When a substance is dissolved in water, it may behave as an acid, behave as a base, or exhibit no acid–base properties. How does the chemical structure of a substance determine which of these behaviors is exhibited by the substance? For example, why do some substances that contain OH groups behave as bases, releasing OH ions into solution, whereas others behave as acids, ionizing to release H+ ions? In this section we discuss briefly the effects of chemical structure on acid–base behavior.

Factors That Affect Acid Strength

A molecule containing H will act as a proton donor (an acid) only if the H—X bond is polarized in this way:

In ionic hydrides, such as NaH, the reverse is true: the H atom possesses a negative charge and behaves as a proton acceptor (a base). Nonpolar H—X bonds, such as the H—C bond in CH4, produce neither acidic nor basic aqueous solutions.

A second factor that helps determine whether a molecule containing an H—X bond will donate a proton is the strength of the bond. Very strong bonds are less easily broken than weaker ones. This factor is important, for example, in the hydrogen halides. The H—F bond is the most polar H—X bond. You therefore might expect HF to be a very strong acid if bond polarity were all that mattered. However, the H—X bond strength increases as you move up the group: 299 kJ/mol in HI, 366 kJ/mol in HBr, 431 kJ/mol in HCl, and 567 kJ/mol in HF. Because HF has the highest bond strength among the hydrogen halides, it is a weak acid, whereas all the other hydrogen halides are strong acids in water.

A third factor that affects the ease with which a hydrogen atom ionizes from HX is the stability of the conjugate base, X. In general, the greater the stability of the conjugate base, the stronger the acid.

The strength of an acid is often a combination of all three factors.

Binary Acids

For a series of binary acids HX in which X represents members of the same group in the periodic table, the strength of the H—X bond is generally the most important factor determining acid strength. The strength of an H—X bond tends to decrease as the element X increases in size. As a result, the bond strength decreases and acidity increases down a group. Thus, HCl is a stronger acid than HF, and H2S is a stronger acid than H2O.

Bond polarity is the major factor determining acidity for binary acids HX when X represents members of the same period. Thus, acidity increases as the electronegativity of the element X increases, as it generally does moving from left to right across a period. For example, the difference in acidity of the period 2 elements is . Because the C—H bond is essentially nonpolar, CH4 shows no tendency to form H+ and CH3 ions. Although the N—H bond is polar, NH3 has a nonbonding pair of electrons on the nitrogen atom that dominates its chemistry, so NH3 acts as a base rather than an acid.

FIGURE 16.17 Trends in acid strength for the binary hydrides of periods 2–4.

The periodic trends in the acid strengths of binary compounds of hydrogen and the nonmetals of periods 2 and 3 are summarized in FIGURE 16.17.


What is the major factor determining the increase in acidity of binary acids going down a group? What is the major factor going across a period?


Many common acids, such as sulfuric acid, contain one or more O—H bonds:

Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom are called oxyacids. At first it may seem confusing that the OH group, which we know behaves as a base, is also present in some acids. Let's take a closer look at what factors determine whether a given OH group behaves as a base or as an acid.

Consider an OH group bound to some atom Y, which might in turn have other groups attached to it:

At one extreme, Y might be a metal, such as Na or Mg. Because of the low electronegativity of metals, the pair of electrons shared between Y and O is completely transferred to oxygen, and an ionic compound containing OH is formed. Such compounds are therefore sources of OH ions and behave as bases, as in NaOH and Mg(OH)2.

When Y is a nonmetal, the bond to O is covalent and the substance does not readily lose OH. Instead, these compounds are either acidic or neutral. Generallyas the electronegativity of Y increasesso does the acidity of the substance. This happens for two reasons: First, as electron density is drawn toward Y, the O—H bond becomes weaker and more polar, thereby favoring loss of H+. Second, because the conjugate base of any acid YOH is usually an anion, its stability generally increases as the electronegativity of Y increases. This trend is illustrated by the Ka values of the hypohalous acids (YOH acids where Y is a halide ion), which decrease as the electronegativity of the halogen atom decreases (FIGURE 16.18).

FIGURE 16.18 Acidity of the hypohalous oxyacids (YOH) as a function of electronegativity of Y

Many oxyacids contain additional oxygen atoms bonded to the central atom Y. These atoms pull electron density from the O—H bond, further increasing its polarity. Increasing the number of oxygen atoms also helps stabilize the conjugate base by increasing its ability to “spread out” its negative charge. Thus, the strength of an acid increases as additional electronegative atoms bond to the central atom Y. For example, the strength of the chlorine oxyacids (Y = Cl) steadily increases as O atoms are added:

Because the oxidation number of Y increases as the number of attached O atoms increases, this correlation can be stated in an equivalent way: In a series of oxyacids, the acidity increases as the oxidation number of the central atom increases.


Which acid has the larger acid-dissociation constant, HIO2 or HBrO3?

