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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.
As the oxidation number of the central atom E increases, the acidity also increases.

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(H 2 O) 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(H 2 O) 3 (OH) 3 , is converted into the soluble ion, [ Al ( H 2 O ) 2 ( OH ) 4 ] , by reaction with hydroxide ion:

Al ( H 2 O ) 3 ( OH ) 3 ( a q ) + OH ( a q ) H 2 O ( l ) + [ Al ( H 2 O ) 2 ( OH ) 4 ] ( a q )

In this reaction, a proton is transferred from one of the aluminum-bound H 2 O molecules to a hydroxide ion in solution. The Al(H 2 O) 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 [ Al ( H 2 O ) 6 ] 3+ by reaction with hydronium ion:

3H 3 O + ( a q ) + Al ( H 2 O ) 3 ( OH ) 3 ( a q ) Al ( H 2 O ) 6 3+ ( a q ) + 3H 2 O ( l )

In this case, protons are transferred from hydronium ions in solution to Al(H 2 O) 3 (OH) 3 , and the compound functions as a base.

Key concepts and summary

The strengths of Brønsted-Lowry acids and bases in aqueous solutions can be determined by their acid or base ionization constants. Stronger acids form weaker conjugate bases, and weaker acids form stronger conjugate bases. Thus strong acids are completely ionized in aqueous solution because their conjugate bases are weaker bases than water. Weak acids are only partially ionized because their conjugate bases are strong enough to compete successfully with water for possession of protons. Strong bases react with water to quantitatively form hydroxide ions. Weak bases give only small amounts of hydroxide ion. The strengths of the binary acids increase from left to right across a period of the periodic table (CH 4 <NH 3 <H 2 O<HF), and they increase down a group (HF<HCl<HBr<HI). The strengths of oxyacids that contain the same central element increase as the oxidation number of the element increases (H 2 SO 3 <H 2 SO 4 ). The strengths of oxyacids also increase as the electronegativity of the central element increases [H 2 SeO 4 <H 2 SO 4 ].

Key equations

  • K a = [ H 3 O + ] [ A ] [ HA ]
  • K b = [ HB + ] [ OH ] [ B ]
  • K a × K b = 1.0 × 10 −14 = K w
  • Percent ionization = [ H 3 O + ] eq [ HA] 0 × 100

Chemistry end of chapter exercises

Explain why the neutralization reaction of a strong acid and a weak base gives a weakly acidic solution.

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Explain why the neutralization reaction of a weak acid and a strong base gives a weakly basic solution.

The salt ionizes in solution, but the anion slightly reacts with water to form the weak acid. This reaction also forms OH , which causes the solution to be basic.

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Use this list of important industrial compounds (and [link] ) to answer the following questions regarding: CaO, Ca(OH) 2 , CH 3 CO 2 H, CO 2, HCl, H 2 CO 3 , HF, HNO 2 , HNO 3 , H 3 PO 4 , H 2 SO 4 , NH 3 , NaOH, Na 2 CO 3 .

(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 H 3 O + and H 2 O.

(c) List those compounds in (a) that can behave as Brønsted-Lowry bases with strengths lying between those of H 2 O and OH .

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Source:  OpenStax, Chemistry. OpenStax CNX. May 20, 2015 Download for free at http://legacy.cnx.org/content/col11760/1.9
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