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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.