Relative strengths of acids and bases: effect of molecular structure

J. Hollett; D. Latimer; D. Vanderwel

94 Relative strengths of acids and bases: effect of molecular structure

Learning Objectives

By the end of this section, you will be able to:

Rationalize trends in acid–base strength in relation to molecular structure

Examples of common acids and bases are listed in Table 94-1. It is important to be able to recognize acids and bases on the basis of their structures. In this section we will consider the effect of molecular structure on the acidity or basicity of a compound.

Table 94-1: Common acids and bases

Factors to consider when evaluating relative acid strength

As discussed in the previous section, K_a is a quantitative indication of acid strength – i.e. a quantitative indication of the extent that the dissociation reaction (shown below for the acid HA) proceeds to the right:

HA(aq) $\rightleftharpoons$ H⁺(aq) + A^–(aq)

The larger the K_a, the stronger the acid. Like all equilibrium constants, K_a is related to the free energy change (ΔG) for the dissociation reaction,

ΔG = -RT ln K_a

Thus, the stronger the acid, the larger the K_a, and the more negative the ΔG for the dissociation reaction. Since we have learned that ΔG is dependent on both the change in entropy (ΔS) and the change in enthalpy (ΔH) of the reaction, we have the tools to qualitatively relate structure to acid strength.

Note that the dissociation reaction of all Brønsted-Lowry acids can be represented by the same dissociation reaction (loss of a proton), and so one would not expect large differences in ΔS: all of the reactions involve the same change in number of species, and all species are aqueous (there is no change in state). The big difference between most dissociation reactions would, thus, be the enthalpy change of the reaction (ΔH_r) – the difference in stability between the reactants and the products. Since one of the products, H⁺(aq), is the same in all cases, the relative strength of acids can be qualitatively assessed by evaluating the relative stabilities of the acid, HA(aq), and the conjugate base, A^–(aq).

Structural factors that are most commonly considered to evaluate the relative strengths of acids would be as follows:

Electrostatic factors:
- like charges repel; opposite charges attract.
- The ionization of an acid involves the loss of a positive charge and, thus, would be subject to electrostatic effects.
Ability of the acid or base (as applicable) to accommodate charge:
- if the ionization reaction involves a neutral acid species, HA, forming a negatively charged conjugate base, A^–, then the ability of the base to accommodate a negative charge would greatly affect the K_a value.
- if the ionization reaction involves a positively charged acid species, HA⁺, forming a neutral conjugate base, A, then the ability of the acid species to accommodate a positive charge would greatly affect the K_a value.
- Factors such as size, resonance, induction and hybridization can affect stability (discussed below).

Electrostatic factors

Consider the relative acid strengths of the hydronium ion (H₃O⁺), water (H₂O), and hydroxide ion (OH^–), as indicated in Figure 94-1:

**Figure 94-1**: Comparison of the structures (and charges) of the hydronium ion, water, and the hydroxide ion.

The acid dissociation reaction of the hydronium ion involves the loss of a positively charged proton from the cation, to form a neutral conjugate base.

H₃O⁺(aq) $\rightleftharpoons$ H₂O(l) + H⁺(aq)

In comparison, the dissociation reaction of the neutral water molecule involves the separation of a positive and a negative charge (i.e., the proton and the hydroxide ion).

H₂O(l) $\rightleftharpoons$ H⁺(aq) + OH¯(aq)

As a rule of thumb, charge separation is an unfavourable process since opposite charges attract (electrostatically). Thus, intuitively, it would be unfavourable for a positive charge to dissociate from a neutral molecule (leaving behind a negative charge). Moreover, if the proton and hydroxide were dissociated, it would be favourable for them reform neutral water. Thus, we would expect the hydronium ion to be a stronger acid than water.

The acid dissociation reaction of the hydroxide ion involves the dissociation of a proton from the negatively charged hydroxide ion, represented below.

OH¯(aq) $\rightleftharpoons$ H⁺(aq) + O^2¯(aq)

Intuitively, this reaction ought to be extremely unfavourable. It would clearly be unfavourable for a positive charge to dissociate from a negatively charge species, since opposite charges attract. Moreover, the species formed – the oxide ion O^2-(aq) – is a highly charged, small ion and so would not be stable in aqueous solution.