SAMPLE EXERCISE 16.20 Predicting Relative Acidities from Composition and Structure

Arrange the compounds in each series in order of increasing acid strength: (a) AsH3, HBr, KH, H2Se; (b) H2SO4, H2SeO3, H2SeO4.


Analyze We are asked to arrange two sets of compounds in order from weakest acid to strongest acid. In (a), the substances are binary compounds containing H, and in (b) the substances are oxyacids.

Plan For the binary compounds, we will consider the electronegativities of As, Br, K, and Se relative to the electronegativity of H. The higher the electronegativity of these atoms, the higher the partial positive charge on H and so the more acidic the compound.

For the oxyacids, we will consider both the electronegativities of the central atom and the number of oxygen atoms bonded to the central atom.


(a) Because K is on the left side of the periodic table, it has a very low electronegativity (0.8, from Figure 8.7). As a result, the hydrogen in KH carries a negative charge. Thus, KH should be the least acidic (most basic) compound in the series.

Arsenic and hydrogen have similar electronegativities, 2.0 and 2.1, respectively. This means that the As—H bond is nonpolar, and so AsH3 has little tendency to donate a proton in aqueous solution.

The electronegativity of Se is 2.4, and that of Br is 2.8. Consequently, the H—Br bond is more polar than the H—Se bond, giving HBr the greater tendency to donate a proton. (This expectation is confirmed by Figure 16.17, where we see that H2Se is a weak acid and HBr a strong acid.) Thus, the order of increasing acidity is KH < AsH3 < H2Se < HBr.

(b) The acids H2SO4 and H2SeO4 have the same number of O atoms and the same number of OH groups. In such cases, the acid strength increases with increasing electronegativity of the central atom. Because S is slightly more electronegative than Se (2.5 vs 2.4), we predict that H2SO4 is more acidic than H2SeO4.

For acids with the same central atom, the acidity increases as the number of oxygen atoms bonded to the central atom increases. Thus, H2SeO4 should be a stronger acid than H2SeO3. We predict the order of increasing acidity to be H2SeO3 < H2SeO4 < H2SO4.


In each pair, choose the compound that give the more acidic (or less basic) solution: (a) HBr, HF; (b) PH3, H2S; (c) HNO2, HNO3(d) H2SO3, H2SeO3.

Answers: (a) HBr, (b) H2S, (c) HNO3(d) H2SO3

Carboxylic Acids

Another large group of acids is illustrated by acetic acid, a weak acid (Ka = 1.8 × 10–5):

The portion of the structure shown in red is called the carboxyl group, which is often written COOH. Thus, the chemical formula of acetic acid is written as CH3COOH, where only the hydrogen atom in the carboxyl group can be ionized. Acids that contain a carboxyl group are calledcarboxylic acids, and they form the largest category of organic acids. Formic acid and benzoic acid are further examples of this large and important category of acids:

Two factors contribute to the acidic behavior of carboxylic acids. First, the additional oxygen atom attached to the carbon of the carboxyl group draws electron density from the O—H bond, increasing its polarity and helping to stabilize the conjugate base. Second, the conjugate base of a carboxylic acid (a carboxylate anion) can exhibit resonance  (Section 8.6), which contributes to the stability of the anion by spreading the negative charge over several atoms:


The general structure of amino acids, the building blocks of proteins, is

where different amino acids have different R groups attached to the central carbon atom. For example, in glycine, the simplest amino acid, R is a hydrogen atom, and in alanine R is a CH3 group:

Amino acids contain a carboxyl group and can therefore serve as acids. They also contain an NH2 group, characteristic of amines (Section 16.7), and thus they can also act as bases. Amino acids, therefore, are amphiprotic. For glycine, we might expect the acid and base reactions with water to be

The pH of a solution of glycine in water is about 6.0, indicating that it is a slightly stronger acid than base. The acid–base chemistry of amino acids is more complicated than shown in Equations 16.47 and 16.48, however. Because the COOH group can act as an acid and the NH2 group can act as a base, amino acids undergo a “self-contained” Brønsted–Lowry acid–base reaction in which the proton of the carboxyl group is transferred to the basic nitrogen atom:

Although the form of the amino acid on the right in this equation is electrically neutral overall, it has a positively charged end and a negatively charged end. A molecule of this type is called a zwitterion (German for “hybrid ion”).

Do amino acids exhibit any properties indicating that they behave as zwitterions? If so, their behavior should be similar to that of ionic substances.  (Section 8.2) Crystalline amino acids have relatively high melting points, usually above 200°C, which is characteristic of ionic solids. Amino acids are far more soluble in water than in nonpolar solvents. In addition, the dipole moments of amino acids are large, consistent with a large separation of charge in the molecule. Thus, the ability of amino acids to act simultaneously as acids and bases has important effects on their properties.



What group of atoms is present in all carboxylic acids?