Acid-Base Equilibria: The Nature of Acids and Bases What makes an Acid an Acid? An acid possess a sour taste An acid dissolves active metals like magnesium An acid causes certain vegetable dyes to turn characteristic colors What makes a Base a Base? A bases possess a bitter taste A base feels slippery to the touch A base causes certain vegetable dues to turn a characteristic color

Acid-Base Equilibria: The Nature of Acids and Bases The Arrhenius Definition of an Acid and a Base: An acid is a substance that produces H+ ions in water solutions A base is a substance that produces OH- ions in a water solution If a solution contains more OH- ions than H+ ions we say that the solution is basic If a solution contains more H+ ions than OH- ions we say that the solution is acidic

Acid-Base Equilibria: The Proton in Water When HCl dissolves in water we write: H2O HCl(g)  H+(aq) + Cl-(aq) H : H+  O : H hydronium ion (H3O+)

Acid-Base Equilibria: The Proton in Water When HCl dissolves in water we write: H2O HCl(g)  H+(aq) + Cl-(aq)

Acid-Base Equilibria: The Proton in Water When HCl dissolves in water we write: H2O HCl(g)  H+(aq) + Cl-(aq) Acidic solutions are formed by a chemical reaction in which and acid transfers a proton (H+) to water. HCl(aq) + H2O(aq)  H3O+(aq) + Cl-(aq)

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases An acid may be defined as a substance that is capable of donating protons and a base may be defined as substance than accepts protons. + H H : _ : : : Cl H + : : N H  Cl + H N H : : H H Bronsted acid Bronsted base

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases Describing the relationship between the Arrhenius and Bronsted definitions of acids and bases, NH3(aq) + H2O(aq)  NH4+(aq) + OH-(aq) The Bronsted definition says that H2O is a an acid because it donated a proton and NH3 is a base because it accepted the proton The Arrhenius definition says that this solution is basic Water is an acid

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2(aq) The Bronsted definition says that acetic acid is a an acid because it donates a proton to water and and water is a base because it accepts the proton from the acetic acid Water is also a base

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2(aq) Bronsted acid Bronsted base Conjugate acid Conjugate base NH3(aq) + H2O(aq)  NH4+(aq) + OH-(aq) Bronsted base Bronsted acid Conjugate acid Conjugate base The stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger its conjugate base.

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases Of course all this Bronsted-Lowry ‘stuff’ raises a number of questions. If water can be an acid and a base, can it act as a proton donor and acceptor to itself? What makes one acid or base strong and another acid or base weak?

Acid-Base Equilibria: The Proton in Water The Relative Strengths of Acids and Bases What makes one acid or base strong and another acid or base weak?

Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale If water can be an acid and a base, can it act as a proton donor and acceptor to itself? Water is capable of auto-ionizing. H2O (l)  H+ (aq) + OH- (aq) H O : +  .. The reaction occurs to a very small extent; about 1 in 108 molecules is ionized at any given moment Protons transfer from one molecule to another at a rate of about 1000 times per second

Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale If this is true: H2O (l)  H+ (aq) + OH- (aq) [H+] [OH-] Than this is true: K = [H2O] And since water is a liquid and its concentration is therefore constant, this expression may be written as: Kw = [H+] [OH-] where Kw is the ion product constant and is equal to 1.0 x 10-14 Note that [H+] = [OH-] = 1.0 x 10-7 M, water is therefore said to be neutral However, in most solutions these concentrations vary. If If [H+] > [OH-] , solution is acidic If [H+] > [OH-] , solution is basic

Acid-Base Equilibria: The Proton in Water The Bronsted-Lowry Concept of Acids and Bases Sample exercise: Indicate whether each of the following solutions is neutral, acidic, or basic: (a) [H+] = 2 x 10-5 M, (b) [OH-] = 0.010 M, ( c) [OH-] = 1.0 x 10-7 M Sample exercise: Calculate the concentration of H+ (aq) in (a) a solution in which the [OH-] is 0.020M, (b) a solution in which the [OH-] = 2.5 x 10-6 M. Indicate whether the solution is acidic or basic

Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale

Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale Because the concentration of H+ ions is often quite small, it can be conveniently expressed in terms of pH = -log [H+] For example, solution with a [H+] = 2. 5 x 105 has a pH of: pH = -log [2. 5 x 10-5 ] = 4.6 Likewise a solution with a pH of 3.8 has a H+ concentration of: Antilog -3.8= 1.58 x 104 M

Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale In a sample of lemon juice, [H+] = 3.8 x 10-4 M. What is the pH. A commonly available window cleaner has a [H+] = 5.3 x 10-9 M In a sample of freshly pressed apple juice has a pH of 3.76. Calculate the [H] Now, you try it!

