Chapter 16 : Acid-Base Equilibria

Chapter 16 : Acid-Base Equilibria Created by Lauren Querido

16.1 Review 16.2 Brønsted-Lowry Acids and Bases 16.3 Autoionization of Water 16.4 pH Scale 16.5 Strong Acids and Bases 16.6 Weak Acids 16.7 Weak Bases 16.8 Relationship Between Kaand Kb 16.9 Acid-Base Properties of Salt Solutions 16.10 Acid-Base Behavior and Chemical Structure 16.11 Lewis Acids and Bases Table of Contents

16.1 Review • Acids • Sour in taste • Litmus paper turns red • Bases • Bitter, slippery • Litmus paper turns blue • When acids and bases mix, their properties disappear!

Arrhenius Acids and Bases • Svante Arrhenius (1880) • In aqueous solutions: • Acids will increase the concentration of H+ ions when dissolved in water. • Bases will increase the concentration of OH-ions when dissolved in water.

16.2 Brønsted-Lowry Acids and Bases • 1923 Brønsted and Lowry made a more general definition • Brønsted-Lowry Acid is a substance that can transfer a proton. It must have a hydrogen atom that can be lost as H+. • Brønsted-Lowry Base is asubstance that can accept a proton. Must have a nonbonding pair of electrons to gain a H+ ion.

Conjugate Acid-Base Pairs • Conjugate base- Removal of proton from the acid • Conjugate acid- Addition of proton to the base

Relative Strengths of Acids and Bases • The stronger the acid, the weaker its conjugate base. • The stronger the base, the weaker its conjugate acid. 1. Strong acids completely transfer protons to water. 2. Weak acids partly dissociate in aqueous solutions and exist as a mixture of acid molecules and component ions. 3. Negligible acidity contain Hydrogen but do not demonstrate acidic behavior. Ex: CH4 • Position of equilibrium favors transfer of proton from stronger acid to stronger base.

16.3 Autoionization of Water • Ion product of water • 1.0 x 10 –14= [H+] [OH-] • This is used to calculate concentrations of H+ and OH- . • If [H+] = [OH-], than neutral equation • If [H+] > [OH-], than acidic equation • If [H+] < [OH-], than basic equation =

16.4 The pH Scale • pH = -log [H+] • pH of 7 is neutral • Acidic solution 0 < pH < 7 • Basic solution 14 > pH > 7 • Other p scales are • pOH = -log [OH-] • pOH + pH = -log Kw = 14.0

Examples on the pH Scale

Measuring pH • A pH meter consists of a pair of electrodes connected to a meter which pH is generated when placed in the solution. • An acid-base indicator turns a color if placed in acid or base. Ex: litmus paper

16.5 Strong Acids and Bases • Strong Acids • 7 most common strong acids are • HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4 • In acidic reactions, equilibrium lies entirely to the right side. • Completely dissociates • Example: • HNO3 => H+ + NO3-

Strong Bases • Most common strong bases are ionic hydroxides of alkali metals (1A) and heavier alkaline earth metals (2A). Examples: LiOH, RbOH, CsOH, NaOH, KOH, and Ca(OH)2, Sr(OH)2, and Ba(OH)2. • Other strong bases react with water to form OH- such as Na2O, CaO. • Also, anions O2-, H-, and N3- are stronger bases than OH- and therefore remove a proton from H2O. • Example: N3- + H2O => NH3 + 3OH-

16.6 Weak Acids • A weak acid only partially ionizes in aqueous solutions. • General weak acid equation • HX  H+ + X- where H is Hydrogen • Many weak acids contain some Hydrogen atoms bonded to carbon atoms and oxygen atoms (organic compounds). • Ka is the acid dissociation constant. • The larger the value of Ka , the stronger the acid.

Calculating Ka from pH • Use and ICE box! • Sample exercise • A student prepared a .10 M solution of formic acid and measures its pH which was 2.38. • A) calculate Ka for formic acid • B) what percentage of the acid is ionized in the .10M solution?

a) HCHO2 H+ + CHO2- Ka = [H+][CHO2-] [HCHO2] pH= -log[H+] =10 –2.38 = 4.2 X 10-3M Ka = [4.2 X 10–3][4.2 X 10–3] [.10] 1.8 X 10-4 = [4.2 X 10–3][4.2 X 10–3] [.10] b) Percent Ionization = Concentration of H+ Initial concentration of component = 4.2% Answer

Using Ka to Calculate pH The best way to explain this is by an example. Calculate the pH of a .30 M solution of acetic acid at 25o C. (Ka = 1.8 X 10-5) So… HC2H3O2 H+ + C2H3O2- Ka = [H+][C2H3O2-] = 1.8 X 10-5 [HC2H3O2] What now?

Ka = (x)(x) =1.8 X 10-5 (.30 – x) Either do the quadratic equation or in this case you can take out x in the denominator. [H+] = x = 2.3 X 10-3 pH = -log 2.3 X 10-3 = 2.64

Polyprotic Acids • Polyprotic acids have more than one ionizable Hydrogen atom. • Example: • H2SO3 H+ + HSO4- • HSO4-  H+ + SO32- • The second Ka (Ka2) is much smaller than Ka1 because it is easier to remove the first proton.

