• Nahkleh Group
• Robinson Group
• Weaver Group
• Bodner Group

Acids and Bases

Factors that Control the Relative Strengths of Acids and Bases

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The Polarity of the XH Bond

When all other factors are kept constant, acids become stronger as the XH bond becomes more polar. The second-row nonmetal hydrides, for example, become more acidic as the difference between the electronegativity of the X and H atoms increases. HF is the strongest of these four acids, and CH₄ is one of the weakest Brnsted acids known.

When these compounds act as an acid, an H-X bond is broken to form H⁺ and X^- ions. The more polar this bond, the easier it is to form these ions. Thus, the more polar the bond, the stronger the acid.

An 0.1 M HF solution is moderately acidic. Water is much less acidic, and the acidity of ammonia is so small that the chemistry of aqueous solutions of this compound is dominated by its ability to act as a base.

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The Size of the X Atom

At first glance, we might expect that HF, HCl, HBr, and HI would become weaker acids as we go down this column of the periodic table because the X-H bond becomes less polar. Experimentally, we find the opposite trend. These acids actually become stronger as we go down this column.

This occurs because the size of the X atom influences the acidity of the X-H bond. Acids become stronger as the X-H bond becomes weaker, and bonds generally become weaker as the atoms get larger as shown in the figure below.

The K_a data for HF, HCl, HBr, and HI reflect the fact that the X-H bond-dissociation enthalpy (BDE) becomes smaller as the X atom becomes larger.

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The Charge on the Acid or Base

The charge on a molecule or ion can influence its ability to act as an acid or a base. This is clearly shown when the pH of 0.1 M solutions of H₃PO₄ and the H₂PO₄^-, HPO₄^2-, and PO₄^3- ions are compared.

Compounds become less acidic and more basic as the negative charge increases.

Acidity: H₃PO₄ > H₂PO₄^- > HPO₄^2-

Basicity: H₂PO₄^- < HPO₄^2- < PO₄^3-

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The Oxidation State of the Central Atom

There is no difference in the polarity, size, or charge when we compare oxyacids of the same element, such as H₂SO₄ and H₂SO₃ or HNO₃ and HNO₂, yet there is a significant difference in the strengths of these acids. Consider the following K_a data, for example.

The acidity of these oxyacids increases significantly as the oxidation state of the central atom becomes larger. H₂SO₄ is a much stronger acid than H₂SO₃, and HNO₃ is a much stronger acid than HNO₂. This trend is easiest to see in the four oxyacids of chlorine.

This factor of 10¹¹ difference in the value of K_a for hypochlorous acid (HOCl) and perchloric acid (HOClO₃) can be traced to the fact that there is only one value for the electronegativity of an element, but the tendency of an atom to draw electrons toward itself increases as the oxidation number of the atom increases.

As the oxidation number of the chlorine atom increases, the atom becomes more electronegative. This tends to draw electrons away from the oxygen atoms that surround the chlorine, thereby making the oxygen atoms more electronegative as well, as shown in the figure below. As a result, the O-H bond becomes more polar, and the compound becomes more acidic.

The relative strengths of Brnsted bases can be predicted from the relative strengths of their conjugate acids combined with the general rule that the stronger of a pair of acids always has the weaker conjugate base.