Thus, on the basis of electrostatic considerations, we would be predict that the relatively acidity of the three species to be OH¯ < H₂O < H₃O⁺. This prediction is correct: the K_a values of OH¯, H₂O, and H₃O⁺ are 10^-36, 10^-14, and ~1, respectively. Note that the hydroxide ion is such a poor acid that it will essentially not dissociate in aqueous solution.

The effect of charge can also be seen when we look at the acid dissociation of sulfuric acid, H₂SO₄ (a strong acid). The position of the equilibrium lies far to the right for the dissociation of one proton in aqueous solution:

H₂SO₄(aq) → H⁺(aq) + HSO₄^–(aq)

Note that the conjugate base formed, hydrogen sulfate (HSO₄^–), is itself a Brønsted-Lowry acid. However, unlike H₂SO₄, HSO₄^– is a weak acid: it is not favourable for the positively charged proton to dissociate from the negatively charged ion:

HSO₄^–(aq) ⇆ H⁺(aq) + SO₄^2-(aq) K_a = 1.2 x 10^-2

The K_a value is relatively large compared to many other weak acids, but hydrogen sulfate is a weak acid: consider that the pH of a 1.5 M solution of HSO₄^– is 1.43: this means that only 2.5% of the acid is dissociated.

Ability of acid or conjugate base to accommodate charge:

In the situation where the ionization of the acid forms a negatively charged base, the ability of the conjugate base to accommodate the negative charge can greatly affect the strength of the acid. Qualitatively, anything that can help to spread out the negative charge can help to stabilize the species. Thus, if the atom (ion) bearing the negative charge is larger, or if the negative charge can be spread out over more than one atom (either through resonance or through inductive effects), then the species will be more stable. Other factors such as the hybridization of the atom bearing the charge can also affect stability. Many of these factors will be illustrated in the discussion below.

Factors to consider when evaluating relative base strength

The principles discussed above also apply when evaluating relative base strength: electrostatic factors and the ability of the base and the conjugate acid to accommodate the charge (through resonance, inductive effects, hybridization, size, etc.) can all greatly affect base strength.

In the discussion below we will focus on acid strength. Recall that when the acid dissociates, a conjugate base is formed. Thus, the strength of an acid is related to the strength of its conjugate base. For example, a strong acid will have a very weak conjugate base (i.e., one that is relatively stable so that it is not as favourable to accept a proton).

Relative Strengths of Binary Acids and Bases

A “binary acid” is one that has only two types of atoms: hydrogen and another element.

Effect of Electron Affinity

Across a row in the periodic table, the acid strength of binary hydrogen compounds increases from left to right: the order of increasing acidity across the second row is CH₄ < NH₃ < H₂O < HF (see Figure 94-2). Hydrogen fluoride (HF) is a weak acid; water is neutral. Ammonia (NH₃) is such a poor acid that it would not act as an acid in aqueous solution; the conjugate base amide (NH₂^–) can be formed in aprotic organic solvents and is commonly used in organic chemistry. Similarly, methane (CH₄) does not act as an acid in aqueous solution.

*Figure 94-2*: Acidity trend across the second row of the periodic table. Acid strength increases from left to right across a row.

The trend in acidity across the row is related to the increasing electronaffinity of the nonmetal atom, from left to right across a row. Recall that electronaffinity is an indication of the stability of the anion of an element: fluorine has the highest electronaffinity of any element in the periodic table. Since the electron affinity increases C < N < O < F, the stability of the anions increases CH₃^– < NH₂^– < HO^– < F^–. Hydrofluoric acid, HF, is the strongest acid in this series, since its conjugate base fluoride, F^–, is the most stable. The electronegativity (a measure of the tendency of an atom to attract a bonding pair of electrons) also increases from left to right, which increases the bond polarity and makes it easier (faster) to lose the proton. This is a kinetic, not a thermodynamic, effect.

Effect of size:

The acid strength of binary compounds of hydrogen with nonmetals (A) increases, moving down a group in the periodic table. For group 17, the order of increasing acidity is HF < HCl < HBr < HI. Likewise, for group 16, the order of increasing acid strength is H₂O < H₂S < H₂Se < H₂Te. (See Figure 94-3).

**Figure 94-3**: Acidity trends down two groups of the Period Table. Acid strength increases from top to bottom down the group. Note that the pKa values for HCl, HBr, and HI are not provided since they are strong acids.

