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. NH3(g) + H2O(l) NH4+ + OH- Eg. HCl(g) + NH3(g) ------> NH4Cl(s) HCl (acid), NH3 (base). 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 C2H3O2NH4+ 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) acid base conjugate acid conjugate base base/conjugate acid pair acid/conjugate base pair H 3+O(aq) + HSO4-(aq) 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-, CH3CH2Acidic 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 Strong acid completely ionized Hydrioidic Hydrobromic Perchloric Hyrdrochloric Chloric Sulfuric Nitric HI HBr HClO4 HCl HClO3 H2SO4 HNO3 Ka ~ 1011 Ka ~ 109 Ka ~ 107 Ka ~ 107 Ka ~ 103 Ka ~ 102 Ka ~ 20 pKa = -11 pKa = -9 pKa = -7 pKa = -7 pKa = -3 pKa = -2 pKa = -1.3 Weak acid partially ionized Hydrofluoric acid HF Formic acid HCOOH Acetic acid CH3COOH Nitrous acid HNO2 Acetyl Salicylic acid C9H8O4 Hydrocyanic acid HCN Ka = 6.6x10-4 Ka = 1.77x10-4 Ka = 1.76x10-5 Ka = 4.6x10-4 Ka = 3x10-4 Ka = 6.17x10-10 pKa = 3.18 pKa = 3.75 pKa = 4.75 pKa = 3.34 pKa = 3.52 pKa = 9.21 Strong Base vs. Weak Base Strong Base completely ionized Lithium hydroxide Sodium hydroxide Potasium hydroxide Rubidium hydroxide Cesium hydroxide LiOH NaOH KOH RbOH CsOH Boarder-line Bases Magnesium hydroxide Mg(OH)2 Calcium hydroxide Ca(OH)2 Strotium hydroxide Sr(OH)2 Barium hydroxide Ba(OH)2 Weak Base Kb~ 102-103 Kb~ 0.01 to0.1 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 nitric hydrochloric sulfuric acetic Bases ammonia sodium hydroxide *undiluted. Formula Molarity* HNO3 16 HCl 12 H2SO4 18 HC2H3O2 18 NH3(aq) NaOH 15 solid Autoionization of Water Autoionization When water molecules react with one another to form ions. H2O(l) + H2O(l) H3O+(aq) + OH-(aq) (10-7M) (10-7M) 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 ion product of water pH and other “p” scales Substance pH need to measure and use 0.0 acids and bases 1 MWe HCl over a very large concentration range. Gastric juices 1.0 - 3.0 Lemon juice 2.2 -track 2.4 of pH and pOH are systems to keep Classic Coke these very large ranges. 2.5 CoffeepH = -log[H3O+5.0 ] Pure Water pOH = -log[OH-]7.0 BloodpH + pOH 7.35 - 7.45 = 14 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0 pH scale A logarithmic scale used to keep track of the large changes in [H+]. 0 7 10-14 M Very acidic Basic 10-7 M 10-14 M Neutral 14 Very 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 Gastric juices Lemon juice Classic Coke Coffee Pure Water Blood Milk of Magnesia Household ammonia 1M NaOH 0.0 1.0 - 3.0 2.2 - 2.4 2.5 5.0 7.0 7.35 - 7.45 10.5 12.0 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 HNO (aq) + H2O(l) H3+O(aq) 1 M 3HCl 0.0+ NO3-(aq) Gastric juices - 3.0 Therefore, the moles of H+ ions 1.0 in the solution is Lemon - 2.4 equal tojuice moles of HNO3 at the 2.2 beginning. Classic [HNO3Coke ] = [H+] = 0.2 mole/L 2.5 Coffee 5.0 pH = -log [H+] Pure Water 7.0 = -log(0.2) Blood 7.35 - 7.45 pH = 0.699 Milk of Magnesia 10.5 Household ammonia 12.0 1M NaOH 14.0 pH of 0.5 M H2SO4 Solution H2SO4(aq) + H2O(l) (aq) H3+O(aq) + HSO4- HSO4-(aq) + H2O(l) H3+O(aq) + SO42-(aq) [H3+O][HSO4-] H2SO4 ; Ka1 = ------------------[H2SO4] [H3+O][SO42-] H2SO4 ; Ka2 = ------------------- ; Ka2 ignored pH of 0.5 M H2SO4 Solution 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 + ABase: BOH + H2O B+ + OH[H3+O][ A-] [B+ ][OH-] Ka = --------------- ; Kb = ----------------[HA] [BOH] Acid Dissociation Constant HCl(aq) + H2O(l) Ka= [H3+O][Cl-] ----------------[HCl] + [H ][Cl-] Ka= ----------------[HCl] H3+O(aq) + Cl-(aq) Base Dissociation Constant NH3 + H2O K = NH4+ + OH- [NH4+][OH-] [NH3] Hydrated Metal Ions as Acids [Fe(H2O)6]3+ (aq) + H2O ( ) [Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq) [Fe(H2 O)5 (OH)2 ][H3 O ] 3 Ka 6.3 10 Fe(H2 O)3 6 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 b. HOCl2 c. HOCl d. HCN ; Ka= 4.0 x 10-4. ; Ka= 1.2 x 10-2. ; Ka= 3.5 x 10-8. ; 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 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 Amount dissociated % 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 0.0270.027 pH = -log [H+] pH = -log(0.027) pH = 1.57 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 [H3O+][Bz-] Ka = [HBz] Weak acid Equilibria HBz Initial conc., M 0.10 Change, DM -x H3O+ 0.00 x Bz- 0.00 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 x2 0.10 = (6.28 x 10-5 )(0.10) H3O+ = 0.0025 M pH = 2.60 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-] Ka = --------------[HCN] = x2 ---------------1.0 - x 1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = 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 x2 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 NH Cl 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) 2) 3) 4) strong acid + strong base = neutral weak acid + strong base = basic 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 • acidic 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 H H H H O C C C C H H H O nonacidic hydrogens H acidic hydrogen butanoic acid O H C 3 electron-attracting oxygen atom C O H C OH acetic acid acidic hydrogen 3 O C H C O 3 - - C OH acetate ion 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 BronstedLowery 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 A Typical Synthetic Detergent Molecule H C H 2CH2C H 2CH2C H 2C H 2C H 2C H 2CH2C H 2CH2C H 2 3CH2C SO3-Na+ Water-soluble part (hydrophilic) Oil-soluble part (hydrophobic) A nonionic detergent H C H 3(C H (C O 2)4CO hydrocarbon chain (hydrophobic) ( 2)2O CH2C H ) 2C H 2C H 2O H 2O alcohol group (hydrophilic) ester link (hydrophilic) ether link ether link (hydrophilic) 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)