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

Nov 1, 2018

Protonic Concept (Bronsted -Lowry Concept)

Protonic Concept (Bronsted -Lowry Concept) :

A great advancement was made in the area of acids and bases when Bronsted and Lowry proposed in 1923 a new concept which is independent of solvents.
According to them, an acid is defined as any hydrogen containing materials (a molecule or a cation or an anion) that can release a proton (H⁺) to any other substance.
Whereas a base is a substance (a molecule or a cation or an anion) that can accept a proton from any other substance.
In short, an acid is a proton donor and a base is proton acceptor.
A neutralization process is therefore involved a release of proton by acid and the acceptance of proton by base. Neutral compounds or even ions could be designated as acids or bases according to this concept.
Describe Bronsted Lowry Theory of Acids and Bases with Examples
Bronsted Acids and Bases

Conjugate Acid - Base Pairs:

When an acid release a proton (H⁺), the residue must be a base and this base can take up a proton (H⁺) to form original acid. Thus acid base neutralization involved two acid or two bases, thus forming conjugated pairs. 
Thus, a conjugate base of an acid is that part left after the proton lost. Similarly the conjugate acid of base is the species formed on the addition of proton to the base.
Describe Bronsted Lowry Theory of Acids and Bases with Examples
Conjugate Acid - Base Pairs
In the above reaction HCl donates a proton to H₂O and is therefore an acid. H₂O on the other hand accepts a proton from HCl and is therefore a base.
In the reverse reaction which at equilibrium proceeds at the same rate as the foreword reaction, H₃O⁺ ion donates a proton to Cl⁻ ion, hence H₃O⁺ ion is an acidCl⁻ 
ion, because it accepts a proton from H₃O⁺ ion, is a base.
HCl + H₂O  ⇆  H₃O⁺ + Cl⁻
The members of which can be formed from each other mutually by the gain or loss of proton are called conjugated acid - base pairs.
Describe Bronsted Lowry Theory of Acids and Bases with Examples
Acid - Base Chart Containing Some Common Conjugate Acid - Base Pairs
A conjugate base of an acid is that part left after the proton is lost. Similarly the conjugate acid of a base is the species formed on the addition of a proton to a base.
An acid exhibits its acid properties only when it is allowed to react with a base. Similarly a base displays its basic properties only when it exposed to an acid.

Examples of Acid - Base Reactions:

NH₄⁺ + OH⁻ NH₃ + H₂O
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
HCl + NH₃ Cl⁻ + NH₄⁺
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
HNO₃ + H₂O NO₃⁻ + H₃O⁺
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
H₂SO₄ + H₂O HSO₄⁻ + H₃O⁺
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
NH₄⁺ + H₂O NH₃ + H₃O⁺
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
H₂O + F⁻ OH⁻ + HF
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
H₂O + CH₃NH₂ OH⁻ + CH₃NH₃⁺
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)
NH₄⁺ + CH₃COO⁻ NH₃ + CH₃COOH
(Acid₁)
(Base₂)
(Base₁)
(Acid₂)

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.





Fate of Proton in Aqueous Medium

Fate of Proton in Aqueous Medium :

The proton plays key role in the acid-base functions. In aqueous medium it is usually represented as H₃O⁺
A naked hydrogen ion has a vanishingly small size ( radius ~10⁻¹³ cm =10⁻¹⁵ A) and therefore has a very high (charge/radius) ratio(~10⁵ and is expected to be the most effective in polarizing other ions or molecules according to Fajan’s rules. In H₃O⁺ there are assumed to coordinate bonds from water oxygen to the proton, thus giving the proton a helium electronic configuration.

Experimental evidence of such formulations:

Perchloric acid (HClO₄) reacts vigorously with water and It gives a series of hydrates which are :
HClO₄ ( M.P. 112°C)
HClO₄, H₂O ( M.P. +50°C)
HClO₄, 2H₂O ( M.P. -17.8°C)
 HClO₄, 3H₂O ( M.P. -37°C)
HClO₄, 3.5H₂O ( M.P. -41.4°C)
Of these hydrates the most remarkable is the mono hydrate, melting at the much higher temperature then the covalent anhydrous acid. It is very stable and can be heated to around 100°C without decomposition. The mono-hydrate is about ten times viscous as the anhydrous acid. It has the same crystal lattice as the ionic ammonium perchlorate, showing that it is a ionic compound, [H₃O⁺][ClO₄].