In this case, the trend for acidity is opposite to the trend for electron affinity (and electronegativity). Fluorine and oxygen have two of the highest electron affinities of the elements of the periodic table; the electron affinity of the elements DECREASES moving down each group. The trend for acidity is, once again, associated with the increased stabilities of the conjugate bases (the anions), but the reason for the increased stability is different from that when moving across a row. Moving down a group, atomic radius increases: the radius of iodine more than twice that of fluorine (see Figure 36.1a). The larger size (lower charge:mass ratio) allows the negative charge of the anion to be more diffuse and, thus, more stable.

The trends for acidity (and the basicity of the conjugate bases) for the binary acids are summarized below in Figure 94-4.

**FIGURE 94-4**: The figure shows trends in the strengths of binary acids and bases. Relative acidity increases according to the relative stabilities of the conjugate bases. Moving across a row (left to right), the anions are more stable due to increased electron affinity; moving down a group, the anions are more stable due to increased size.

Ternary Acids and Bases

Ternary compounds composed of hydrogen, oxygen, and some third element (“E”) may be structured as depicted in the image below. In these compounds, the central E atom is bonded to one or more O atoms, and at least one of the O atoms is also bonded to an H atom, corresponding to the general molecular formula O_mE(OH)_n. These compounds may be acidic, basic, or amphoteric depending on the properties of the central E atom. Examples of such compounds include sulfuric acid, O₂S(OH)₂, sulfurous acid, OS(OH)₂, nitric acid, O₂NOH, perchloric acid, O₃ClOH, aluminum hydroxide, Al(OH)₃, calcium hydroxide, Ca(OH)₂, and potassium hydroxide, KOH:

A diagram is shown that includes a central atom designated with the letter E. Single bonds extend above, below, left, and right of the E. An O atom is bonded to the right of the E, and an arrow points to the bond labeling it, “Bond a.” An H atom is single bonded to the right of the O atom. An arrow pointing to this bond connects it to the label, “Bond b.”

Hydroxides

If the central atom, E, has a low electronegativity, its attraction for electrons is low. Little tendency exists for the central atom to form a strong covalent bond with the oxygen atom, and bond a between the element and oxygen is more readily broken than bond b between oxygen and hydrogen as shown below:

Hence bond a is ionic, hydroxide ions are released to the solution, and the material behaves as a base. Lower electronegativity is characteristic of the more metallic elements; hence, the metallic elements form ionic hydroxides that are by definition basic compounds. Such examples include the strong bases such as sodium hydroxide (NaOH), potassium hydroxide (KOH), and calcium hydroxide [Ca(OH)₂].

Oxyacids (oxoacids)

If, on the other hand, the atom E has a relatively high electronegativity, it strongly attracts the electrons it shares with the oxygen atom, making bond a relatively strongly covalent. The oxygen-hydrogen bond, bond b, is thereby very polar since electrons are displaced toward E. Hydrogen ions can be released (i.e., the material behaves as an acid), as shown below:

High electronegativities are characteristic of the more nonmetallic elements. Thus, nonmetallic elements form covalent compounds containing acidic −OH groups that are called oxyacids (oxoacids). This is a large group of acids that includes both inorganic acids (phosphoric acid, H₃PO₄, sulfuric acid, H₂SO₄, etc) and organic acids (e.g. the carboxylic acids, R-COOH).

The strength of the oxyacid depends on the identity of the central atom E, as well as the characteristics of the groups attached to E. The more stable the conjugate base (the anion), the stronger the acid will be. The oxyanion could be stabilized by inductive or resonance effects, as discussed below.

Carboxylic acids

The effect of structure on the relative acidity of oxyacids can be nicely illustrated with the carboxylic acids. Carboxylic acids are organic molecules that contain a “carboxyl” group (-COOH). The simplest carboxylic acid is formic acid (methanoic acid), which has only one carbon, but there are literally thousands (or more?) of different carboxylic acids. Although the structure of formic acid is very similar to methanol (which also contains the “OH” group and has only one carbon), formic acid is a much strong acid than methanol, as you can see through comparison of the K_a values:

Methanol can lose a proton to form the alkoxide anion, where the negative charge is accommodated on the oxygen. However, methanol is not called an “acid” since it is not appreciably dissociated in aqueous solution (it is an extremely weak acid). Formic acid is a much stronger acid than methanol: the negative charge is again accommodated on an electronegative oxygen, but in the case of formic acid TWO oxygens are involved in carrying the negative charge, since the carboxylate anion is resonance stabilized.

