺ (iv) H⁺

(i) The hydride ion has a negative charge and is far too large in size compared to the hydrogen atom. Its electronegativity is quite low and it will be highly polarisable by virtue of its large size. Hence it is Soft Base.
(ii) Quadrivalent nickel has quite a high positive charge. Compared to bivalent nickel its size will be much smaller. Its electronegativity will be very high and polarisability will be low. Hence it is Hard Acid.
(iii) Mono-positive iodine has a low positive charge and has a large size. It has a low electronegativity and a high polarisability. Hence it is a Soft Acids.
(iv) H⁺ has the smallest size with a high positive charge density. It has no unshared pair of electrons in its valence shell. All these will give a high electronegativity and very low polarisability. Hence H⁺ is a Hard Acid.

Dec 23, 2018

Acids and Bases

Acids and Bases:

Acids:

Any of the class of substances whose aqueous solutions are characterized by the sour taste, the ability to turns blue litmus red, and the ability to react with bases and certain metals to form salts.

Bases:

In chemistry, a base is a substance that in aqueous solution is slippery to touch, that's bitter, changed the color of the indicator, (turns red litmus to blue) react with acid to form salts.
The concepts of acid and base is developed one after another tend to make the definitions more and more broad based.
We need to familiar with the following five concepts :
1. The Arrhenius Concept.
2. The Solvent System Concept.
3. The Protonic Concept.
4. The electronic Theory.

1. The Arrhenius Concept :

Arrhenius was one of the early exponent of electrolytic dissociation theory to define acids and bases.
His classification of acids and bases was based on the theory that, acids when dissolved in water dissociate hydrogen ions and anions where as bases when dissolved in water dissociate into hydroxyl ions and cations. Thus HCl is an acid and NaOH is a base and the neutralization process can be represented by a reaction involving the combination of H⁺ and OH⁻ ions to form H₂O.
Acid :   HCl H⁺ + Cl⁻ 
Base :   NaOH Na⁺ + OH⁻
Acid Base Equilibrium: H⁺ + OH⁻  H₂O

Acid base neutralization reaction in aqueous medium :

The Arrhenius concept was highly useful to explaining acid base reaction in aqueous medium.
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Representation of Acid Base Neutralization Reaction
The heat liberated during neutralization of strong acid (HCl, HClO₄, HNO₃, HBr, HI and H₂SO₄) and strong base (NaOH, KOH, RbOH, and Ca(OH)₂) is the same, namely 13.4 kcal/mole (56 KJ/mole) indicating that the ion participating in the reaction must be the same (that is H⁺ and OH⁻) and that other ions (Na⁺, K⁺, Rb⁺, Ca⁺², Cl⁻, Br⁻, I⁻, ClO₄⁻ or SO₄⁻²) take no part.
This can be explain as follows.Let, 
HA(aq)+BOH(aq) BA(aq)+H₂O
Where, HA and BOH are strongly electrolytes and completely dissociate in the Solution.
Hence, (H⁺+A⁻)+( B⁺+OH⁻)(B⁺+A⁻)+ H₂O
Cancelling likes ions, we have,
H⁺+OH⁻  H₂O
Since neutralization of all strong acids and bases reduced to the formation of 1 mole water from H⁺ ion and OH⁻ ion.The enthalpy change (ΔH) will also be same.

Limitation of Arrhenius Concept:

(i) According to this concept, HCl is regarded as an acid only when dissolved in water and not in some other solvent such as benzene or when it exists in the gaseous state.
(ii) It cannot account for the acidic and basic character of the materials in non aqueous solvents, as for example, NH₄NO₃, in liq NH₃ acts as an acid, through it does not give H⁺ ions. Similarly many organic materials in NH₃ , which does not give OH⁻ ions at all, are actually known to show basic character.
(iii) The neutralization process limited to those reactions which can occurs in aqueous solutions only, although the reactions involving salt formation do occur in many other solvents and even in the absence of solvents.
(iv) It cannot explain the acidic character of certain salts such as AlCl₃ in aqueous solution.

