Factors affecting acid strength

Effect of charge :  It becomes more difficult to take a proton away from something that is already negative.  Thus, a neutral parent acid is always stronger as an acid than its anion.  Here is an example:

H₃PO₄ : (K_a = 7.5 x 10^-3)
H₂PO₄^- : (K_a = 6.2 x 10^-8)
HPO₄^2- : (K_a = 2.2 x 10^-13)

Binary Acids (HX) -

polarity of the HX bond - the more electronegative X is, the stronger is HX as an acid CH₄ < NH₃ < H₂O < HF strength of the HX bond - the weaker the HX bond is, the stronger is HX as an acid
stability of the conjugate base X-, the more stable X- is, the stronger is HX as an acid.

hydrogen halide	bond energy (kJ/mol)
HF	567	weak acid
HCl	431	strong acid
HBr	366	strong acid
HI	299	strong acid

OxyAcids (HXO_n)

The acidity of a hypohalous acid increases as the electronegativity of the halogen increases:

HClO : (K_a = 3.0 x 10^-8)
HBrO : (K_a = 2.5 x 10^-9)
HIO : (K_a = 2.3 x 10^-11)

The acidity of an oxyacid increases with the number of oxygen atoms:

H₂SO₄ : (strong acid)
H₂SO₃ : (Ka = 1.7 x 10^-2)

Stability of the conjugate base

This factor (considering the others equal) can play a vital role in deciding the strength of an acid.  The conjugate base results from a loss of a proton, thus, the base is usually an anion, negatively charged.  Therefore, addressing the stability of a conjugate base means considering how well the base can carry an extra electron.

Compare:

CH₃OH (methanol) versus HCOOH (methanoic or formic acid) The conjugate bases are: CH₃O^- (methoxide) versus HCOO^- (methanoate or formate) Clearly, in this case, formate will be more stable.

Compare:

H₃CCH₃ (ethane), H₂CCH₂ (ethene), HCCH (ethyne) In this set, the difference lies in the hybridization of the C.  In ethane, the C-H bond arises from an sp³ hybrid orbital, in ethene, it is sp² , and for ethyne, it is sp.  If any one of the above compounds loses a proton, the electron that stays behind will reside in this hybrid orbital.  The greater the s character of a hybrid orbital, the closer are the electrons, on the average, from the nucleus.  Thus, an extra electron in an sp hybrid orbital will see a higher nuclear charge, rendering the anion to be more stable than it would be if the electron is in an sp² , or sp³ hybrid orbital.  Thus, in terms of acidity:

HCCH > H₂CCH₂ > H₃CCH₃

 

More about Lewis Acids:

The most general definition of acids and bases, which encompasses the Arrhenius and Bronsted-Lowry definitions is due to our old friend, Lewis and his dot structures. A Lewis acid is defined to be any species that accepts lone pair electrons. A Lewis base is any species that donates lone pair electrons. Thus, H⁺ is a Lewis acid, since it can accept a lone pair, while OH^- and NH₃ are Lewis bases, both of which donate a lone pair:

H⁺ + OH^- <=> H₂O

Interestingly, however, is that species which have no hydrogen to donate (a la the Bronsted-Lowry scheme) can still be acids according to the lewis scheme. As an example, consider the molecule BF₃. If we determine Lewis structure of BF₃, we find that B is octet deficient and can accept a lone pair. Thus it can act as a Lewis acid. Thus, when reacting with ammonia, the reaction would look like:

In fact octet deficient molecules are often strong Lewis acids because they can achieve an octet configuration by accepting a lone pair from a Lewis base. Compounds involving elements in periods lower then the second period can act as Lewis acids as well by expanding their valence shells. Thus, SnCl₄ acts as a Lewis acid according to the reaction:

Figure 2:

The central tin atom is surrounded by a valence shell of 12 electrons rather than 8.

 

Titration:

The variation of pH versus the volume of added titrant is a titration curve. Such a curve can be recorded (by an automatic potentiometer) or be established point by point. The curve is sigmoïd and presents a significant variation of pH at the equivalence point, which allows an easy determination of the latter.

 Note : The more diluted is the acid (or bases) the smaller the pH jump.

 Titration of 50 ml of a strong acid HA 0.1000 M by NaOH 0.1000 M :

 

Determining the Equialence Point :

a) Graphical method

During the titration of an acid by a base the pH of the solution is recorded versus the volume of added base and the equivalence point can be determined from the graph pH = V_titrant taking advantage of the approximate symmetry of the curve.

For example, if 25 ml of HCl 0.1 M are titrated by NaOH 0.1 M, the equivalence point occurs at pH = 7.00.