Chapter 16 Dr. Nabil EL-Halabi

Chapter 16 Dr. Nabil EL-Halabi Acids and Bases

Brønsted Acids and Bases: Brønsted Acid - Any substance that can donate a proton, H+ ion to a base. (hydrogen-ion donors or proton donors). Brønsted Base - Any substance that can accept a proton, H+ ion from an acid. (hydrogen-ion acceptor or proton acceptor). Conjugate Acid-Base Pairs: can be defined as an acid and its conjugate base or a base and its conjugate acid. -The part of the acid remaining when an acid donates a H+ ion is called the conjugate base. The acid formed when a base accepts a H+ ion is called the conjugate acid. -The formula of a conjugate base always has one fewer hydrogen atom and one more negative charge (or one fewer positive charge) than the formula of the corresponding acid. -Every Brønsted acid has a conjugate base, and every Brønsted base has a conjugate acid. - + Ex.: Cl is the conjugate base formed from HCl, and H2O is the conjugate base of the acid H3O .  The ionization of acetic acid can be represented as: - + CH3COOH(aq) + H2O(l) ⇄ CH3COO (aq) + H3O (aq) Acid1 base2 base1 acid2  Ammonia can be classified as Brønsted base: + - NH3(aq) + H2O(l) ⇄ NH4 (aq) + OH (aq) base1 acid2 Acid1 base2  NaOH is not a Brønsted base because it cannot accept proton. However, NaOH is a strong electrolyte that ionized completely in solution. The OH- ion formed is a Brønsted base: + - H3O (aq) + OH (aq) ⇄ 2 H2O(l) Thus, when we call NaOH or any other metal hydroxide a base, we actually reffering to the OH- species derived from the hydroxide. + + H3O represent a hydrated proton in water. H may be associated to more than one H2O molecule + + + + and has formulas such as H5O2 or H9O4 . In calculation, keep in mind that both H and H3O represent the same species in solution. Example 16.1. The Acid-Base Properties of Water: Water is a very weak electrolyte and therefore a poor conductor of electricity, but it does undergo ionization to a small extent: - + H2O(l) ⇄ OH (aq) + H (aq) Autoionization reaction: Two molecules of the solvent react with each other to form ions. - + H2O + H2O ⇄ OH + H3O acid1 base2 base1 acid2 The Ion Product of Water: The equilibrium constant for autoionization of water: + - + - Kc = [H3O ][OH ]/[H2O] or Kc = [H ][OH ]/[H2O]

The concentration of H2O remains unchanged = 55.6 M + - Kc[H2O] = Kw = [H ][OH ] Kw is called the ion-product constant, which is the product of the molar concentrations of H+ and OH- ions at a particular temperature. In pure water at 25oC: [H+] = 1.0 * 10-7 M [OH-] = 1.0 * 10-7 M -7 -7 -14 Kw = (1.0 * 10 )(1.0 * 10 ) = 1.0 * 10 + - -14 Kw = [H ][OH ] = 1.0 * 10

1 When [H+] = [OH-], solution is neutral. When [H+] > [OH-], solution is acidic. When [H+] < [OH-], solution is basic. If we adjust the solution so that [H+] = 1.0 * 10-6 M, the [OH-] must change to: - + -14 -6 -8 [OH ] = Kw/[H ] = 1.0 * 10 / 1.0 * 10 = 1.0 * 10 M. Example 16.2. pH-A measure of Acidity: Because [H+] and [OH-] ions in aqueous solutions are frequently very small numbers therefore we use a more practical measure called pH. The pH of the solution is defined as the negative logarithm of the hydrogen ion concentration (in moles per liter): + + pH = -log [H3O ] or pH = -log [H ]. Because pH is simply a way to express hydrogen ion concentration, acidic and basic solutions at 25oC can be identified by their pH values, as: Acidic solution [H+] > 1.0 * 10-7 M, pH < 7. Basic solution [H+] < 1.0 * 10-7 M, pH > 7. Neutral solution [H+] = 1.0 * 10-7 M, pH = 7. The pH of concentrated acid solutions can be negative. For example, the pH of a 2.0 M HCl solution is –0.30. pOH = -log[OH-] + - -14 [H ][OH ] = Kw = 1.0 * 10 -(log [H+] + log[OH-]) = -log (1.0 * 10-14) -log [H+] - log[OH-] = 14.00 pH + pOH = 14.00 Examples 16.3, 16.4, 16.5. Strength of Acids and Bases: Strong acids: are strong electrolytes, which are assumed to ionize completely in water. - + HCl(aq) + H2O(l)  Cl (aq) + H3O (aq) - + HNO3(aq) + H2O(l)  NO3 (aq) + H3O (aq) - + HClO4(aq) + H2O(l)  ClO4 (aq) + H3O (aq - + H2SO4(aq) + H2O(l)  HSO4 (aq) + H3O (aq