Water as an Acid and as a Base :

We know that water dissociates weakly to H⁺ and OH⁻ ions. Regardless of what other ions are present in water, there will always be equilibrium between H and OH ions.
H₂O H⁺ + OH⁻
The proton, however, will be solvated and is usually written as [H₃O⁺] . For simplicity we will write H⁺ only. The above equilibrium will have its own equilibrium constant:
K=([H⁺]×[OH⁻])/[H₂O]
or, K×[H₂O]=[H⁺]×[OH⁻]
The square brackets indicate concentrations. Recognizing the fact that in any dilute aqueous solution, the concentration of water molecules (55.5 moles / liter) greatly exceed that of any other ion, [H₂O] can be taken as a constant. Hence,
K×[H₂O]=KW=[H⁺]×[OH⁻]
Where KW is the dissociation constant of water (ionic product of water). 
The value of H⁺ in pure water has been determined as 10⁻⁷ M, so that KW becomes,
KW=[H⁺]×[OH⁻
10⁻⁷×10⁻⁷ = 1.0×10⁻¹⁴ M 
Definition of pH and pOH in Aqueous Solution Mathematical Relationship between pH and pOH
Representation of Dissociation Constant of Water
The above relation tells us that in aqueous solution the concentration of H⁺ and OH⁻ are inversely proportional to each other. 
If H⁺ concentration increases 100 fold, that of OH⁻ has to decrease 100 fold to maintain KW constant. 

Dissociation of H₂O into H⁺ and OH⁻ ions:

Dissociation of water into H⁺ and OH⁻ ions are an endothermic reaction.
Endothermic Reaction: The Reactions in which heat is absorbed by the system from the surroundings are known as endothermic reactions.
H₂O+13.7 kcal H⁺+OH⁻

Le-Chatelier's Principle: 

If a system is in equilibrium, a change in any factors that determine the condition of equilibrium will cause the equilibrium to shift in such way as to minimize the effect of this change. Thus according to Le-Chatelier’s principle, increasing temperature will facilitate dissociation, thus giving higher values of KW. The value of KW at 20°C, 25°C and 60°C are 0.68×10⁻¹⁴, 1.00×10⁻¹⁴ and 9.55×10⁻¹⁴ respectively.

The Concept of pH:

The dissociation of water, KW, has such low value that expressing the concentrations of H⁺ and OH⁻ ions of a solution in terms of such low figures is not much convenient and meaningful. Such expressions necessarily have to involve negative power of the base 10. Sorensen proposed the use of a term known as pH , defined as:
pH = -log10[H⁺]
= -log101/[H⁺]
Thus for a solution having H⁺ concentration 10−¹M has a pH=1.
And for a solution having H⁺ concentration 10⁻¹⁴M has a pH =14.
For such solutions having H⁺ concentration in the range of 10⁻¹ M to 10⁻¹⁴ M is more convenient and meaningful to express the acidity in terms of pH rather then H⁺ concentrations. The use of small fractions or negative exponents can thus be avoided. For a mono basic acid molarity and normality are the same while they are different for poly-basic acid.
Examples:
0.1M H₂SO₄ is really 0.2N H₂SO₄ 
So that, pH = -log₁₀⁡[H⁺
= -log₁₀⁡(0.2) = 0.699

Calculation of pH:

Solution:
Since HCl is strong electrolyte and completely dissociated.
Thus, [H⁺] = [HCl
= 0.002M = 2×10⁻³M
So, pH = -log[H⁺
= -log(2×10⁻³) = (3-log2) =2.7 
(ii) For 0.002M H₂SO₄ solution:
For H₂SO₄, [H⁺
= 2[H₂SO₄
= 2×0.002 = 4×10⁻³M
Thus, pH = -log[H⁺
= -log(4×10⁻³) 
= (3 - log4) = 2.4
(iii) For 0.002M Acetic Acid Solution:
For Acetic Acid [H⁺
= √(Ka×[CH₃COOH
√(2×10⁻⁵×2×10⁻³) 
= 2×10⁻⁴
Thus, pH = -log[H⁺
= -log2×10⁻⁴ 
= (4 - log2) = 3.7

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

The Concept of pOH :

The corresponding expression for the hydroxide ion is,
pOH = - log[OH⁻]

It is follows from these relations that the lower the pH, the solution is  acidic.
If the acidity of a solution goes down 100 fold its pH goes up by two units.
Example:
A solution of pH = 1 has [H⁺] which is 100 times greater than that of pH = 3. Taking the case of OH⁻ ions, the pOH will go down by two unit(from 13 to 11).
The product of [H⁺] and [OH⁻] is 10⁻¹⁴ and that this has to remain constant. 
Thus, [H⁺] [OH⁻] = 10⁻¹⁴
or, log[H⁺] [OH⁻] = log10⁻¹⁴
or, log[H⁺] + log[OH⁻] = -14
or, -log[H⁺] - log[OH⁻] = 14
or, pH + pOH = 14
We have from the definition, pH= - log[H⁺] and pOH = -log [OH⁻].

