George Mason University General Chemistry 212 Chapter 18 Acid-Base Equilibria Acknowledgements

1. George Mason University General Chemistry 212 Chapter 18 Acid-Base Equilibria Acknowledgements Course Text: Chemistry: the Molecular Nature of Matter and Change, 6thed, 2011, Martin S. Silberberg, McGraw-Hill The Chemistry 211/212 General Chemistry courses taught at George Mason are intended for those students enrolled in a science /engineering oriented curricula, with particular emphasis on chemistry, biochemistry, and biology The material on these slides is taken primarily from the course text but the instructor has modified, condensed, or otherwise reorganized selected material.Additional material from other sources may also be included. Interpretation of course material to clarify concepts and solutions to problems is the sole responsibility of this instructor.

2. Chapter 18 – Acid-Base Equilibria • Acids & Bases in Water • Arrhenius Acid-Base Definition • Bronsted-Lowry Acid-Base Definition • Lewis Acid-Base Definition • Acid Dissociation Constant • Relative Strengths of Acids & Bases • Autoionization of Water and the pH Scale • Autoionization and Kw • The pH Scale • Proton Transfer and the Bronsted-Lowry Acid-Base Definition • The Conjugate Acid-Base Pair • Met Direction of Acid-Base Reactions

3. Chapter 18 – Acid-Base Equilibria • Solving Problems Involving Weak-Acid Equilibria • Finding Ka Given Concentrations • Finding Concentrations Given Ka • Extent of Acid Dissociation • Polyprotic Acids • Weak Bases and Their Relation to Weak Acids • Ammonia & The Amines • Anions of Weak Acids • The Relation Between Ka & Kb • Molecular Properties and Acid Strength • Nonmetal Hydrides • Oxoacids • Acidity of Hydrated Metal Ions

4. Chapter 18 – Acid-Base Equilibria • Acid-Base Properties of Salt Solutions • Salts That Yield Neutral Solutions • Salts That Yield Acidic Solutions • Salts That Yield Basic Solutions • Salts of Weakly Acidic Cations and Weakly Basic Anions • Salts of Amphoteric Anions • Generalizing the Bronsted-Lowry Concept: The Leveling Effect • Electron-Pair Donation and the Lewis Acid-Base Definition • Molecules as Lewis Acids • Metal Cations as Lewis Acids • Overview of Acid-Base Definitions

5. Acid & Bases • Antoine Lavoisier was one of the first chemists to try to explain what makes a substance acidic • In 1777, he proposed that oxygen was an essential element in acids • The actual cause of acidity and basicity was ultimately explained in terms of the effect these compounds have in water by Svante Arrhenius in 1884

6. Acids & Bases • Several concepts of acid-base theory: • The Arrhenius concept • The Bronsted-Lowry concept • The Lewis concept

7. Acids & Bases • According to the Arrheniusconcept of acids and bases • Anacidisa substance that, when dissolved in water, increases the concentration of Hydrogen ion, H+(aq) • The H+(aq) ion, called the Hydrogen ion is actually chemically bonded to water, that is, H3O+, called the Hydronium ion • Abaseisa substance that when dissolved in water increases the concentration of the Hydroxide ion, OH-(aq)

8. Acids & Bases More About the Hydronium Ion The Hydronium Ion (H3O+) actually exists as a hydrogen bonded H9O4+ cluster. The formation of Hydronium ions is a complex process in aqueous solution The Hydronium ion has trigonal pyramidal geometry because all three H-O-H dihedral bond angles are identical

9. Acids & Bases • An Arrhenius acidis any substance which increasesthe concentration of Hydrogen ions (H+) in solution. Acids generally have a sour taste • HBr(aq)  H+(aq) + Br-(aq) • AnArrhenius baseis any substance which increases the concentration of Hydroxide ions (OH-) in solution. Bases generally have a bitter taste • KOH(s) K+(aq) + OH-(aq) • Hydrogen and Hydroxide ions are important in aqueous solutions because - Water reacts to form both ions • H2O(l)  H+(aq) + OH-(aq) • HBr and KOH are examples of a strong acidand a strong base; strong acids & bases dissociate completely

10. Acids & Bases • Monoprotic acids are those acids that are able to donate one proton per molecule during the process of dissociation (sometimes called ionization) as shown below (symbolized by HA): HA(aq) + H2O(l) ⇌ H3O+(aq) + A−(aq) Common examples of monoprotic acids in mineral acidsinclude Hydrochloric Acid (HCl) and Nitric Acid (HNO3) • Polyprotic acids are able to donate more than one proton per acid molecule. • Diprotic acids have two potential protons to donate H2A(aq) + H2O(l) ⇌ H3O+(aq) + HA−(aq) • Triprotic acids have three potential protons to donate H3A(aq) + H2O(l) ⇌ H3O+(aq) + H2A−(aq)

