In general, for acids H x A in the same group, i.e, HCl, HBr, etc, the bond strength is the more important of these two factors.

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1 15.15 Factors that Affect Acid Strength There are several factors that can affect acid strength, and the importance of the factors can be variable. However, some trends are notable. 1) Bond strength The strength of the bond between the acidic proton and the rest of the molecule will have an effect on acidity. The weaker the bond, the more acidic the acid will be generally. 2) Bond polarity The polarity of a bond is the distribution of the electrons between the two bonded atoms. If the electrons are fairly equally distributed, the bond is not very polar. As the electron distribution gets weighted towards one atom, the bond becomes more polar. A highly polar bond between an acidic hydrogen and another atom tends to make it more easy for the proton to leave the molecule than would happen for a non-polar bond. In general, for acids H x A in the same group, i.e, HCl, HBr, etc, the bond strength is the more important of these two factors. Polarity versus bond strength Here we see that the bond between hydrogen and halogen atoms is most polar in HF and least polar in HI, but HF is a weak acid, while the rest are strong acids. This is because the bond strength of HF is much greater than the other bond strengths. 1

2 For acids of elements in the same row, the bond strengths tend to be more similar to each other, and so the polarity of the bond plays a greater role in determining acid strength. Since polarity depends on the electronegativity of the atoms involved, acid strength tends to increase with electronegativity, because the bond tends to be more polar. Here we see that methane is nonpolar, while polarity, electronegativity, and acid strength, increase as we move along the row to HF. Polarity versus strength( continued) In general, 1) since electronegativity tends to increase from left to right on the periodic table, and 2)bond strength with hydrogen tends to decrease as we move down the periodic table, stronger acids (H x A) tend to involve elements in the lower right of the table. Oxoacids and acid strength Another category of acids are oxoacids, where the acidic proton is bonded to oxygen, which is in turn bonded to another atom (HOX). For this category of acids, acid strength tends to increase with the electronegativity of the other atom bonded to the oxygen, and with an increase in its oxidation number. 2

3 Oxoacids and acid strength Oxoacids with the same atom X will be strongest when many other atoms are bonded to X. Here the oxidation number of chlorine increases as more atoms are bonded to the chlorine. The larger positive oxidation number means the polarity of the O-Cl bond increases as Cl tries to reduce the large positive charge it is experiencing. Oxygen then tries to replace some of this electron density by pulling electron density away from the hydrogen. Problem Identify the stronger acid in each of the following pairs: a) H 2 S or H 2 Se H x A acids from the same group tend affected more by bond strength, which decreases as we go down the table. H 2 Se is stronger. b) HI or H 2 Te H x A acids from the same row tend affected more by electronegativity, which increase as we go right on the table. HI is stronger. c) HNO 2 or HNO 3 While the formulas make it difficult to see it directly, these are oxoacids (HONO x ) where an oxygen is bonded to the hydrogen and the nitrogen. Since oxoacids of the same element get stronger as the oxidation number increase, HNO 3 should be the stronger acid. d) H 2 SO 3 or H 2 SeO 3 While the formulas make it difficult to see it directly, these are oxoacids (HOXO 2 ) where an oxygen is bonded to the hydrogen and the other atom. Since oxoacids of the different elements get stronger as electronegativity increases, H 2 SO 3 should be the stronger acid, as S is more electronegative. 3

4 15.16 Lewis Acids and Bases Around the same time that Brønsted-Lowry acids and bases were defined, G.N. Lewis proposed a different definition that is more general in scope. A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor. These definitions are more general than the Brønsted-Lowry definitions because protons do not need to be involved for a compound to be a Lewis acid. More specifically, a proton is the ultimate electron pair acceptor, but it is not the only species capable of being an electron pair acceptor. There exist Lewis acids that are not Brønsted-Lowry acids. However, we saw that Brønsted-Lowry bases must all have at least one lone pair of electrons, so any Brønsted-Lowry base is also a Lewis base, or vice versa. acid-base adducts In the picture, BF 3 is a Lewis acid, as the B atom can accept 2 electrons to complete its octet of electrons. The N of the NH 3 has a lone pair that can be donated, making NH 3 a Lewis base. The product of the Lewis acid-base reaction is an acid-base adduct. 4

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Since we will be dealing with aqueous acid and base solution, first we must examine the behavior of water.

Acids and Bases Know the definition of Arrhenius, Bronsted-Lowry, and Lewis acid and base. Autoionization of Water Since we will be dealing with aqueous acid and base solution, first we must examine the