Sooooo… if Kw = [H+] [OH-] and Kw = 1.0 x 10-14, then: Acid-Base Equilibria: The Proton in Water The Dissociation of Water and the pH Scale Because the concentration of OH- ions is often quite small, it can be conveniently expressed in terms of p)H = -log [OH-] Now lets think about this, if the [H] = [OH-], then the pH = pOH = 7 Sooooo… if Kw = [H+] [OH-] and Kw = 1.0 x 10-14, then: -log Kw = 14 = pH + pOH Sample exercise: What is the pH of a solution with a pOH of 2.5? Is the solution acidic or basic?

Acid-Base Equilibria: The Proton in Water Measuring the pH Using Indicators

Acid-Base Equilibria: The Differences Between Strong and Weak Acids HX H+ X- HX HX H+ X- Initial Equilibrium Initial Equilibrium

Acid-Base Equilibria: The Differences Between Strong and Weak Acids

Acid-Base Equilibria: The Differences Between Strong and Weak Acids Dealing with a Strong Acid: What is the pH of 0.010 M solution of HCl? Dealing with a weak acid that is only partially ionizable: [H+][X-] Since HX (aq)  H+(aq) + X-(aq), then Ka = [HX] The smaller the value of the acid dissociation constant Ka, the weaker the acid What is the Ka of a 0.10 M solution of formic acid (HCHO2) which has a pH = 2.38? HCHO2  H+ + CHO2 I C E

Acid-Base Equilibria: The Differences Between Strong and Weak Acids What is the concentration of H+ ions in a 0.10 M solution of HC2H3O2 (Ka = 1.8 x 10-5) HC2H3O2  H+ + C2H3O2 I C E What is the pH of the solution? What is the percent ionization of this solution?

Acid-Base Equilibria: The Differences Between Strong and Weak Acids What is the pH and percent ionization of a 0.20 M solution of HCN? Ka = 4.9 x 10 -10 I C E

Acid-Base Equilibria: Dealing with Polyprotic Acids Substances that are capable of furnishing more than one proton to water are called polyprotic acids. H2SO3(aq)  H+(aq) + HSO3-(aq) K a1 = 1.7 x 10-2 HSO3-(aq)  H+(aq) + SO32-(aq) K a2 = 6.4 x 10-8 Because Ka1 is so much larger than subsequent dissociation constants for most polyprotic acids, almost all the H+ (aq) in the solution come from the first ionization reaction.

Acid-Base Equilibria: Dealing with Polyprotic Acids

Acid-Base Equilibria: Dealing with Polyprotic Acids What is the pH of 0.0037 M solution of carbonic acid (H2CO3) H2CO3  H+ + HCO3- I C E HCO3 -  H+ + CO32- I C E

Acid-Base Equilibria: Strong Bases The most common soluble strong Bases are the hydroxides of group IA and heavier group 2A metals What is the pH of a 0.010 M solution of Ba(OH)2?

Acid-Base Equilibria: Dealing with Weak Bases The base dissociation constant Kb refers to the equilibrium in which a base reacts with H2O to from the conjugate acid and OH- Weak base + H2O  conjugate acid + OH- NH3 (aq) + H2O (l)  NH4(aq) + OH-(aq) [NH4+] [OH-] Kb = [NH3] Calculate the [OH-] in a 0.15 M solution of NH3. NH3 + H2O  NH4+ + OH- I C E