Since the negative charge is distributed over two oxygens, the anion is more stable. Thus formate (the conjugate base of formic acid) is more stable than methoxide (the conjugate base of methanol); formic acid is a stronger acid than methanol.

The carboxyl group (-COOH) is always acidic, which is why organic compounds that contain the carboxyl group are referred to as “carboxylic acids“. However, the strength of the carboxylic acids will vary according to what groups are attached to the carboxyl group. Consider the effect of the substitution of 1, 2 or 3 fluorine atoms in acetic acid (a 2-carbon carboxylic acid), as shown in Table 94-1.

Table 94-1: Relative acid strength of acetic acid and fluorinated derivatives.
Name	Structure	K_a
acetic acid, CH₃COOH		1.8 x 10^-5
fluoroacetic acid, CH₂FCOOH		2.6 x 10^-3
difluoroacetic acid, CHF₂COOH		4.7 x 10^-2
trifluoroacetic acid, CF₃COOH		5.9 x 10^-1

Note that as the number of fluorine substituents increases, the molecules become more acidic, so that the trifluorinated derivative is more than 10,000 times stronger than acetic acid itself. This effect is not due to resonance stabilization of the conjugate base: rather, the conjugate base is stabilized by the electronegative fluorine atoms drawing electron density away from the negatively charged carboxylate through an inductive effect (i.e., through the bonds). As seen below, this effect is additive: each of the fluorine atoms can “pull” electrons through the Ω bonds to help to disperse the negative charge of the carboxylate anion.

Note that the same sort of trend as observed in Table 94-1 occurs if, instead of adding extra fluorine atoms, we adjust the stabilization of the carboxylate anion by varying the electronegativity of the attached group. As shown in Table 94-2, looking at monohalogenated derivatives of acetic acid, as we add different halogens (moving from the least electronegative, iodine, through to the most electronegative, fluorine), we see that the increased inductive withdrawal of electron density will increase the acidity of the carboxylic acid.

Table 94-2: Relative acid strength of acetic acid and monohalogenated derivatives.
Name	Structure	K_a
acetic acid, CH₃COOH		1.8 x 10^-5
iodoacetic acid, CH₂ICOOH		6.7 x 10^-4
bromoacetic acid, CH₂BrCOOH		1.3 x 10^-3
chloroacetic acid, CH₂ClCOOH		1.4 x 10^-3
fluoroacetic acid, CH₂FCOOH		2.6 x 10^-3

Another trend that we can observe is that, as we increase the distance between the electronegative substituent and the carboxylic acid group, the effect on acidity decreases (see Table 94-3). The effect of the inductive withdrawal of electron density though σ bonds is dependent on distance: the farther the electronegative group is from the carboxylic acid moiety, the less of an effect it will have on acidity.

Table 94-3: Effect of position of substituent on acidity of butanoic acid and monohalogenated derivatives.
Name	Structure	K_a
Butanoic acid, CH₃CH₂CH₂COOH		1.5 x 10^-5
2-bromobutanoic acid, CH₃CH₂CHBrCOOH		1.1 x 10^-3
3-bromobutanoic acid, CH₃CHBrCH₂COOH		1.1 x 10^-4
4-bromobutanoic acid, CH₂BrCH₂CH₂COOH		1.1 x 10^-4

In summary, carboxylic acids are acidic due to resonance stabilization of the conjugate base. The acidity of the carboxylic acids can be affected by the presence of electronegative (or electropositive) substituents. The closer the substituents are to the carboxylic acid moiety, the greater the effect will be.

Other oxyacids:

There are many other types of oxyacids besides the organic carboxylic acids. The factors affecting acidity are the same: anything that can stabilize the conjugate base will increase acidity.

Table 94-4 summarizes the relative acidities of various oxyacids of the form Cl-O-H, with different numbers of oxygens attached to the central chlorine atom. As you can see perchloric acid (HClO₄) is a strong acid, and as the number of oxygens attached to the central chlorine are decreased, the acidity decreases so that hypochlous acid (HClO) is a very weak acid.

