Concept of Accid and Bases

CONCEPTS OF ACIDS AND BASES Strength of acids and bases (qualitative treatment), balancing equations of chemical reaction, (including oxidation-reduction using ion-electron and oxidation number methods)

Introduction There are several so-called theories of acids and bases, but they are not really theories but merely different definitions of what we choose to call an acid or a base. Since it is only a matter of definition, no theory is more right or wrong than any other, and we use the most convenient theory for a particular chemical situation. So before we talk of strength of acids and bases, we need to know several theories.

Various Theories Regarding Acids & Bases i)

Arrhenius Theory: According to Arrhenius, substances producing H+ ions in solution are acids and those producing OH- ions in solution are bases. Therefore, substances like H2O, HCl, H2SO4, CH3 COOH etc. are acids and the ones like NH 4OH, NaOH, KOH, H2O etc. are bases.

ii)

Bronsted-Lowry Theory: In 1923, Bronsted and Lowry independently defined acids as proton donors, and bases as proton acceptors. For aqueous solutions the definition does not vary much for acids from the Arrhenius theory but it widens the scope of bases. In this, the bases need not contain OH- ions and simply have to accept protons. So ions like Cl-, CH3COO-, Br- etc. which do not contain OH- ions can be considered as bases under this definition. Levelling Solvents: Whenever an acid is dissolved in water, it acts as an acid only if the solvent acts as a base. That is, if we dissolve acids like HCl, HNO 3 etc in water, their acidic strength is almost the same since water acts as a base for both these acids. Infact, it is known that all strong acids show equal acidic strength when dissolved in water. This is because, water acts as a base to all these acids and thus forces them to donate almost the same amount of protons irrespective of their chemical nature. Since water levels the acidic strength of strong acids, it is referred to as a levelling solvent. In order to measure the strength of strong acids, they are dissolved in glacial acetic acid and the amount of protons measured by conductometry. It is found that the strength of acids varies as HClO4 > HBr> H2SO4 > HCl > HNO3 Amphiprotic species: Many molecules and ions can behave like water and may either gain or lose a proton under the appropriate conditions. Such species are said to be amphiprotic. eg. Acid1 HSHBr HCO3−

+ + +

Base2 OHHSCN-

Acid2 H2O H2S HCN

+ + +

Base1 S2BrCO32 −

H3O+

+

HCO3−

H2CO3

+

H2O

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The hydroxides of metals near the boundary between metals and non-metals in the periodic table, are amphiprotic and so react either as acids or as bases. [Al(H2O)3

(OH)3] + OH-

H3O+ + [Al(H2O)3(OH)3] Exercise 1:

iii)

H2O + [Al(H2O)2(OH)4][Al(H2O)4(OH)2]+ + H2O.

Identify conjugate acid base pair in the following reactions C6H5OH + H2O H3O+ + C6H5O– – HCl + OH H2O + Cl–

Lewis Theory: Lewis developed a definition of acids and bases that did not depend on the presence of protons, nor involve reactions with the solvent. He defined acids as materials which accept electron pairs, and bases as substances which donate electron pairs. Thus a proton is Lewis acid and ammonia is Lewis base since, the lone pair of electrons on the nitrogen atom can be donated to a proton: H+ : NH3 → [H ← : NH3]+ 4NH3 + lewis base

NH3

2+

Cu lewis acid

NH3

2+

Cu NH3

NH3

acid - base adduct.

Conditions to be a Lewis Acid i)

Compounds whose central atoms have an incomplete octet. E.g. BF3, AlCl3, GaCl3 etc.

ii)

Compounds in which the central atom has available d-orbital and may acquire more than an octet of valence electrons. E.g. SiF4 + 2F- → SiF62Other examples are : PF3, SF4, SeF4, TeCl4.

iii)

All simple cations : Na+, Ag+, Cu2+, Al3+, Fe3+ etc.

Conditions to be a Lewis Base i)

All simple negative ions, Cl- , F- etc.

ii)

Molecules with unshared pair of electrons, H2O, NH3 etc.

iii)

Multiple bonded compounds which form co-ordination compounds with metals . E.g., CO, NO, Ethylene, Acetylene etc.

Exercise 2:

transition-

In the following reaction, identify each of the reactant as a Lewis acid or Lewis base i) Cr+3 + 6H2O → Cr(H2O)6+3 ii) BF3 + (C2H5)2O →F3B: O(C2H5)2

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Hard and soft acids and bases (Principle of HSAB) Lewis acids & bases are classified as hard & soft acids & bases. Hardness is defined as the property of retaining valence electrons very strongly. Thus hard acid is that in which electron-accepting atom is small with a high positive charge and no electron which are easily polarised or removed. Soft acid is that in which the acceptor atom is large, carries a low positive charge or it has electrons in orbitals which are easily polarised or distorted. A lewis base which holds its electrons strongly is called hard base eg. OH -, F-, NH3, H2O etc. A lewis base in which the position of electrons is easily polarised or removed is called a soft base eg. I-, CO, CH3S-, (CH3)3P etc. A hard acid prefer to bind to hard bases and soft acids prefer to bind to soft bases. The bonding between hard acids & hard bases is ionic & that between soft acids & bases is mainly covalent. Exericise-3: Identify the hard acid and hard bases among the following OH–, H2O, Al+3, Cr+3, Sn+4, PO −43

Relative Strength of Acids and Bases i)

Predict the relative acidic strength among the following: CH3 - H, H2N - H, OH - H, F -H Assume that all these compounds have lost their OH, F-

protons. So we get -CH3 , -NH2,

We shall qualitatively analyse the charge density in each of the species and we shall follow the rule given below: •

Larger the volume over which the charge is spread.

•

Smaller is the charge density.

•

Smaller is the basic character of the ion to attract a proton.

•

Larger is the acidity of the conjugate acid.

In the above case one can see that, the size of the central atom over which the negative charge is present is decreasing from left to right. But it can also be seen that 3/4 th of the volume of C is overlapped by hydrogens in -CH3, 2/3 of the volume of N is overlapped by hydrogens in -NH2, 1/2 the volume of oxygen is overlapped by hydrogens in -OH where as, the entire volume of F is available for the charge in -F SO, actually the space available for the charge is increasing in the order -CH3 < -NH2 < -OH < -F. Therefore, the following conclusion may be arrived at. CH4, NH3, H2O, HF.

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1. Volume available for the negative charge is increasing in the conjugate bases from left to right. 2. Charge density of the conjugate bases is decreasing from left to right. 3. Basicity of the conjugate bases is decreasing from left to right. 4. Acidity of the acids is increasing from left to right. Therefore the increasing acidic character is CH4 < NH3