Chemistry II Notes Chapter 16: Acids & Bases Acid-Base Equilibria 16.1. Acids and Bases: A Brief Review Acid and base behavior in an aqueous medium is briefly reviewed. Arrhenius defined an acid as a substance that produces H+ ions in water; he defined a base as a substance that produces OH– ions in water. HCl—one of the strong acids—is an Arrhenius acid. Sodium hydroxide—one of the strong bases— is an Arrhenius base. 16.2 Brønsted-Lowry Acids and Bases The Arrhenius definitions of acids and bases are limited to aqueous reactions. A broader definition is the Brønsted-Lowry approach. According to Brønsted-Lowry, an acid is a substance that donates an H+ ion to another substance; a base is a substance that accepts an H+ ion. Chemists use the term proton to refer to the aqueous hydrogen ion, H+. They also refer to the hydronium ion, H3O+ in the context of acids. All four of these terms, H+, hydrogen ion, H3O+, and hydronium ion, are used interchangeably. In water a proton interacts with the lone pairs on oxygen atoms of water molecules and becomes hydrated. So it is slightly more realistic to represent an aqueous proton as H3O+, which depicts the proton attached to a water molecule. In order for one substance to behave as an acid, another substance must behave as a base (and vice versa). A Brønsted-Lowry acid must have a proton to donate, and a Brønsted-Lowry base must have a lone pair of electrons in order to accept the proton. Some substances are capable of acting as an acid in one reaction and as a base in another. Such substances are called amphoteric. When an acid molecule ionizes in water, it donates its proton to a water molecule, producing a hydronium ion and an anion. The anion produced by the ionization is the conjugate base of that acid. (HX is used to denote a generic acid.) HX (aq) + H2O (l) H3O+ (aq) + X- (aq) HX (the acid) and X– (the conjugate base) differ only by a proton. They are a conjugate acid-base pair. Every acid has a conjugate base. By the same token, every base has a conjugate acid. 16.3 The Autoionization of Water Pure water has a very small tendency to ionize, acting both as an acid (donating a proton) and as a base (accepting a proton). At 25°C, the Kc for this process is 1.0 10–14, which means that only about one molecule per billion undergoes this autoionization. The equilibrium expression for the autoionization of water is (Recall that a liquid does not appear in the equilibrium expression.) Because the autoionization of water is a very important equilibrium, its equilibrium constant is given a special subscript, w. For any aqueous solution at 25°C, the product of hydronium and hydroxide ion concentrations is equal to Kw. In neutral water, where the only source of either ion is the autoionization, the hydronium and hydroxide ion concentrations are equal. So the concentrations of both hydronium ion and hydroxide ion in neutral water is 1.0 10–7 M. 16.4 The pH Scale 16.5 Strong Acids and Bases Strong Acids Hydrochloric acid Hydrobromic acid Hydroiodic acid Strong Bases HCl LiOH Lithium hydroxide HBr NaOH Sodium hydroxide HI KOH Potassium hydroxide HNO3 Nitric acid RbOH Rubidium hydroxide HClO3 Chloric acid CsOH Cesium hydroxide HClO4 Perchloric acid Ba(OH)2 Barium hydroxide H2SO4 Sulfuric acid Sr(OH)2 Strontium hydroxide Strong acids are those that ionize completely in water. Strong acids are also strong electrolytes. Strong bases are ionic compounds that dissociate completely in water. They include the hydroxides of group 1A metals and group 2A metals. 16.6 Weak Acids A weak acid is one that ionizes partially in water to produce hydronium ion and a conjugate base. Acetic acid is a weak acid. The ionization equilibrium of a weak acid has an associated equilibrium constant called the acid-dissociation constant. The equilibrium constant for a weak acid ionization is subscripted with an a for acid. The equilibrium expression for the above equation is As with all equilibria, the larger the value of K, the further the equilibrium lies to the right. This means that the larger the value of Ka, the stronger the acid. Table 16.2 lists some weak acids, their conjugate bases, and their acid-dissociation constants. 16.7 Weak Bases A weak base is one that ionizes partially in water to produce hydroxide ion and a conjugate acid. Ammonia is an example of a group of weakly basic compounds called amines. An amine contains a nitrogen atom with a lone pair of electrons. The site of the lone pair is where a proton can be accepted. In general, Being an equilibrium, the ionization of a weak base has an equilibrium constant called the base-dissociation constant associated with it. The equilibrium constant for a weak base ionization has the subscript b for base. The equilibrium expression for the above equation is As with all equilibria, the larger the value of K, the further the equilibrium lies to the right. This means that the larger the value of Kb, the stronger the base. Table 16.4 lists some weak bases, their conjugate acids, and their base-dissociation constants. 