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
Difference Between Acidic, Basic and Neutral solution on the basis of Con-centration:
We can now proceed to differentiate between neutral, acidic or basic on the basis of relative concentrations of H⁺ and OH⁻ ions on the one hand, and on the basis of pH on the other. 
A neutral solution is one in which the concentrations of H⁺ and OH⁻ ions are equal.
Thus, [H⁺] = [OH⁻] = 10⁻⁷M

An acidic solution is one in which concentration of H⁺ exceeds that of  OH⁻:

[H⁺][OH⁻]
or, [H⁺] 〉10⁻⁷M and [OH⁻]〈 10⁻⁷M


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)
An basic solution is one in which concentration of OH⁻ exceeds that of H⁺ : 
[H⁺][OH⁻]
or, [OH⁻] 〉10⁻⁷M and [H⁺]〈 10⁻⁷M
In terms of pH we have the following relations:
Neutral Solution [H⁺] = 10⁻⁷M or, pH = 7 
[H+] pH [OH-] pOH
10-7 7 10-7 7
Acidic Solution :  [H⁺] 〉10⁻⁷M or, pH〈 7.0 
[H+] pH [OH-] pOH
100 0 10-14 14
10-1 1 10-13 13
10-2 2 10-12 12
10-3 3 10-11 11
10-4 4 10-10 10
10-5 5 10-9 9
10-6 6 10-8 8
Basic Solution : [H⁺]〈10⁻⁷M or, pH 〉7 
[H+] pH [OH-] pOH
10-8 8 10-6 6
10-9 9 10-5 5
10-10 10 10-4 4
10-11 11 10-3 3
10-12 12 10-2 2
10-13 13 10-1 1
10-14 14 100 0
Mathematical definition of pH provides a negative value when [H⁺] exceeds 1M. However pH measurements of such concentrated solutions are avoided as these solutions are not likely to be dissociated fully. Concentration of such strongly acid solutions is best expressed in terms of molarity than in terms of pH.








Electronic Theory of Acids and Bases

Lewis Concept:

Each of the concepts had its own limitations. The need was therefore felt for a concept which can cover all the earlier aspects of acids and bases. Lewis, (1923) explains the acid-base phenomenon not in terms of ionic reactions but in terms of electronic structure of the acid and base along with the formation of the coordinate covalent bond.
A base was to be considered according to Lewis as an electron pair donor and an acid as an electron pair acceptor. 
An acid therefore combines with a base leading to a coordinate link from the base to the acid.
An acid must have a vacant orbital into which an electron pair donated by a base can be accommodated. A base, then, is a substance which has at least one unshared (lone) pair of electrons.
Lewis Acid and Base Definition and Examples
Lewis Representation of Acid and Base
According to Lewis theory, the process of neutralisation is simply the formation of a coordinate bond between an acid and a base. The neutralisation product, termed as coordinate complex or adduct, may be either non-ionisable or may undergoes dissociation or condensation reaction depending on its stability.
 
H+ + :NH3 [H  :NH3]+
Acid
Base
Adduct
 
F3B + :NH3 F3B :NH3
Acid
Base
Adduct

Bronsted Concept Vs Lewis Concept:

According to Bronsted Concept, a base is any molecules or ion which accepts proton. Lewis Concept recognises that the base is an electron pair to donate to H⁺. So that during neutralization a coordinate bond is formed between the donor atom of the base and the proton.
H+ + :NH3 [H  :NH3]+
Acid
Base
Adduct
Thus all Bronsted bases are bases in Lewis sense also.

Arrhenius Concept Vs Lewis Concept:

According to Arrhenius, an acid when dissolved in water dissociate into hydrogen ions and anions, that is it release hydrogen ions (proton). Since the proton can receive electron pair from a base. Thus all the Arrhenius acids are also acids under Lewis definition.

