Chemistry 102(01) spring 2009 Instructor: Dr. Upali Siriwardane e-mail: upali@chem.latech.edu Office: CTH 311 Phone 257-4941 Office Hours: M,W, 8:00-9:00 & 11:00-12:00 a.m.; Tu,Th,F 9:00 - 10:00 a.m.   Test Dates: March 25, April 26, and May 18; Comprehensive Final Exam: May 20,2009 9:30-10:45 am, CTH 328. March 30, 2009 (Test 1): Chapter 13 April 27, 2009 (Test 2): Chapters 14 & 15 May 18, 2009 (Test 3): Chapters 16, 17 & 18 Comprehensive Final Exam: May 20,2009 :Chapters 13, 14, 15, 16, 17 and 18

Chapter 16. Acids and Bases 16.1 The Brønsted-Lowry Concept of Acids and Bases 16.2 Types of acids/bases:Organic Acids and Amines 16.3 The Autoionization of Water 16.4 The pH Scale 16.5 Ionization Constants of Acids and Bases 16.6 Problem Solving Using Ka and Kb 16.7 Molecular Structure and Acid Strength 16.8 Acid-Base Reactions of Salts 16.9 Practical Acid-Base Chemistry 16.10 Lewis Acid and Bases

Types of Reactions a) Precipitation Reactions. Reactions of ionic compounds or salts b) Acid/base Reactions. Reactions of acids and bases c) Redox Reactions. reactions of oxidizing & reducing agents

What are Acids &Bases? Definition? a) Arrhenius b) Bronsted-Lowry c) Lewis

Arrhenius Definitions Arrhenius, Svante August (1859-1927), Swedish chemist, 1903 Nobel Prize in chemistry Acid Anything that produces hydrogen ions in a water solution. HCl (aq) H+ ( aq) + Cl- ( aq) Base Anything that producs hydroxide ions in a water solution. NaOH (aq) Na+ ( aq) + OH- ( aq) Arrhenius definitions are limited proton acids and hydroxide bases to aqueous solutions.

Brønsted-Lowry definitions Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions Acid Proton donor Base Proton acceptor This definition explains how substances like ammonia can act as bases. Eg. HCl(g) + NH3(g) ------> NH4Cl(s) HCl (acid), NH3 (base). NH3(g) + H2O(l) NH4+ + OH-

Lewis Definition G.N. Lewis was successful in including acid and bases without proton or hydroxyl ions. Lewis Acid: A substance that accepts an electron pair. Lewis base: A substance that donates an electron pair. E.g. BF3(g) + :NH3(g) F3B:NH3(s) the base donates a pair of electrons to the acid forming a coordinate covalent bond common to coordination compounds. Lewis acids/bases will be discussed later in detail

Dissociation Strong Acids: HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) Dissociation Equilibrium Weak Acid/base: H2O(l) + H2O(l) H3+O(aq) + OH-(aq) This dissociation is called autoionization of water. HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq) NH3 (aq) + H2O(l) NH4+ + OH-(aq) Equilibrium constants: Ka, Kb and Kw

Brønsted-Lowry Definitions Conjugate acid-base pairs. Acids and bases that are related by loss or gain of H+ as H3O+ and H2O. Examples. Acid Base H3O+ H2O HC2H3O2 C2H3O2- NH4+ NH3 H2SO4 HSO4- HSO4- SO42-

Bronsted acid/conjugate base and base/conjugate acid pairs in acid/base equilibria HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) HCl(aq): acid H2O(l): base H3+O(aq): conjugate acid Cl-(aq): conjugate base H2O/ H3+O: base/conjugate acid pair HCl/Cl-: acid/conjugate base pair

Select acid, base, acid/conjugate base pair, base/conjugate acid pair H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4-(aq) acid base conjugate acid conjugate base base/conjugate acid pair acid/conjugate base pair

Types of Acids and Bases Binary acids: HCl, HBr, HI, H2S More than two elements: HCN Oxyacid: HNO3, H2SO4, H3PO4 Polyprotic acids: H2SO4, H3PO4 Organic acids: R-COOH, R= CH3-, CH3CH2- Acidic oxides: SO3, NO2, CO2, Basic oxides: Na2O, CaO Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary R2-NH : secondary, R3-N: tertiary Lewis acids & bases: BF3 and NH3