Table 94-4: Relative strength of a family of oxyacids
Name	Structure	K_a
perchloric acid, HClO₄		very large
chloric acid, HClO₃		~1
chlorous acid, HClO₂		1.2 x 10^-2
hypochlorous acid, HClO		2.9 x 10^-8

In this case, the oxygens provide tremendous stability to the conjugate bases since they help to diffuse the negative charge through resonance. Shown below are the four resonance contributors to perchloric acid: the negative charge is spread over four oxygen atoms:

The conjugate base of chloric acid would be able to spread the negative charge over three oxygens; chlorous acid over two oxygens; and hypochlorous acid only has one oxygen atom. Thus, the acid strength decreases from perchloric acid to chlorous acid.

A similar trend can be seen with other oxyacids (see Figure 94-2). Sulfuric acid, H₂SO₄, or O₂S(OH)₂, is more acidic than sulfurous acid, H₂SO₃, or OS(OH)₂. After losing one proton, Nitric acid, HNO₃, or O₂NOH, is more acidic than nitrous acid, HNO₂, or ONOH. In each of these pairs, more oxygens are attached to the central atom for the stronger acid, allowing for greater resonance stabilization.

A diagram is shown that includes four structural formulas for acids. A red, right pointing arrow is placed beneath the structures which is labeled “Increasing acid strength.” At the top left, the structure of Nitrous acid is provided. It includes an H atom to which an O atom with two unshared electron pairs is connected with a single bond to the right. A single bond extends to the right and slightly below to a N atom with one unshared electron pair. A double bond extends up and to the right from this N atom to an O atom which has two unshared electron pairs. To the upper right is a structure for Nitric acid. This structure differs from the previous structure in that the N atom is directly to the right of the first O atom and a second O atom with three unshared electron pairs is connected with a single bond below and to the right of the N atom which has no unshared electron pairs. At the lower left, an O atom with two unshared electron pairs is double bonded to its right to an S atom with a single unshared electron pair. An O atom with two unshared electron pairs is bonded above and an H atom is single bonded to this O atom. To the right of the S atom is a single bond to another O atom with two unshared electron pairs to which an H atom is single bonded. This structure is labeled “Sulfurous acid.” A similar structure which is labeled “Sulfuric acid” is placed in the lower right region of the figure. This structure differs in that an H atom is single bonded to the left of the first O atom, leaving it with two unshared electron pairs and a fourth O atom with two unshared electron pairs is double bonded beneath the S atom, leaving it with no unshared electron pairs. — **FIGURE 94-2**: The acids with conjugate bases that are more resonance stabilized (due to more oxygen substituents on the central atom) are stronger acids.

Amphoteric compounds

Hydroxy compounds of elements with intermediate electronegativities and relatively high oxidation numbers (for example, elements near the diagonal line separating the metals from the nonmetals in the periodic table) are usually amphoteric. This means that the hydroxy compounds act as acids when they react with strong bases and as bases when they react with strong acids. The amphoterism of aluminum hydroxide, which commonly exists as the hydrate Al(H₂O)₃(OH)₃, is reflected in its solubility in both strong acids and strong bases. In strong bases, the relatively insoluble hydrated aluminum hydroxide, Al(H₂O)₃(OH)₃, is converted into the soluble ion, ${\left[\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{2}\left({\text{OH}\right)}_{4}\right]}^{\text{-}},$ by reaction with hydroxide ion:

$\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{3}\left({\text{OH}\right)}_{3}\left(aq\right)+{\text{OH}}^{\text{−}}\left(aq\right)\phantom{\rule{0.2em}{0ex}}\rightleftharpoons\phantom{\rule{0.2em}{0ex}}{\text{H}}_{2}\text{O}\left(l\right)+{\left[\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{2}\left({\text{OH}\right)}_{4}\right]}^{\text{-}}\left(aq\right)$

In this reaction, a proton is transferred from one of the aluminum-bound H₂O molecules to a hydroxide ion in solution. The Al(H₂O)₃(OH)₃ compound thus acts as an acid under these conditions. On the other hand, when dissolved in strong acids, it is converted to the soluble ion ${\left[\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{6}\right]}^{3+}$ by reaction with hydronium ion:

${\text{3H}}_{3}{\text{O}}^{\text{+}}\left(aq\right)+\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{3}\left({\text{OH}\right)}_{3}\left(aq\right)\phantom{\rule{0.2em}{0ex}}\rightleftharpoons\phantom{\rule{0.2em}{0ex}}\text{Al}{\left({\text{H}}_{2}\text{O}\right)}_{6}{}^{3+}\left(aq\right)+{\text{3H}}_{2}\text{O}\left(l\right)$

In this case, protons are transferred from hydronium ions in solution to Al(H₂O)₃(OH)₃, and the compound functions as a base.