Acid-Base Equilibria: Classes of Weak Acids Amines Anions of Weak Acids

Acid-Base Equilibria: Anions of Weak Acids HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2- (aq) Bronsted acid Bronsted base Conjugate acid Conjugate base A second class of weak base is composed of the anions of weak acids Anions of weak acids can be incorporated into salts NaC2H3O2  Na+(aq) + C2H3O2- (aq) C2H3O2- + H2O  HC2H3O2 + OH- Kb = 5.6 x 1010

Acid-Base Equilibria: Anions of Weak Acids Calculate the pH of a 0.01 M solution of sodium hypochlorite (NaClO) + H2O  + OH- I C E

Acid-Base Equilibria: Anions of Weak Acids Now it’s you turn: the Kb for BrO- is 5.0 x 10-6. Calculate the pH of a 0.050 M solution of NaBrO

Acid-Base Equilibria: Relationship Between Ka and Kb NH4+(aq)  NH3(aq) + H+ (aq) NH3(aq) + H2O NH4+(aq) + OH- (aq) [H+][NH3] [NH4][OH- ] Ka = Kb = [NH4+] [NH3] NH4+(aq)  NH3(aq) + H+ (aq) NH3(aq) + H2O(l) NH4+(aq) + OH- (aq) H2O  H+(aq) + OH-(aq) When two reactions are added to give a third reaction, the equilibrium constant for the third reaction reaction is given by the product of the equilibrium constants for the two added reactions Ka x Kb = Kw pKa + pKb = pKw

Acid-Base Equilibria: Relationship Between Ka and Kb Calculate the (a) base-dissociation constant, Kb, for the fluoride ion, is the pKa of HF = 3.17 pKa = -log Ka 3.17 = -log Ka Antilog -3.17 = 6.76 x 10-4 Since Ka x Kb = Kw (6.76 x 10-4)x Kb = 1.0 x 10-14 Kb = 1.0 x 10-14/ 6.76 x 10-4 = 1.5 x 10-11

Acid-Base Equilibria: Relationship Between Ka and Kb Calculate the pKb for carbonic acid (Ka = 4.3 x 10-7) Now it’s your turn

Acid-Base Equilibria: Acid-Base Properties of Salt Solutions Anions of weak acids, HX, are basic and will react with H2O to produce OH- X- (aq) + H2O (l)  HX(aq) + OH-(aq) Anions of strong acids, such as NO3-, exhibit no basicitiy, these ions do not react with water and consequently do not influence the pH Anions of polyprotic acids, such as HCO3-, that still have ionizable protons are capable of acting as either proton donors or acceptors depending upon the magnitudes of the Ka or Kb This last one requires a bit of an explanation

Acid-Base Equilibria: Acid-Base Properties of Salt Solutions Anions of polyprotic acids, such as HCO3-, that still have ionizable protons are capable of acting as either proton donors or acceptors depending upon the magnitudes of the Ka or Kb Predict whether the salt Na2HPO4 will form an acidic or basic solution on dissolving in water. Na2HPO4  2Na+ (aq) + HPO4- HPO4- acting like an acid HPO4- (aq) + H2O  H3O+ + PO43-(aq) K3 = 4.2 x 10-13 HPO4- acting like an base HPO4- (aq) + H2O  H2PO42-(aq) + OH-(aq) So HPO- is the conjugate base of H2PO4-. Since the K2 of H2PO4- = 6.2 x 10-8 then: Kw 1.0 x 10-14 Kb = = 1.6 x 10-7 = Ka 6.2 x 10-8 Since Kb is larger than Ka, HPO4- will act like a base

Acid-Base Equilibria: Acid-Base Properties of Salt Solutions Salt derived from a strong base and a strong acid will have a pH of 7 Salt derived from a strong base and a weak acid will have a pH above 7 Salt derived from a weak acid and a weak base depends upon whether the dissolved ion acts as an acid or a base as determined by the size of the Ka or Kb

Acid-Base Equilibria: Acid-Base Character and Chemical Structure A substance HX will transfer a proton only if the H X bond, is already polarized in the following way: H X In ionic compounds such as NaH, the H atom possess a negative charge and behaves like a proton acceptor. Very strong bonds in an HX are less easily ionizable than weak bonds

Acid-Base Equilibria: Acid-Base Character and Chemical Structure