11. Acids & Bases • Oxoacids • Oxoacids with one oxygen have the structure: HOY HOClHOBr HOI • Oxoacids with multiple oxygen have the structure: (OH)mYOn HOClO3 (HClO4) HOClO2 (HOCLO3) HOClO (HCLO2) HOCL (HClO) The more oxygen atoms, the greater the acidity; thus HOCLO3 is a stronger acid than HOCl Acidity refers to the relative acid strength, i.e. the degree to which the acid dissociates to form H+ (H3O+) & OH-

12. Bronsted-Lowry Acids & Bases • Brønsted-Lowry Concept of Acids and Bases • In the Brønsted-Lowry concept: • An Acid is a species that donates protons • A Base is a species that accepts protons • Acids and bases can be ions as well as molecular substances • Acid-base reactions are not restricted to aqueous solution • Some species can act as either acids or bases (Amphoteric species) depending on what the other reactant is

13. Bronsted-Lowry Acids & Bases • Amphoteric Species • A species that can act as either an acid or base is: Amphoteric • Water is an important amphoteric species in the acid-base properties of aqueous solutions • Water can react as an acid by donating a proton to a base • Water can also react as a base by accepting a proton from an acid H+ H+

14. Bronsted-Lowry Acids & Bases • Brønsted-Lowry Concept of Acids and Bases • Proton transfer as the essential feature of a • Brønsted-Lowry acid-base reaction Lone pair binds H+ Acid H+ donor Base H+ acceptor Lone pair binds H+ Acid H+ donor Base H+ acceptor

15. Acids & Bases • Bronsted-Lowery Concept – Conjugate Pairs • In Bronsted theory, acid and base reactants form: • Conjugate Pairs • The acid (HA) donates a proton to water leaving the conjugate Base (A-) • The base (H2O) accepts a proton to form the conjugate acid (H3O+) • HA(aq) + H2O(l) ⇌ A−(aq) + H3O+(aq) • Acid Base Conjugate Conjugate Base Acid

16. Bronsted-Lowry Acids & Bases Bronsted-Lowery Acid-Base Conjugate Pairs HNO2(aq) + H2O(l)  H3O+aq) + NO2-(aq) acid base acid base BronstedAcid BronstedBase H3O+ is conjugate acid of base H2O NO2- is conjugate base of acid HNO2 • Conjugate acid-base pairs are shown connected • Every acid has a conjugate base • Every base has a conjugate acid • The conjugate base has one fewer H and one more “minus” charge than the acid • The conjugate acid has one more H and one fewer minus charge than the base

17. Bronsted-Lowry Acids & Bases • Bronsted-Lowery Acid-Base Conjugate Pairs NH3(aq) + H2O(l)  NH4+aq) + OH-(aq) base acid acid base BronstedBase BronstedAcid NH4+ is conjugate acid of base NH3 OH- is conjugate base of acid H2O Accepts Proton Donates Proton

18. Bronsted-Lowry Acids & Bases Bronsted-Lowery Acid-Base Conjugate Pairs acid base acid base H2S + NH3 NH4++HS- BronstedAcid BronstedBase NH4+ is conjugate acid of base NH3 HS- is conjugate Base of acid H2S Donates Proton Accepts Proton

19. Bronsted-Lowry Acids & Bases • Bronsted_Lowry Concept – The Leveling Effect • All Bronsted_Lowry acids yield H3O+ ions (Cation) • All Bronsted_Lowry bases yield OH- ions (Anion) • All strong acids are equally strong because all of them form the strongest acid possible - H3O+ • Similarly, all strong bases are equally strong because they form the strongest base possible - OH- • Strong acids and bases dissociate completely yielding H3O+ and OH- • Any acid stronger than H3O+ simply donates a proton to H2O • Any base stronger than OH- simply accepts proton from H2O • Water exerts a leveling effect on any strong acid or base by reacting with it to form water ionization products

20. Practice Problem • Identify the conjugate acid-base pairs in the following: Base Conjugate Acid Acid Conjugate base

21. Lewis Acids & Bases • Lewis Acids & Bases • Some acid base reactions don’t fit the Bronsted-Lowry or Arrhenius classifications • The Lewis acid-base concepts expands the acid class • Such reactions involve a “sharing” of electron pairs between atoms or ions • Lewis Acid – An electron deficient species (Electrophile) that accepts an electron pair • Lewis Base – An electron rich species (Nucleophile) that donates an electron pair