Key Concepts and Summary

The relative strengths of acids can be predicted through consideration electrostatic factors (like charges repel; opposites attract), as well as consideration of the relative stabilities of the acids and their conjugate bases. If the acid ionization creates an anionic conjugate base, factors that stabilize the anion will provide for a more stable conjugate base (and, thus, a stronger acid). Factors discussed that can stabilize the conjugate base include stabilization through the dispersal of the negative charge due to the size of the anion; resonance effects and inductive effects.

The strengths of the binary acids increase from left to right across a period of the periodic table (e.g. CH₄ < NH₃ < H₂O < HF), and they increase down a group (e.g. HF < HCl < HBr < HI). The strengths of oxyacids increases with increased stabilization of the anionic conjugate base through increased electron affinity of the central element (e.g. H₂SeO₄ < H₂SO₄). The strengths of oxyacids that contain the same central element also increases with increased stabilization of the anionic conjugate base, through resonance stabilization (e.g. H₂SO₃< H₂SO₄) and/or increased inductive withdrawal of negative charge (e.g. CH₃COOH < CF₃COOH).

Section 94 Practice Problems

Click on this link to access the worked solutions to these problems (as a downloadable pdf).

1. Use this list of important industrial compounds (and Figure 93-2) to answer the following questions regarding: CaO, Ca(OH)₂, CH₃CO₂H, CO_2, HCl, H₂CO₃, HF, HNO₂, HNO₃, H₃PO₄, H₂SO₄, NH₃, NaOH, Na₂CO₃.

(a) Identify the strong Brønsted-Lowry acids and strong Brønsted-Lowry bases.

(b) List those compounds in (a) that can behave as Brønsted-Lowry acids with strengths lying between those of H₃O⁺ and H₂O.

(c) List those compounds in (a) that can behave as Brønsted-Lowry bases with strengths lying between those of H₂O and OH⁻.

2. Explain why the ionization constant, K_a, for H₂SO₄ is larger than the ionization constant for H₂SO₃.

3. Explain why HI is a stronger acid than HF.

4. Predict which acid in each of the following pairs is the stronger and explain your reasoning for each.

(a) H₂O or HF

(b) ${\text{HSO}}_{3}{}^{\text{-}}$ or ${\text{HSO}}_{4}{}^{\text{-}}$

(c) NH₃ or H₂S

(d) H₂O or H₂Te

5. Predict which compound in each of the following pairs of compounds is more acidic and explain your reasoning for each.

(a) HSO₄^– or HSeO₄^–

(b) NH₃ or H₂O

(c) PH₃ or HI

(d) NH₃ or PH₃

(e) H₂S or HBr

6. Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign.

(a) acidity: HCl, HBr, HI

(b) basicity: H₂O, OH⁻, H⁻, Cl⁻

(c) basicity: Mg(OH)₂, Si(OH)₄, ClO₃(OH) (Hint: Formula could also be written as HClO₄.)

(d) acidity: HF, H₂O, NH₃, CH₄

7. Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign.

(a) acidity: NaHSO₃, NaHSeO₃, NaHSO₄

(b) basicity: ${\text{BrO}}_{2}{}^{\text{-}},$ ${\text{ClO}}_{2}{}^{\text{-}},$ ${\text{IO}}_{2}{}^{\text{-}}$

(c) acidity: HOCl, HOBr, HOI

(d) acidity: HOCl, HOClO, HOClO₂, HOClO₃

(e) basicity: ${\text{NH}}_{2}{}^{\text{-}},$ HS⁻, HTe⁻, ${\text{PH}}_{2}{}^{\text{-}}$

(f) basicity: BrO⁻, ${\text{BrO}}_{2}{}^{\text{-}},$ ${\text{BrO}}_{3}{}^{\text{-}},$ ${\text{BrO}}_{4}{}^{\text{-}}$

Glossary

oxyacid: ternary compound with acidic properties, molecules of which contain a central nonmetallic atom bonded to one or more O atoms, at least one of which is bonded to an ionizable H atom

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