16.8 Relationship Between Ka and Kb As we saw in Section 16.2, a weak acid has a strong conjugate base. It is important to distinguish between the meaning of the word strong in the context of acids and bases and its meaning in the context of conjugate acids and bases. A strong acid is one that ionizes completely in water. A strong conjugate acid is one that is sufficiently acidic compared with water to protonate the water to some extent. Note again the difference in equations used to represent each. The ionization of HCl is written with an arrow in one direction. The ionization of the ammonium ion (the conjugate acid of the weak base ammonia) shows a double arrow, denoting an equilibrium. 16.9 Acid-Base Properties of Salt Solutions The pH of a salt solution can be predicted qualitatively by considering the ionic constituents of the salt. 1. Salts derived from a strong base and a strong acid: Examples are NaCl and Ca(NO3)2, which are derived from NaOH and HCl and from Ca(OH)2 and HNO3, respectively. Neither cation nor anion hydrolyzes. The solution has a pH of 7. 2. Salts derived from a strong base and a weak acid: In this case the anion is a relatively strong conjugate base. Examples are NaClO and Ba(C2H3O2)2. The anion hydrolyzes to produce OH–(aq) ions. The cation does not hydrolyze. The solution has a pH above 7. 3. Salts derived from a weak base and strong acid: In this case the cation is a relatively strong conjugate acid. Examples are NH4Cl and Al(NO3)3. The cation hydrolyzes to produce H+(aq) ions. The anion does not hydrolyze. The solution has a pH below 7. 4. Salts derived from a weak base and a weak acid: Examples are NH4C2H3O2, NH4CN, and FeCO3. Both cation and anion hydrolyze. The pH of the solution depends on the extent to which each ion hydrolyzes. The pH of a solution of NH4CN is greater than 7 because CN– (Kb = 2.0 10–5) is more basic than NH4+ (Ka = 5.6 10–10) is acidic. Consequently, CN– hydrolyzes to a greater extent than NH4+ does. 16.10 Acid-Base Behavior and Chemical Structure As we move from left to right across a row in the periodic table, there is less change in bond strength. In this case, what determines acid strength is the polarity of the H- X bond. The electronegativity of elements increases from left to right across a period in the periodic table. As the electronegativity of X increases, the polarity of the H- X bond increases, increasing acidity. The strength of an oxyacid depends on the electronegativity of the central nonmetal to which the OH groups are bound and on the number of oxygen atoms bound to the central nonmetal atom. For a series of oxyacids with the same number of oxygen atoms, the acidity increases with the electronegativity of the nonmetal. For a series of acids with the same central nonmetal atom, the acidity increases with the number of oxygen atoms bound to the central atom. (This also relates increasing acidity to increasing oxidation number on the central atom.) Carboxylic acids are all weak acids. Their acidity stems partly from the fact that the second oxygen withdraws electron density from the O–H bond, making it more polar than it would otherwise be. The acidity is increased by the presence of additional electron-withdrawing (electronegative) atoms, such as halogens. In addition, the anion left behind after donation of a proton (the conjugate base) is fairly stable. The negative charge on the anion can be accommodated on either of the two oxygen atoms, stabilizing the anion via resonance. 16.11 Lewis Acids and Bases The Brønsted-Lowry acid-base theory broadens the definitions of acids and bases to include reactions that do not occur in aqueous media. Lewis acid-base theory further broadens the definition to include reactions other than proton-transfer reactions. A Lewis acid is defined as an electron-pair acceptor, and a Lewis base as an electron-pair donor. In the examples we have seen of Arrhenius and BrønstedLowry acid-base behavior, the bases are all acting as Lewis bases. In the ionization of an acid in water, the water molecule donates electrons (a lone pair on the oxygen atom) to the hydrogen atom of the acid. As a bond forms between the oxygen atom on water and the acidic proton, the bond between the acidic proton and its original molecule is broken. The Lewis definition of an acid does not require that an acid have a proton to donate, only that it be able to accept a pair of electrons from a Lewis base.