16.7 Weak Bases • Weak base + water => conjugate acid + hydroxide ion • Kb is the base-dissociation constant (equilibrium in which base reacts when H2O to form conjugate acid and OH- ion). • Types of weak bases: • Neutral substances that have atoms with a non-bonding pair of electrons that can serve as a proton acceptor. • Most of these contain amines, N-H which is sometimes replaced with a bond between C or N Ex: NH2CH3 • Anions of weak acids • Ex: ClO- + H2O  HClO + H+ • ClO- is the weak base

16.8 Relationship Between Ka and Kb • Reaction 1 + reaction 2 = reaction 3 Which leads to K1 x K2 = K3 Which leads to Ka x Kb = Kw • Kw is the ion-product constant for water • Kw = 1 x 10-14 • As the strength of the acid increases, the strength of the base decreases and visa-versa. • pKa + pKb = pKw = 14.00

16.9 Acid-Base Properties of Salt Solutions • Hydrolysis is the process at which ions react with water and produce H+ or OH- X- + H2O HX + OH- • Anions of strong acids do not influence pH • Ex: NO3- • Anions that still have ionizable protons are amphoteric • Ex: HSO3- from H2SO4 • Most cations (except 1A elements and Ca+2, Sr+2. Ba+2) act as weak acids in solution.

Predicting the pH of a Solution 1. Salts derived from a strong acid and a strong base makes a neutral pH (pH of 7). • NaOH + HCl => NaCl + H2O 2. Salts derived from a strong base and a weak acid will yield a pH of above 7 because the anion hydrolyzes to produce OH- ions and the cation does not hydrolyze. • NaOH + HClO => NaClO + H2 3. Salts derived from a weak base and a strong acid will result in a pH that is below 7 because the cation hydrolyzes to produce H+ ions and the anion does not hydrolyze. • Al(OH)3 + 3HNO3 => Al(NO3)3 + 3H2O

4. Salts derived from a weak base and a weak acid will yield a pH that is dependant on the constant value of the constant dissociations (Ka and Kb). • if the base is more basic than the acid is acidic, then the solution will have a pH that is greater than 7. • If the acid is more acidic, than the pH will be less than 7. • NH4+ + CN- NH4CN • NH4+ Ka = 5.6 X 10-10 • CN- Kb= 2.0 X 10-5 • Therefore, the pH of NH4CN is greater than 7

16.10 Acid-Base Behavior and Chemical Structure • Factors that effect acid strength • If H-X bond is polarized (X is more electronegative) the H acts as a proton acceptor. • Non-polar bonds (CH4) produce neutral solutions. • Weaker bonds dissociate more easily than very strong bonds. • HF is a weak acid because of this. • The greater the stability of the conjugate base, the weaker the acid. • Ultimately, there are three factors effecting acid strength: • Polarity of H-X bond • Strength of H-X bond • Stability of conjugate base, X-

Binary Acids • Binary acids are composed of Hydrogen and a non-metal. • Ex: HCl, HF, H2S, etc. • The more polar the bond,the stronger it is • The weaker the bond, the stronger the acid. • Strength of the bond decreases (acidity increases) as the element increases in size or moves down a group. • Acid strength increases (acidity decreases) moving from left to right

Increasing base strength Increasing acid strength Increasing acid strength Increasing base strength

Oxyacids • Oxyacids are acids with an OH group is bound to a central atom. • Example: H2SO4

OH- Bonding • To determine if an OH group acts as an acid or base, consider this: • If Y is a metal than sources of OH- behave as bases. • If Y is a non-metal than the compound will not readily lose the OH- ion. • The electronegativity will increase and so will the acidity. • The increasing number of Oxygen atoms stabilizes the conjugate base and thus increases the strength of the acid.

Oxyacid Rules of Thumb • Oxyacids that have the same number of OH groups and the same number of Oxygen atoms, acid strength increases with increasing elecronegativity of the central atom • Example: HClO > HBrO > HIO (> = more acidic) 2. For oxyacids with the same central atom, acid strength increases with increasing number of Oxygen atoms that are attached. • Example: HClO < HClO2 < HClO3 < HClO4 ( < less acidic)

Carboxylic Acid • Carboxylic acids are organic compounds. • -COOH is the functional group • -R is either a Hydrogen or Carbon based group • If an extra Oxygen is added than it stabilizes the conjugate base and increases the acidity. • If conjugate base has resonance structures, it spreads the negative charge evenly over the compound. • Acid strength of carboxylic acid increases as the number of electronegative atoms increase.

6.11 Lewis Acids and Bases • G.N. Lewis proposed this: • Lewis Acids have an incomplete octet of electrons. Function as electron pair acceptors • Lewis Bases act as electron pair donators

Hydrolysis of Metal Ions • Hydration is a process when when metals attract unshared electron pairs of water molecules. • The metal acts as Lewis acid • The water acts as Lewis base • Ex: Fe(H2O)6+3 Fe(H2O)5(OH)2+ + H+ • So, general equation • M(H2O)nc  M(H2O)n-1(OH)c-1 + H+

Chapter 16 : Acid-Base Equilibria