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Acid Dissociation Constant

Chemical compounds that are bonded together by electrovalent bond generally dissociate in to ions in a suitable polar solvent.

Some of them dissociate completely and some others only to some extent. The point where no more dissociation takes place under given condition is called equilibrium point. At equilibrium point the solution contains both the anions, cations of the compound along with the undissociated compound. Acids and bases also have this property of dissociating and in various proportions. This dissociation takes place in the aqueous medium. Thus the acid strength and base strength is relevant only in their aqueous solutions.

There are certain acids which dissociate completely in aqueous medium. They are called strong acids. There are some other acids which show a limited dissociation in aqueous medium. These are called weak acids. The same is the case with bases. There are weak bases and strong bases based on their capacity to dissociate.

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Acid Dissociation Constant ka

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  1. According to the Arrhenius theory of acids and bases, acids are those which contain hydrogen ion and bases or alkalis are those which contain hydroxyl ion.
  2. The chemical activity and the electrical conductivity experiments are the basis of this theory.
  3. Thus the proton (H+) and the hydroxyl ion (OH-) are responsible for acidic and basic properties of compounds respectively.
  4. The reaction product of an acid and base is called a salt. This theory was a distinct advancement in the study of compounds but it failed to explain the similar properties exhibited by some other compounds which are not having either a proton to liberate or a hydroxyl ion.
  5. Water is considered as a very good solvent for ionic compounds because of the reason that it can stabilize the ions that are formed due to the dissociation by attracting them strongly.
  6. This attraction is strong particularly because of the asymmetry in the shape of water molecule which results in the asymmetrical distribution of the charges.
  7. Ions get hydrated by many water molecules that surround them and hence the stability of the ions.
  8. A more powerful theory was proposed to explain the general concept of acids and bases and it is known as Lowry-Bronsted definition.
  9. According to this theory, an acid is a species having a tendency to donate or lose a proton and a base is a species having the tendency to accept or add a proton.
  10. Water can act as an acid or a base. When it hydrates a proton it acts as a base and when it the hydrated water gets dehydrated it acts as an acid.
In acid-base reaction H+ ion is transferred from acid to a base. Acid is generally represented as HA while corresponding base represented as A-.

HA(aq) acid + H2O(l) base $\to$ H3O+(aq) corresponding acid + A-(aq) corresponding base

The strength of the acid can be defined from the equilibrium position of the above equation.
Ka = $\frac{[H_3O^+] [A^-]}{[HA]}$


Ka = $\frac{[H^+] [A^-]}{[HA]}$
The equilibrium constant Ka, in the case of acid is called dissociation constant of acid or simply strength constant. There is no difference between H3O+ and H+ since the proton in aqueous medium always exists as hydronium ion.Similarly for bases too there is a dissociation constant which is given by Kb. Based on the above equation the by expressing in logarithmic expression:

pKa = - log [Ka]

Acid Dissociation Constant of Water

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As indicated above water acts as both acid and base. In presence f a stronger acid, it acts as a base by accepting a proton while in presence of a stronger base it donates the proton. This property is often referred to as ampholyte. This may also give information that water reacts in itself and is in equilibrium.

H2O (liquid) + H2O (liquid) ↔ H3O+ (aq) + OH- (aq)

The equilibrium position in the above equation is far to the left in pure conditions.

Hence the equilibrium constant of water is

Kw = 1.0 X 10-14 M

Since the dissociation equilibrium is more towards H2O, and by the dissociation of water the ratio of H3O+ and OH- ions is 1:1,

[H3O]+ = [OH]- = 1.0 X 10-7 M

Determination of the Dissociation Constant of a Weak Acid

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Weak acids are those which will not get dissociated completely in aqueous solution. An non dissociated amount of acid will still remain at the point of equilibrium It is difficult to find the Dissociation constant for such equilibrium where all the entities are unknowns.
If it is a monobasic acid it is easy to equate the formed H3O+ ions and the A- ions and take it as [H3O+]2 so that the two unknowns can be converted to one. But the non dissociated acid still remains an unknown which will leave all unknowns.In order to solve this predicament, an approximation is required. This is done by taking excess of pure weak acid so that the concentration of the remaining undissociated can be taken as negligible in the whole amount. Since we know the amount and purity of the added acid we can get the approximate value for the un dissociated weak acid. Thus we can find the dissociation constant of the weak acid.

Dissociation Constant of Acetic Acid

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Acetic acid is a weak acid. Acetic acid will not dissociate completely in water. Its reaction with water is,

CH3COOH + H2O ↔ H3O+ + CH3COO-

The equilibrium constant for the equation is

Ka =
$\frac{[H_3O]^+ [CH_3COO]^-}{[CH_3COOH]}$

Since the acetic acid is a weak acid the dissociation is not complete and it will be partial. Thus the denominator faction will be considerable. The amount of non dissociated acetic acid in the medium can be found out in several ways of experiments.

The dissociation constant of acetic acid is thus

Ka (acetic acid) = 1.8 X 10-5 M

Acid Dissociation Constant Table

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 The Acid Dissociation Constant Table has Given Below:-

Acid Formula Ka1
Acetic CH3COOH 1.75 x 10-5
Ammonium Ion NH4+ 5.70 x 10-10
Arsenic H3AsO4 5.8 x 10-3 1.1x10-7 3.2 x 10-12
Benzoic C6H5COOH 6.28 x 10-5
Boric H3BO3 5.81 x 10-10
Carbonic H2CO3 4.45 x 10-7 4.69 x 10-11
Citric HOOC(OH)C(CH2COOH)2 7.45 x 10-4 1.73 x 10-5 4.02 x 10-7
Formic HCOOH 1.80 x 10-4
Glycolic HOCH2COOH 1.47 x 10-4
Hydrazinium Ion H2NNH3+ 1.05 x 10-8
Hydrogen Peroxide H2O2 2.2 x 10-12
Hydrogen Sulfide H2S 9.6 x 10-8 1.3 x 10-14
Hydroxyl Ammonium IonHONH3+1.10 x 10-6
Tartaric acidHOOC(CHOH)2COOH9.20 x 10-4 4.31 x 10-5
SuccinicHOOCCH2CH2COOH6.21 x 10-52.31 x 10-6
Iodic acidHIO31.7 x 10-1
Phosphoric acidH3PO47.11 x 10-36.32 x 10-84.5 x 10-13
SulfamicH2NSO3H1.03 x 10-1
Lactic acidCH3CHOHCOOH1.38 x 10-4
Maleic acidcis-HOOCCH:CHCOOH1.3 x 10-25.9 x 10-7
Nitrous acidHNO27.1 x 10-4
Hypochlorous HOCl 3.0 x 10-8
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