Strong Acid vs. Weak Acids completely ionized Hydrioidic HI Ka ~ 1011 pKa = -11 Hydrobromic HBr Ka ~ 109 pKa = -9 Perchloric HClO4 Ka ~ 107 pKa = -7 Hyrdrochloric HCl Ka ~ 107 pKa = -7 Chloric HClO3 Ka ~ 103 pKa = -3 Sulfuric H2SO4 Ka ~ 102 pKa = -2 Nitric HNO3 Ka ~ 20 pKa = -1.3 Weak acid partially ionized Hydrofluoric acid HF Ka = 6.6x10-4 pKa = 3.18 Formic acid HCOOH Ka = 1.77x10-4 pKa = 3.75 Acetic acid CH3COOH Ka = 1.76x10-5 pKa = 4.75 Nitrous acid HNO2 Ka = 4.6x10-4 pKa = 3.34 Acetyl Salicylic acid C9H8O4 Ka = 3x10-4 pKa = 3.52 Hydrocyanic acid HCN Ka = 6.17x10-10 pKa = 9.21

Strong Base vs. Weak Base completely ionized Lithium hydroxide LiOH Sodium hydroxide NaOH Potasium hydroxide KOH Kb~ 102-103 Rubidium hydroxide RbOH Cesium hydroxide CsOH Boarder-line Bases Magnesium hydroxide Mg(OH)2 Calcium hydroxide Ca(OH)2 Strotium hydroxide Sr(OH)2 Kb~ 0.01 to0.1 Barium hydroxide Ba(OH)2 Weak Base partially ionized Ammonia NH3 Kb=1.79x10-5 pKb = 4.74 Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25

Acid and Base Strength Strong acids Ionize completely in water. HCl, HBr, HI, HClO3, HNO3, HClO4, H2SO4. Weak acids Partially ionize in water. Most acids are weak. Strong bases Ionize completely in water. Strong bases are metal hydroxides - NaOH, KOH Weak bases Partially ionize in water.

Common Acids and Bases Acids Formula Molarity* nitric HNO3 16 hydrochloric HCl 12 sulfuric H2SO4 18 acetic HC2H3O2 18 Bases ammonia NH3(aq) 15 sodium hydroxide NaOH solid *undiluted.

Autoionization of Water Autoionization When water molecules react with one another to form ions. Acids and bases alter the dissociation equilibrium of water based on Le Chaterlier’s principle Kw = [ H3O+ ] [ OH- ] = 1.0 x 10-14 at 25oC Note: [H2O] is constant and is included in Kw. H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M) ion product of water

pH and other “p” scales Substance pH 1 M HCl 0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0 We need to measure and use acids and bases over a very large concentration range. pH and pOH are systems to keep track of these very large ranges. pH = -log[H3O+] pOH = -log[OH-] pH + pOH = 14

pH scale A logarithmic scale used to keep track of the large changes in [H+]. 0 7 14 10-14 M 10-7 M 10-14 M Very Neutral Very acidic Basic When you add an acid to, the pH gets smaller. When you add a base to, the pH gets larger.

pH of some common materials Substance pH 1 M HCl 0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0

pH of Aqueous Solutions

What is pH? Kw = [H3+O][OH-] = 1 x 10-14 [H3+O][OH-] = 10-7 x 10-7 Extreme cases: Basic medium [H3+O][OH-] = 10-14 x 100 Acidic medium [H3+O][OH-] = 100 x 10-14 pH value is -log[H+] spans only 0-14 in water.