Weak acids: are weak electrolytes, which are assumed to ionize partially in water. - + HF(aq) + H2O(l) ⇄ F (aq) + H3O (aq) - + HNO2(aq) + H2O(l) ⇄ NO2 (aq) + H3O (aq) - 2- + HSO4 (aq) + H2O(l) ⇄ SO4 (aq) + H3O (aq - + H2O + H2O ⇄ OH + H3O

Strong bases: are all strong electrolytes that ionize completely in water. NaOH(aq)  OH-(aq) + Na+(aq) KOH(aq)  OH-(aq) + K+(aq)

Weak bases: are weak electrolytes. + - NH3(aq) + H2O(l) ⇄ NH4 (aq) + OH (aq) Table 16.2 lists some important conjugate acid-base pairs in order of their relative strengths.

2 It is useful to keep the following points in mind:  The conjugate base of a strong acid has no measurable strength. + +  H3O is the strongest acid that can exist in aqueous solution. Acids stronger than H3O react + with water to produce H3O and their conjugate bases. - + HCl(aq) + H2O(l)  Cl (aq) + H3O (aq) + + Acids weaker than H3O react with water to a much smaller extent, producing H3O and their conjugate bases. - + HF(aq) + H2O(l) ⇄ F (aq) + H3O (aq)  The OH-ion is the strongest base that can exist in aqueous solution. Bases stronger than OH- react with water to produce OH- and their conjugate acids.

2- - O (aq) + H2O(l)  2OH (aq) For this reason the oxide ion does not exist in aqueous solutions. Examples 16.6, 16.7.

Weak Acids and Acid Ionization Constants: Consider a weak monoprotic acid HA. Its ionization in water is represented by - + HA(aq) + H2O(l) ⇄ A (aq) + H3O (aq) Or simply - + HF(aq) ⇄ F (aq) + H (aq) The equilibrium constant for this acid ionization, which we call the acid ionization constant, Ka, is + - + - given by: Ka = [H3O ][A ]/[HA] orKa = [H ][A ]/[HA]

At a given temperature, the strength of the acid HA is measured by the magnitude of Ka. Table 16.3 + The larger Ka, the stronger the acid—that is, the greater the concentration of H ions at equilibrium because of its ionization. Solving weak acid ionization problems: 1- Express the equilibrium concentrations of all species in terms of the initial concentrations and a single unknown x, which represents the change in concentrations. 2- Write the acid ionization constant, Ka in terms of equilibrium concentrations. Knowing the value of Ka, we can solve for x. 3- Calculate the equilibrium concentrations of all species and/or pH of the solution. Examples 16.8, 16.9, 16.10.

Percent Ionization Percent Ionization = (ionized acid concentration at equilibrium / initial concentration of acid)*100% The stronger the acid, the greater the percent ionization. For a monoprotic acid HA, we can write the percent ionization as: + Percent Ionization = ([H ]/[HA]o)*100% + [H ] is the concentration at equilibrium and [HA]o is the initial concentration.

