ANALYTICAL CHEMISTRY CHEM 3811 CHAPTER 8 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 8 ACIDS AND BASES

ARRHENIUS ACIDS - Acids are substances that ionize in aqueous solutions to produce hydrogen ions (proton, H + ) HCl, HNO 3, H 2 SO 4 - Arrhenius acids are covalent compounds in the pure state Properties sour taste, change blue litmus paper to red, corrosive

ARRHENIUS BASES - Bases are substances that ionize in aqueous solutions to produce hydroxide ions (OH - ) NaOH, KOH, Ca(OH) 2 - Arrhenius bases are ionic compounds in the pure state Properties bitter taste, change red litmus paper to blue, slippery to touch

BRONSTED-LOWRY ACIDS - Acids are proton (H + ) donors - Not restricted to aqueous solutions HCl, HNO 3, H 2 SO 4

- Bases are proton acceptors - Not restricted to aqueous solutions NH 3, dimethyl sulfoxide (DMSO) - Proton donation cannot occur unless an acceptor is present BRONSTED-LOWRY BASES

LEWIS ACIDS - Acids are electron pair acceptors - Not restricted to protons or aqueous solutions BF 3, B 2 H 6, Al 2 Cl 6, AlF 3, PCl 5

- Bases are electron pair donors - Not restricted to protons or aqueous solutions NH 3, ethers, ketones, carbon monoxide, sulfoxides - The product of a Lewis acid-base reaction is known as an adduct - The base donates an electron pair to form coordinate covalent bond LEWIS BASES

ACIDS Monoprotic Acid - Donates one proton per molecule (HNO 3, HCl) Diprotic Acid - Donates two protons per molecule (H 2 SO 4, H 2 CO 3 ) Triprotic Acid - Donates three proton per molecule (H 3 PO 4, H 3 AsO 4 ) Polyprotic Acid - Donates two or more protons per molecule

CONJUGATE ACID BASE PAIRS - Most Bronsted-Lowry acid-base reactions do not undergo 100% conversion - Acid-base equilibrium is established - Every acid has a conjugate base associated with it (by removing H + ) - Every base has a conjugate acid associated with it (by adding H + )

CONJUGATE ACID BASE PAIRS HX(aq) + H 2 O(l)X - (aq) + H 3 O + (aq) - HX donates a proton to H 2 O to form X - HX is the acid and X - is its conjugate base - H 2 O accepts a proton from HX H 2 O acts as a base and H 3 O + is its conjugate acid

CONJUGATE ACID BASE PAIRS NH 3 (aq) + H 2 O(l)NH 4 + (aq) + OH - (aq) HF(aq) + H 2 O(l)H 3 O + (aq) + F - (aq) HNO 3 (aq) + H 2 O(l)H 3 O + (aq) + NO 3 - (aq)

AMPHOTERIC SUBSTANCES - A substance that can lose or accept a proton - A substance that can function as either Bronsted-Lowry acid or Bronsted-Lowry base - H 2 O is the most common (refer to previous slide for examples)

REACTIONS OF ACIDS AND BASES Arrhenius acid + Arrhenius base → salt + water HCl + NaOH → NaCl + H 2 O B-L acid + B-L base → conjugate base + conjugate acid H 3 PO 4 + H 2 O → H 2 PO H 3 O +

SALTS - Salts are ionic compounds - The positive ion is a metal or polyatomic ion - The negative ion is a nonmetal or polyatomic ion [exception is the hydroxide ion (OH - )] - Salts dissociate completely into ions in solution - A reaction between an acid and a hydroxide base produces salt (cation from the base and anion from the acid)

SALTS - Solutions of salts may be acidic, basic, or neutral - Acidity depends on relative values of K a of the cation and K b of the anion - The conjugate base of a strong acid (anion from a strong acid) has no net effect on the pH of a solution (spectator ion) Cl - from HCl, NO 3 - from HNO 3 - Cation from a strong base has no net effect on the pH of a solution (spectator ion) Na + from NaOH, K + from KOH

SALTS - NaCl solution contains Na + and Cl - ions - Both ions are spectator ions and do not affect the pH of the solution - pH is determined by autoionization of water

AUTOPROTOLYSIS OF WATER H 2 O + H 2 OH 3 O + + OH - KwKw - Pure water molecules (small percentage) interact with one another to form equal amounts of H 3 O + and OH - ions reduces to H + + OH - H2OH2O KwKw

- The number of H 3 O + and OH - ions present in a sample of pure water at any given time is small - At equilibrium (24 o C) [H 3 O + ] = [OH - ] = 1.00 x M - [H 3 O + ] = hydronium ion concentration - [OH - ] = hydroxide ion concentration AUTOPROTOLYSIS OF WATER

- The ion product constant of water = [H 3 O + ] x [OH - ] = (1.00 x ) x (1.00 x ) = 1.00 x Valid in all solutions (pure water and water with solutes) AUTOPROTOLYSIS OF WATER

