PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 15 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 15 SOLUTIONS OF 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, Metal ions Can accept four or six pairs of electrons from Lewis bases Fe H 2 O(l) → Fe(H 2 O) 6 3+

- 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 + )

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) CONJUGATE ACID BASE PAIRS

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 +

AUTOPROTOLYSIS OF WATER H 2 O + H 2 OH 3 O + + OH - KwKw - Autoionization (self-ionization) of water - 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 (25 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 (K w ) = [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 - Commonly expressed to 2 decimal places (2 significant figures) 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 = THE pH CONCEPT

- 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

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 ionizes 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 NaOH(aq) → Na + (aq) + OH - (aq) [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

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

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

- 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 pH OF STRONG ACIDS AND BASES

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 KaKa WEAK ACID EQUILIBRIUM

- 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 - KbKb WEAK BASE EQUILIBRIUM

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 KbKb WEAK BASE EQUILIBRIUM

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

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

HYDROLYSIS OF SALTS - Reaction of salt with water to produce hydronium ion or hydroxide ion or both (do not go to 100% completion) - Not all salts hydrolyze - The salt of a strong acid and a strong base does not hydrolyze - Neutral solution is the result - The salt of a strong acid and a weak base hydrolyzes - Acidic solution is the result

- The salt of a weak acid and a strong base hydrolyzes - Basic solution is the result - The salt of a weak acid and a weak base hydrolyzes - Slightly acidic, neutral, or basic, depending on relative weaknesses of acid and base HYDROLYSIS OF SALTS

Acidic Hydrolysis positive ion of salt + H 2 O Conjugate base + H 3 O + - The hydronium ion makes the solution acidic NH H 2 O → NH 3 + H 3 O + HYDROLYSIS OF SALTS

Basic Hydrolysis negative ion of salt + H 2 O Conjugate acid + OH - - The hydroxide ion makes the solution basic F - + H 2 O → HF + OH - HYDROLYSIS OF SALTS

- 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