3 Diprotic and Polyprotic Acids Diprotic and polyprotic acids may yield more than one hydrogen ion per molecule. These acids ionize in a stepwise manner, that is, they lose one proton at a time. An ionization constant expression can be written for each ionization stage. For example, for H2CO3 we write: - + + - H2CO3(aq) ⇄ HCO3 (aq) + H (aq) Ka1 = [H ][HCO3 ]/[H2CO3] - 2- + + 2- - HCO3 (aq) ⇄ CO3 (aq) + H (aq) Ka2 = [H ][ CO3 ]/[HCO3 ] Note that the conjugate base in the 1st ionization stage becomes the acid in the second ionization stage. + + Ka1 > Ka2 because it is easier to remove H ion from a neutral molecule than to remove another H from a negatively charged ion derived from the molecule. Table 16.4 shows the ionization constants of several diprotic and a polyprotic acids. Example 16.11 Weak Bases and Base Ionization Constants: When ammonia dissolves in water, it undergoes the reaction: + - NH3(aq) + H2O(l) ⇄ NH4 (aq) + OH (aq) [OH-] > [H+], and therefore pH > 7. + - Kb = [NH4 ][OH ]/[NH3] Kb is the base ionization constant. Table 16.5 lists a number of common weak bases and their ionization constants. Example 16.12 The Relation Between Conjugate Acid-Base Ionization Constants: Use acetic acid as an example: - + CH3COOH(aq) ⇄ CH3COO (aq) + H (aq) + - Ka = [H ][CH3COO ]/[CH3COOH] - The conjugate base, CH3COO , reacts with water according to the equation: - - CH3COO (aq) + H2O(l) ⇄ CH3COOH(aq) + OH (aq) - - Kb = [CH3COOH][OH ]/[CH3COO ] + - - - + - KaKb = ([H ][CH3COO ]/[CH3COOH])*([CH3COOH][OH ]/[CH3COO ]) = [H ][OH ] = Kw This result can be understood by realizing that sum of reactions (1) and (2) here is simply the autoionization of water. - + 1 CH3COOH(aq) ⇄ CH3COO (aq) + H (aq) Ka - - 2 CH3COO (aq) + H2O(l) ⇄ CH3COOH(aq) + OH (aq) Kb - + 3 H2O(l) ⇄ OH (aq) + H (aq) Kw

Thus for any conjugate acid-base pair it is always true that: Kw = KaKb Ka = Kw / Kb Kb = Kw / Ka The stronger the acid (the larger Ka), the weaker its conjugate base (the smaller Kb), and vice versa (see Tables 16.3, 16.4, 16.5). All anions can act as weak bases except those that are the conjugate bases of the strong acids. Acid-Base Properties of Salts: Salt hydrolysis: the reaction of an anion or a cation of a salt, or both, with water. Salt hydrolysis usually affects the pH of a solution. Salts that produce Neutral Solutions: Salts containing an alkali metal ion or alkaline earth metal ion (except Be2+) and the conjugate base - - - of a strong acid (for example Cl , Br , and NO3 ) do not undergo hydrolysis to appreciable extent, and their solutions are neutral.

4 NaNO3 dissolves in water and dissociates completely as: - + NaNO3(aq)  NO3 (aq) + Na (aq) + - A solution containing Na and NO3 ions is neutral, with a pH of 7. Salts That Produce Basic Solutions: The dissociation of CH3COONa (salt derived from a weak acid and a strong base) in water is given by: - + CH3COONa(aq)  CH3COO (aq) + Na (aq) + - Na has no acidic or basic properties. CH3COO however, is conjugate base of weak acid + CH3COOH and therefore has an affinity for H ions. The hydrolysis reaction is given by: - - CH3COO (aq) + H2O(l) ⇄ CH3COOH(aq) + OH (aq) - CH3COONa solution is basic, because this reaction produces OH . - - Kb = [CH3COOH][OH ]/[CH3COO ] Salts That Produce Acidic Solutions: When a salt derived from a strong acid and a weak base dissolves in water , the solution become acidic. + - NH4Cl(aq)  NH4 (aq) + Cl (aq) - + + The Cl ion has no affinity for H ions. The NH4 is the weak conjugate acid of the weak base. + + NH4 (aq) + H2O(l) ⇄ NH3(aq) + H3O (aq) Or simply + + NH4 (aq) ⇄ NH3(aq) + H (aq) + NH4Cl solution is acidic, because this reaction produces H . + + -14 -5 -10 Ka = [H ][NH3]/[NH4 ] = Kw / Ka = 1.0*10 /1.8*10 = 5.6*10 Example 16.13. Metal ion hydrolysis: Salts containing small and highly charged metal cations (Al3+, Fe3+, Cr3+, Bi3+, and Be2+) and conjugate bases of strong acids produce an acidic solution. 3+ 3+ AlCl3 dissolves in water, the Al ion take hydrated form Al(H2O)6 .