2. The Solvent system Concept:

The non- aqueous solvent molecules may also dissociate into two oppositely charged ions.
Examples:
(i)We consider the solvent H₂O, its characteristic cation and anion are H⁺ and OH⁻ respectively are, 
H₂O H⁺ + OH⁻
Since we know that a bare proton will readily polarize other anions or molecules we write a H⁺ as H₃O⁺ indicating that it is a solvated proton that exists in the solution.
So that the overall dissociation of solvent will be:
H₂O + H₂O H₃O⁺ + OH⁻
Thus all those compounds which can give H₃O⁺ ions in H₂O will act as an acids and all those compounds which can give OH⁻ ions in H₂O will behave as bases.
(ii) Similarly liquid NH₃ can written as 
NH₃ + NH₃ NH₄⁺ + NH₂⁻
Thus those compounds which give NH₄⁺ ions in liquid NH₃ will act as an acids and all those compounds which can give NH₂⁻ ions in liquid NH₃ acts as bases.
Thus the dissociation (or autoionisation) of non-aqueous solvents is directly responsible for the nature of the chemical reactions that can be initiated in such solvents.
According to solvent system concept,
An acid is a substance which by dissociation in the solvent forms the same cation as does the solvent itself due to auto-ionization.
A base is one that, gives on dissociation in the solvent the same anion as does the solvent itself on its ionization.
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Acid-Base Neutralisation Reaction According to Solvent System Concept
The auto-ionisation of some protonic and non-protonic solvents are, Protonic Solvents:
H₂O + H₂O H₃O⁺ + OH⁻
Acid
Base
Acid
Base
NH₃ + NH₃ NH₄⁺ + NH₂⁻
Acid
Base
Acid
Base
CH₃COOH + CH₃COOH CH₃COOH₂⁺ + CH₃COO⁻
Acid
Base
Acid
Base
Non-Protonic Solvents:
SO₂ + SO₂ SO⁺² + SO₃⁻²
Acid
Base
Acid
Base
BrF₃ + BrF₃ BrF₂⁺ + BrF₄⁻
Acid
Base
Acid
Base
N₂O₄ + N₂O₄ 2NO⁺ + 2NO₃⁻
Acid
Base
Acid
Base
Neutralization Reactions according to Solvent System Concept: 
Just as with the Arrhenius definition, neutralisation is a reaction between an acid and a base to produce salt and solvent. 
Neutralisation of some non-aqueous solvents are, 
(i) Neutralisation reaction in liquid NH₃: 
Dissociation of Solvent:    
NH₃ + NH₃  NH₄⁺ + NH₂⁻
Dissociation of Acid:         
NH₄Cl NH₄⁺ + Cl⁻ 
Dissociation of Base:         
KNH₂  K⁺ + NH₂⁻
Acid Base Neutralization reaction: 
NH₄Cl + KNH₂ KCl + 2NH₃
Acid
Base
Salt
Solvent
NH₄Cl(Acid) + KNH₂(Base)  KCl(Salt) + 2NH₃(Solvent)
(ii) Neutralization reaction in liquid SO₂:
Dissociation of Solvent:
SO₂SO₂  SO⁺² + SO₃⁻²
Dissociation of Acid:
SOBr₂ SO⁺² + 2Br⁻ 
Dissociation of Base: 
K₂SO₃ 2K⁺ + SO₃⁻²
Acid Base Neutralization reaction:
SOBr₂ + K₂SO₃ 2KBr + 2SO₂
Acid
Base
Salt
Solvent
Arrhenius concept of acid and base and Solvent System concept of Acids and Bases
Neutralization Reactions according to Solvent System Concept

Utility of the Concept:

Evidently this concept of solvent system can be used to explain the acid-base reactions occurring in aqueous and non aqueous solvents(protonic or non-protonic both).
Thus this theory can simply be said to be an extension of the Arrhenius Theory.
Limitations:
(i) This theory does not consider a number of acid-base reactions included in the protonic definition.
(ii) It limits acid-base phenomena to solvent system only. Thus it does not explain the acid base reactions which may occurs in the absence of solvent.
(iii) It can not explain the neutralisation reactions occurring without the presence of ions.