Addition of Acidic Solute - increases [H 3 O + ] - [OH - ] decreases by the same factor to make product 1.00 x Addition of Basic Solute - increases [OH - ] - [H 3 O + ] decreases by the same factor to make product 1.00 x AUTOPROTOLYSIS OF WATER

Acidic Solution - An aqueous solution in which [H 3 O + ] is higher than [OH - ] Basic Solution - An aqueous solution in which [OH - ] is higher than [H 3 O + ] Neutral Solution - An aqueous solution in which [H 3 O + ] is equal to [OH - ] THE pH CONCEPT

pH - Negative logarithm of the hydronium ion concentration [H 3 O + ] in an aqueous solution pH = - log[H 3 O + ] [H 3 O + ] = 10 -pH

THE pH CONCEPT - For [H 3 O + ] coefficient of Expressed in exponential notation - The pH is the negative of the exponent value [H 3 O + ] = 1.0 x M, then pH = 5.00 [H 3 O + ] = 1.0 x M, then pH = 3.00 [H 3 O + ] = 1.0 x M, then pH = 11.00

- For neutral solutions pH is equal to For acidic solutions pH is less than For basic solutions pH is greater than Increasing [H 3 O + ] lowers the pH THE pH CONCEPT

- A change of 1 unit in pH corresponds to a tenfold change in [H 3 O + ] pH = 3.00 implies [H 3 O + ] = 1.0 x M = M pH = 2.00 implies [H 3 O + ] = 1.0 x M = M which is tenfold - The pH meter and the litmus paper are used to determine pH values of solutions THE pH CONCEPT

pK w = -log(K w ) = -log(1.00 x ) = 14 pOH = -log[OH - ] [H 3 O + ][OH - ] = K w Implies that pH + pOH = pK w pH + pOH = THE pH CONCEPT

STRENGTH OF ACIDS Strong Acids - Transfer 100% (or very nearly 100%) of their protons to H 2 O in aqueous solution - Completely or nearly completely ionize in aqueous solution - Strong electrolytes HCl, HBr, HClO 4, HNO 3, H 2 SO 4 Weak Acids - Transfer only a small percentage (< 5%) of their protons to H 2 O in aqueous solution Amino acids, Organic acids: acetic acid, citric acid

- Equilibrium position lies to the far right for strong acids HA(aq) + H 2 O(l)H 3 O + (aq) + A - (aq) - Equilibrium position lies to the far left for weak acids HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) - Predominant species are H 3 O + and A - - Predominant species is HA STRENGTH OF ACIDS

- Equilibrium constant for the reaction of a weak acid with water - Represented by K a (acid dissociation constant) HA(aq) + H 2 O(l)H 3 O + (aq) + A - (aq) - H 2 O is a pure liquid so not included - Acid strength increases with increasing K a value - For polyprotic acids, K a for each dissociation step is smaller than the previous step (weaker acid) STRENGTH OF ACIDS

Strong Bases - Completely or nearly completely ionize in aqueous solution - Strong electrolytes Hydroxides of Groups IA and IIA are strong bases LiOH, CsOH, Ba(OH) 2, Ca(OH) 2 most common in lab: NaOH and KOH Weak bases - produce small amounts of OH - ions in aqueous solution Organic bases, methylamine, cocaine, morphine most common: NH 3 STRENGTH OF BASES

- Weak bases produce small amounts of OH - ions in aqueous solution (NH 3 ) NH 3 (g) + H 2 O(l)NH 4 + (aq) + OH - (aq) - Equilibrium position lies to the far left - Small amounts of NH 4 + and OH - ions are produced - The name aqueous ammonia is preferred over ammonium hydroxide STRENGTH OF BASES

- Equilibrium constant for the reaction of a weak base with water - Represented by K b (base hydrolysis constant) B(aq) + H 2 O(l)BH + (aq) + OH - (aq) - H 2 O is a pure liquid so not included STRENGTH OF BASES

K a x K b = [H 3 O + ][OH - ] = K w = 1.00 x Reaction goes to completion when K a value is very large - Weak acids have small K a values WEAK ACIDS AND BASES

pK a = - logK a pK b = - logK b pK a + pK b = pK w - The stronger an acid the smaller its pK a - The stronger the acid the weaker its conjugate base - The stronger the base the weaker its conjugate acid

METAL IONS - Metal ions with +1 charge have negligible acidity - Metal ions with +2 charge or higher are acidic - The higher the charge the more acidic the metal Example Fe 2+ : K a = 4 x , Fe 3+ : K a = 6.5 x Solutions of metal salts are usually acidic - Metal ions are Lewis acids