pH, pKw and pOH The relation of pH, Kw and pOH Kw = [H+][OH-] log Kw = log [H+] + log [OH-] -log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14 since Kw =1 x 10-14 14 = pH + pOH pH = 14 - pOH pOH = 14 - pH

pH and pOH calculations of acid and base solutions a) Strong acids/bases dissociation is complete for strong acid such as HNO3 or base NaOH [H+] is calculated from molarity (M) of the solution b) weak acids/bases needs Ka , Kb or percent(%)dissociation

pH of Strong Acid/bases Substance pH 1 M HCl 0.0 Gastric juices 1.0 - 3.0 Lemon juice 2.2 - 2.4 Classic Coke 2.5 Coffee 5.0 Pure Water 7.0 Blood 7.35 - 7.45 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0 HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq) Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning. [HNO3] = [H+] = 0.2 mole/L pH = -log [H+] = -log(0.2) pH = 0.699

pH of 0.5 M H2SO4 Solution HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) [H3+O][HSO4-] H2SO4 ; Ka1 = ------------------- [H2SO4] [H3+O][SO42-] H2SO4 ; Ka2 = ------------------- ; Ka2 ignored [HSO4-]

pH of 0.5 M H2SO4 Solution pH = -log(0.5) pH = 0.30 H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning. [H2SO4] = [H+] = 0.5 mole/L pH = -log [H+] pH = -log(0.5) pH = 0.30

1.5 x 10-2 M NaOH. 1.5 x 10-2 M NaOH. NaOH is also a strong base dissociates completely in water. [NaOH] = [HO- ] = 1.5 x 10-2 mole/L pOH = -log[HO-]= -log(1.5 x 10-2) pOH = 1.82 As defined and derived previously: pKw= pH + pOH; pKw= 14 pH = pKw + pOH pH = 14 - pOH pH = 14 - 1.82 ; pH = 12.18

Mixtures of Strong and Weak Acids the presence of the strong acid retards the dissociation of the weak acid

Measuring pH Arnold Beckman inventor of the pH meter father of electronic instrumentation

Equilibrium, Constant, Ka & Kb Ka: Acid dissociation constant for a equilibrium reaction. Kb: Base dissociation constant for a equilibrium reaction. Acid: HA + H2O H3+O + A- Base: BOH + H2O B+ + OH- [H3+O][ A-] [B+ ][OH-] Ka = --------------- ; Kb = ----------------- [HA] [BOH]

Acid Dissociation Constant HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq) [H3+O][Cl-] Ka= ----------------- [HCl] [H+][Cl-]

Base Dissociation Constant NH3 + H2O NH4+ + OH- [NH4+][OH-] K = [NH3]

Hydrated Metal Ions as Acids

Ionization Constants for Acids

Comparing Kw and Ka & Kb Any compound with a Ka value greater than Kw of water will be a an acid in water. Any compound with a Kb value greater than Kw of water will be a base in water.

WEAKER/STRONGER Acids and Bases & Ka and Kb values A larger value of Ka or Kb indicates an equilibrium favoring product side. Acidity and basicity increase with increasing Ka or Kb. pKa = - log Ka and pKb = - log Kb Acidity and basicity decrease with increasing pKa or pKb.

Which is weaker? a. HNO2 ; Ka= 4.0 x 10-4. b. HOCl2 ; Ka= 1.2 x 10-2. c. HOCl     ;  Ka= 3.5 x 10-8. d. HCN      ;  Ka= 4.9 x 10-10.

What is Ka1 and Ka2? H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq)

Ka Examples H2SO4(aq) + H2O(l) H3+O(aq) + HSO4-(aq) HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) [H3+O][HSO4-] H2SO4 ; Ka1 = ------------------- [H2SO4] [H3+O][SO42-] H2SO4 ; Ka2 = ------------------- [HSO4-]

Ka Examples HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2-(aq) [H+][C2H3O2-] H C2H3O2; Ka= ------------------ [H C2H3O2] NH3 (aq) + H2O(l) NH4+ + OH-(aq) [NH4+][OH-] NH3; Kb= -------------- [ NH3]

How do you calculate pH of weak acids/bases From % dissociation From Ka or Kb What is % dissociation Amount dissociated % Dissoc. = ------------------------- x 100 Initial amount

How do you calculate % dissociation from Ka or Kb 1.00 M solution of HCN; Ka = 4.9 x 10-10 What is the % dissociation for the acid?