Solvent System Concept Vs Lewis Concept:

The solvent system concept recognizes that ammonium chloride should behave as an acid in liquid ammonia. It may be noted that ammonium ion is but an ammoniated proton and the proton can accept a lone pair of electron from a base. Therefore an ammonium ion in liquid ammonia is also an acid in Lewis sense.
The solvent system considers amide ion is a base in liquid ammonia.The amide ion is capable of donating a lone pair to an acid, hence this is a base in Lewis sense.
The reaction between the ammonium ion and amide ion in liquid ammonia is represented in Lewis sense as follows:
NH₃ + [H⁺ NH₂⁻]  NH₄⁺  + NH₂⁻

Classification of Lewis Acids and Bases:

A wide variety of compounds are recognized as acids and bases according to Lewis. Any Lewis acid must contain at least one empty orbital in the valence shell of one of its atoms to accept an electron pair from a Lewis base.

(i) Molecules containing a central atom with an incomplete octet:

Typical examples of this class of acids are electron deficient molecules such as alkyl and halides of Be, B and Al
Some reactions of this types of Lewis acid with Lewis base are,
 
F₃B + O(C₂H₅)₂ F₃BO(C₂H₅)₂
Acid
Base
Adduct
 
Cl₃Al + NC₅H₅ Cl₃AlNC₅H₅
Acid
Base
Adduct
 
Me₃B + N₂H₄ Me₃BN₂H₄
Acid
Base
Adduct

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
Acid
Base
Adduct
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.

(ii) Molecules Containing a Central atom with Vacant d-Orbitals:

The central atom of the halides such as SiX₄, GeX₄, TiCl₄, SnX₄, PX₃, PF₅, SF₄, SeF₄, TeCl₄, etc. have vacant d-orbitals. These substances can, therefore, accept an electron pair from the Lewis base to accommodate in their vacant d-orbital and can thus form adducts with a number of halide ions and organic bases. These substances is thus Lewis acids.
These halides are vigorously hydrolyzed the appropriate HX. The hydrolysis reaction take place through the formation of an unstable adduct(intermediate) with H₂O.

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.

(iii) Simple Cations:

Theoretically all simple cations are potential Lewis acids. Reactions of some cations as Lewis acids with Lewis base are,
 
Lewis acid + Lewis base Adduct 
Ammonation: 
Ag⁺ + 2(:NH₃) [NH₃AgNH₃]⁺ 
Acid
Base
Adduct
Cu⁺² + 2(:NH₃) [Cu(NH₃)₄]⁺² 
Acid
Base
Adduct
Hydration:
Co⁺³ + 6(:OH₂) [Co(OH₂)₆]⁺³ 
Acid
Base
Adduct
Alcoholation:
Li⁺ + :OHCH₃ [LiOHCH₃] 
Acid
Base
Adduct
The Lewis acid strength or coordinating ability of the simple cations are increases with-
(a) An increases in the positive charge carried by the cation.
(b) An increases in the nuclear charge for atoms in any period of the periodic table.
(c) A decreases in ionic radius.
(d) A decreases in the number of the shielding shells.
Evidently the acid strength of simple cations increases for the elements on moving from left to right in a period and from bottom to top in a group on the periodic table.
Examples:
Strength of Lewis acid increasing:
Fe⁺²Fe⁺³
( Positive charge increases from +2 to +3)
K⁺Na⁺ 
(On moving from bottom to top in a group)
Li⁺Be⁺² 
(On moving from left to right in a period)

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. 

(iv) Molecules having a multiple bond between atoms of dissimilar electronegativity:

Example of these class are CO₂, SO₂ and SO₃. In these compounds the oxygen atom are more electronegative than sulphur or carbon atom. 
As a result the electron density of π-electrons is displaced away from carbon to sulphur atoms which are less electronegative than oxygen, towards the O-atom.Thus carbon and sulphur are electron deficient able to accept the electron pair from a Lewis base such as OH⁻ ions to from dative bond.
Lewis Acid and Base Definition and Examples
Examples of Dative Bond

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.
(v) Elements with an Electron sextet:
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

Utility of Lewis Concept:

(i) This concept includes those reactions in which no proton are involved.
(ii) Lewis concept is more general than the Bronsted-Lowery concept(that is Protonic Concept) in that acid base behavior is not dependent on the presence of one particular element or on the presence or absence of a solvent.
(iii) It explains the long accepted basic properties of metallic oxides and acidic properties of non-metallic oxides.
(iv) This theory explain many reactions such as gas-phase high temperature and non-solvent reaction as neutralisation process.