3+ 2+ + Al(H2O)6 (aq) ⇄ Al(OH)(H2O)5 (aq) + H (aq) 2+ + 3+ Ka = [Al(OH)(H2O)5 ][H ]/[Al(H2O)6 ] 2+ Al(OH)(H2O)5 can undergo further ionization: 2+ + + Al(OH)(H2O)5 (aq) ⇄ Al(OH)2(H2O)4 (aq) + H (aq)

5 Salts in which Both the cation and the Anion hydrolyze: Table 16.6 is very important; it summarizes the behavior in aqueous solution of salts discussed in this section.

 pH is determined by the stronger Ka or Kb of the ions.

 If Kb = Ka, the solution is neutral.

 If Kb > Ka, the solution is basic; the anion will hydrolyze to a greater extent than the cation. At equilibrium, OH- ions > H+ ions.

 If Kb < Ka, the solution is acidic; the cation will hydrolyze to a greater extent than the anion. At equilibrium, OH- ions < H+ ions. - Some anions can act either as an acid or as a base as HCO3 - 2- + -11 HCO3 (aq) + H2O(l) ⇄ CO3 (aq) + H3O (aq) Ka = 4.8 *10 - - -8 HCO3 (aq) + H2O(l) ⇄ H2CO3 (aq) + OH (aq) Kb = 2.4 *10 Because Kb > Ka, we predict that the hydrolysis reaction will outweigh the ionization process. Thus, a solution of NaHCO3 will be basic. Example 16.14.

Acidic, Basic, and Amohoteric Oxides: In the representative elements, all alkali metal oxides and alkaline earth metal oxides except BeO are basic. BeO and several metallic oxides in Groups 3A and 4A are amphoteric. Nonmetallic oxides in which the oxidation number of the representative element is high are acidic (for example, N2O5, SO3, and Cl2O7). Those in which the oxidation number is low (for example, CO and NO) show no measurable acidic properties.  The basic metallic oxides react with water to form metal hydroxides:

Na2O(s) + H2O(l)  2NaOH(aq)  The reaction between acidic oxides and water are: CO2(g) + H2O(l) ⇄ H2CO3(aq) SO3(g) + H2O(l) ⇄ H2SO4(aq) The reaction between SO3 and H2O is responsible for acid rain.  Reaction between acidic oxides and bases and those between basic oxides and acids, produce a salt and water, like acid-base reaction.

CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)

BaO(s) + 2HNO3  Ba(NO3)2(aq) + H2O(l) Amphoteric Oxides: can behave either as acidic oxide or as a basic oxide depending on the reaction condition

Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l)

Al2O3(s) + 2NaOH(aq) + 3H2O(l)  2NaAl(OH)4(aq)  Some transition metal oxides in which the metal has a high oxidation number act as acidic oxides (for example, Mn2O7 and CrO3).  The higher the oxidation state of the metal, the more covalent the compound; the lower the oxidation state, the more ionic the compound. Lewis Acids and Bases: Lewis Acids: a substance that can accept a pair of electrons. Lewis base: a substance that can donate a pair of electrons. F H F H F B + :N H F B N H F H F H acid base

6 Selected Problems: 5 – 8, 15 – 24, 30 – 34, 41 – 45, 47, 51, 52, 55, 56, 67, 68, 70, 73 – 77.