1.00 M solution of HCN; Ka = 4.9 x 10-10 First write the dissociation equilibrium equation: HCN(aq) + H 2O(l) <===> H 3+O(aq) + CN-(aq) [HCN] [H+ ] [CN- ] Ini. Con. 1.00 M 0.0 M 0.00 M Cha. Con -x x x Eq. Con. 1.0 - x x x [H 3+O ][CN-] x2 Ka = ----------- = ---------------- [HCN] 1.0 - x

1.00 M solution of HCN; Ka = 4.9 x 10-10 1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2 1.0 x = 4.9 x 10-10 = 2.21 x 10 -5 Amount disso. 2.21 x 10 -5 ----------------- x 100 =- ------------- x 100 Ini. amount 1.00 % Diss. =2.21 x 10 -5 x 100 = 0.00221 %

% Dissociation gives x (amount dissociated) need for pH calculation % Dissoc. = ------------------------- x 100 Initial amount/con. x % Dissoc. = --------------------------- x 100 concentration

Calculate the pH of a weak acid from % dissociation 1 M HF, 2.7% dissociated Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027

Calculate the pH of a weak acid from % dissociation HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq) [H+][F-] Ka = ----------- [HF] [HF] [H+ ] [F- ] Ini. Con. 1.00 M 0.0 M 0.00 M Chg. Con -x x x Eq.Con. 1.0-0.027 0.027 0.027 pH = -log [H+] pH = -log(0.027) pH = 1.57

HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq) Weak acid Equilibria Example Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = 6.28 x 10-5 HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq) The first step is to write the equilibrium expression Ka = [H3O+][Bz-] [HBz]

Weak acid Equilibria HBz H3O+ Bz- Initial conc., M 0.10 0.00 0.00 Change, DM -x x x Eq. Conc., M 0.10 - x x x [H3O+] = [Bz-] = x We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.

Weak Acid Equilibria Solve the equilibrium equation in terms of x Ka = 6.28 x 10-5 = x = (6.28 x 10-5 )(0.10) H3O+ = 0.0025 M pH = 2.60 x2 0.10

pH from Ka or Kb 1.00 M solution of HCN; Ka = 4.9 x 10-10 First write the dissociation equilibrium equation: HCN(aq) + H 2O(l) H 3+O(aq) + CN-(aq) [HCN] [H+ ] [CN- ] Ini. Con. 1.00 M 0.0 M 0.00 M Chg. Con -x x x Eq. Con. 1.0 - x x x

Weak Acid Equilibria [H 3+O ][CN-] x2 1.0 - x ~ 1.00 since x is small [HCN] 1.0 - x 1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2 1.0 x = 4.9 x 10-10 = 2.21 x 10 -5 pH = -log [H+] pH = -log(2.21 x 10-5) pH = 4.65

The Conjugate Partners of Strong Acids and Bases The conjugate acid/base of a strong base/acid has no net effect on the pH of a solution The conjugate base of a weak acid hydrolyze in water and basic or pH of a solution > 7.00 E.g. Na+C2H3O2- sodium acetate The conjugate acid of a weak base hydrolyze in water and acidic or pH of a solution < 7.00 E.g NH4Cl

Hydrolysis Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-Lowry acid-base reaction. CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq) NH4+(aq) + H2O(l) NH3 (aq) + H3O+(aq) This type of reaction is given a special name. Hydrolysis The reaction of an anion with water to produce the conjugate acid and OH-. The reaction of a cation with water to produce the conjugate base and H3O+.

Acid-Base Properties of Typical Ions

What salt solutions would be acidic, basic and neutral? 1) strong acid + strong base = neutral 2) weak acid + strong base = basic 3) strong acid + weak base = acidic weak acid + weak base = neutral, basic or an acidic solution depending on the relative strengths of the acid and the base.