METAL IONS - Many metal ions bind with four or six water molecules M(H 2 O) 6 n+ KaKa M(H 2 O) 5 (OH) (n-1)+ + H + M n+ + 6H 2 O M(H 2 O) 6 n+

pH OF STRONG ACIDS - Differences in acidities of strong acids cannot be measured since they all ionize completely - This phenomenon is known as leveling effect Find the pH of 3.9 x M HCl HCl is a strong acid and dissociates completely HCl(aq) → H + (aq) + Cl - (aq) pH = - log(3.9 x ) = 1.41

pH OF STRONG BASES Find the pH of 3.9 x M NaOH [H 3 O + ][OH - ] = K w = 1.0 x [H 3 O + ][3.9 x ] = 1.0 x [H 3 O + ] = 2.6 x pH = - log(2.6 x ) = 12.59

pH OF STRONG BASES Find the pH of 3.9 x M NaOH Alternatively pOH = - log[OH - ] pOH = - log(3.9 x ) = 1.41 pH + pOH = 14 pH = = 12.59

pH OF STRONG ACIDS AND BASES - For dilute solutions the contribution of H 2 O should not be neglected - Acids and bases suppress water ionization What concentrations of H + and OH - are produced by H 2 O dissociation in 1.0 x M HCl? pH = 3 [OH - ] = K w /[H 3 O + ] = 1.0 x OH - is produced from the dissociation of H 2 O Implies H 2 O dissociation = [OH - ] = [H 3 O + ] = 1.0 x

pH OF STRONG ACIDS AND BASES - For dilute solutions the contribution of H 2 O should not be neglected - Acids and bases suppress water ionization What concentrations of H + and OH - are produced by H 2 O dissociation in 1.0 x M KOH? [H 3 O + ] = K w /[OH - ] = 1.0 x H 3 O + (or H + ) is produced from the dissociation of H 2 O Implies H 2 O dissociation = [OH - ] = [H 3 O + ] = 1.0 x

WEAK ACID EQUILIBRIUM For a weak acid HA HAA - + H + c HA = total concentration = analytical concentration = [HA] + [A - ] KaKa

WEAK ACID EQUILIBRIUM For a weak acid HA HAA - + H + - Fraction of dissociation increases with increasing acid strength - Fraction of dissociation increases with dilution KaKa

For a weak acid HA HAA - + H + - Assume [H + ] ≈ [A - ] - F is the initial (formal) concentration of HA - Initial concentration of H + and A - is 0 each - Final concentration of H + and A - is x each - The iCe table may be used for such problems WEAK ACID EQUILIBRIUM KaKa

- The equation reduces to WEAK ACID EQUILIBRIUM - If x ≤ 5% of F That is F – x ≈ F if x ≤ 0.05F

For a weak base B B + H 2 OBH + + OH - WEAK BASE EQUILIBRIUM KbKb

For a weak base B B + H 2 OBH + + OH - - Assume [BH + ] ≈ [OH - ] - F is the initial (formal) concentration of B - Initial concentration of BH + and OH - is 0 each - Final concentration of BH + and OH - is x each - The iCe table may be used for such problems WEAK BASE EQUILIBRIUM KbKb

- The equation reduces to WEAK BASE EQUILIBRIUM - If x ≤ 5% of F That is F – x ≈ F if x ≤ 0.05F

- When determining the pH of a mixture of acids only the pH of the strongest acid is considered - Contributions by the weaker acids towards pH are neglected - A weak acid produces fewer protons in the presence of a strong acid Similarly - A weak base produces fewer hydroxide ions in the presence of a strong base MIXTURES OF ACIDS

- Key factors are the strength of the H – A bond and the stability of the A - ion Binary Acid (HA) - An acidic compound composed of hydrogen and one other element (mostly a nonmetal) HCl, HI, HBr, H 2 S, H 2 O FACTORS AFFECTING STRENGTH OF ACIDS

Bond Strength of Binary Acids - Generally decreases down the groups of the periodic table - Due to increasing size of the other element - Acidity increases down the groups of the periodic table - Due to decreasing bond strength FACTORS AFFECTING STRENGTH OF ACIDS

Example Bond strength of hydrogen halides HF > HCl > HBr > HI Acidity of hydrogen halides HF < HCl < HBr < HI FACTORS AFFECTING STRENGTH OF ACIDS

Stability of the A - Anion - Depends on the ability of the A atom to accept additional negative charge - Electronegativity is the factor - A more electronegative atom results in a stronger acid - Acidity of nonmetal hydrides increases across periods of the periodic table CH 4 < NH 3 < H 2 O < HF FACTORS AFFECTING STRENGTH OF ACIDS

- Bond strength and electronegativity sometimes predict opposite trends - Bond strength dominates down a group - Electronegativity dominates across a period FACTORS AFFECTING STRENGTH OF ACIDS

Oxyacids - Acids containing hydrogen, oxygen, and a third element The third element may be a - Nonmetal: HNO 3, H 2 SO 4, H 3 PO 4 - A transition metal with high oxidation state: H 2 CrO 4 - Carbon in organic acids: CH 3 COOH - Acidity increases with electronegativity of the third element - Hypohalous acids (H – O – X), X = Cl, Br, I FACTORS AFFECTING STRENGTH OF ACIDS