What pH? Neutral, basic or acidic? a)NaCl neutral b) NaC2H3O2 basic c) NaHSO4 acidic d) NH4Cl

How do you calculate pH of a salt solution? Find out the pH, acidic or basic? If acidic it should be a salt of weak base If basic it should be a salt of weak acid if acidic calculate Ka from Ka= Kw/Kb if basic calculate Kb from Kb= Kw/Ka Do a calculation similar to pH of a weak acid or base

What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5) Find out the pH, acidic if acidic calculate Ka from Ka= Kw/Kb Ka= Kw/Kb = 1 x 10-14 /1.8 x 10-5) Ka= 5.56. X 10-10 Do a calculation similar to pH of a weak acid

Continued NH4+ + H2O H 3+O + NH3 [NH4+] [H3+O ] [NH3 ] Ini. Con. 0.5 M 0.0 M 0.00 M Change -x x x Eq. Con. 0.5 - x x x [H 3+O ] [NH3 ] Ka(NH4+) = -------------------- = [NH 4+] x2 ---------------- ; appro.:0.5 - x . 0.5 (0.5 - x)

Continued x2 Ka(NH4+) = ----------- = 5.56 x 10 -10 0. 5 x2 = 5.56 x 10 -10 x 0.5 = 2.78 x 10 -10 x= 2.78 x 10 -10 = 1.66 x 10-5 [H+ ] = x = 1.66 x 10-5 M pH = -log [H+ ] = - log 1.66 x 10-5 pH = 4.77 pH of 0.5 M NH4Cl solution is 4.77 (acidic)

Types of Acids and Bases Binary acids Oxyacid Organic acids Acidic oxides Basic oxides Amine Polyprotic acids

Influence of Molecular Structure on Acid Strength Binary Hydrides hydrogen & one other element Bond Strengths weaker the bond, the stronger the acid Stability of Anion higher the electronegativity, stronger the acid

Binary Acids Compounds containing acidic protons bonded to a more electronegative atom. e.g. HF, HCl, HBr, HI, H2S The acidity of the haloacid (HX; X = Cl, Br, I, F) Series increase in the following order: HF < HCl < HBr < HI

Oxyacids Compounds containing acidic - OH groups in the molecule. Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens) The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend. HClO4 > HClO3 > HClO2 > HClO perchloric chloric chlorus hyphochlorus

Influence of Molecular Structure on Acid Strength Oxyacids hydrogen, oxygen, & one other element H-O-E higher the electronegativity on E, stronger the acid as this weakens the bond between the O and H

Oxo Acid < < < <

Acidic Oxides These are usually oxides of non-metallic elements such as P, S and N. E.g. NO2, SO2, SO3, CO2 They produce oxyacids when dissolved in water SO3 + H2O ---> H2SO4 CO2 + H2O ---> H2CO3 NO2 + H2O ---> HNO3

Basic Oxides Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water. e.g. CaO + H2O ---> Ca(OH)2 Na2O + H2O ---> 2 NaOH

Protic Acids Monoprotic Acids: The form protic refers to acidity due to protons. Monoprotic acids have only one acidic proton. e.g. HCl. Polyprotic Acids: They have more than one acidic proton. e.g. H2SO4 - diprotic acid H3PO4 - triprotic acid.

Polyprotic Acids acids where more than one hydrogen per molecule is released

Polyprotic Acids

Organic or Carboxylic Acids

Organic or Carboxylic Acids FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid). Acid Ka pKa HCOOH (formic acid) 1.78 X 10-43 0.75 CH3COOH (acetic acid) 1.74 X 10-54 0.76 CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86

Amines Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted-Lowery or Lewis acid/base definitions.

Amines

Acid-Base Chemistry of Some Antacids

Acid-Base in the Kitchen vinegar - acetic acid lemon juice (citrus juice) - citric acid baking soda - NaHCO3 milk - lactic acid baking powder - H2PO4- & HCO3-

Household Cleaners

Dishwashing Detergent

Lewis Definition G.N. Lewis was successful in including acid and bases without proton or hydroxyl ions. Lewis Acid: A substance that accepts an electron pair. Lewis base: A substance that donates an electron pair. E.g. BF3(g) + :NH3(g) F3B:NH3(s) the base donates a pair of electrons to the acid forming a coordinate covalent bond common to coordination compounds. Lewis acids/bases will be discussed later in detail

Lewis Acids and Bases Reactions H+ + NH3  NH4+ acid base Cu+2 + 4 NH3  [Cu(NH3)4+2] acid base

What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions: a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l) b)HCl(g) + NH3(g) ---> NH4Cl(s) c)BF3(g) + NH3(g) ---> F3B:NH3(s) d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)