22. Lewis Acids & Bases F H N + B H F H F • Lewis Acids & Bases • The product of any Lewis acid-base reaction is called an “Adduct”, a single species that contains a new covalent bond (shared electron pair) An acid is an electron-pair acceptor A base is an electron-pair donor Adduct

23. Lewis Acids & Bases • Species that do not contain Hydrogen in their formulas, such as CO2 and Cu++ function as Lewis acids by accepting an electron pair • The donated proton of a Bronsted-Lowry acid acts as a Lewis acid by accepting an electron pair donated by the base: • The Lewis acids are Yellow and the Lewis bases are Blue in following reaction BF3 + :NH3  BF3NH3 Fe3+ + 6 H2O  Fe(H2O)63+ HF + H2OF-+ H3O+ + -

24. Lewis Acids & Bases • Solubility Effects • A Lewis acid-base reaction between nonpolar Diethyl Ether and normally insoluble Aluminum Chloride - + • The solubility of Aluminum Chloride in Dimethyl Ether results from the Oxygen acting as a base by donating an electron pair to the Aluminum acting as an acid forming a water-soluble polar covalent bond

25. Practice Problem The following shows ball-and-stick models of the reactants in a Lewis acid-base reaction. Write the complete equation for the reaction, including the product Identify each reactant as a Lewis acid or Lewis base Donates e- pair Accepts e- pair AlCl3 + :NH3 AlCl3:NH3 Lewis acid Lewis base Adduct

26. Acids & Bases (Summary) • Arrhenius acid concept • acid= proton (H+) producer; base = OH- producer • Formation of Water (H2O) from H+ & OH- • Bronsted-Lowry concept • Stronger Acid (HA) transfers proton to a stronger base (H2O)to form weaker acid H3O+) and weaker base (A-) • Stronger base (NH3) accepts proton from stronger acid (H2O) • Lewis Concept • Donation and acceptance of an electron pair to form a covalent bond in an adduct

27. Acids & Bases • Self Ionization of Water • Pure water undergoes auto-ionization to produce Hydronium and Hydroxide ions 2 H2O(l)  H3O+(aq) + OH-(aq) • The extent of this process is described by an auto-ionization (ion-product) constant, Kw • The concentrations of Hydronium and Hydroxide ions in any aqueous solution must obey the auto-ionization equilibrium • Classification of solutions using Kw: Acidic: [H3O+] > [OH-] Basic: [H3O+] < [OH-] Neutral: [H3O+] = [OH-]

28. Acids & Bases • Because of the auto-ionization constant of water, the amounts of Hydronium and Hydroxide ions are always related • For example, pure water at 25 oC, • The Definition of pH • The amount of Hydronium ion (H3O+) in solution is often described by the pH of the solution • The Definition of pOH • Hydroxide (OH-) can be expressed as pOH

29. Acids & Bases • pH scale ranges from 0 to 14 • The pH of pure water = -log(1.00 X 10-7) = 7.00 pH = 0.00 - 6.999+ = acidic [H3O+] > [OH-] pH = 7.00 = neutral [H3O+] = [OH-] pH = 7.00+ - 14.00 = basic [H3O+] < [OH-] • Acidic solutions have a lower pH (higher [H3O+] and a higher pOH (lower [OH-] than basic solutions

30. Acids & Bases • Relations among pH, pOH, and pKw

31. Equilibrium & Acid-Base Conjugates • Equilibrium Constants (K) for acids & bases can also be expressed as pK • Reaction of an Acid (HA) with water as a base • Reaction of a Base (A-) with water as an acid

32. Equilibrium & Acid-Base Conjugates • Relationship between: Ka and HA Kb and A- • Setup two dissociation reactions: • The sum of the two dissociation reactions is the autoionization of water • The overall equilibrium constant is the product of the individual equilibrium constants (autoionization)

33. Equilibrium & Acid-Base Conjugates • From this relationship, Ka of the acid in a conjugate pair can be computed from Kb and vice versa • Reference tables typically have Ka & Kb values for molecular species, but not for ionssuch as F- or CH3NH3+

34. Equilibrium & Acid-Base Conjugates • Acid strength [H3O+] increases with increasing value of Ka • Base strength [OH-] increases with increasing value of kb • A lowpK corresponds to a high value of K • A reaction that reaches equilibrium with mostly products present (proceeds to far right) has a “low pK” (high K)

35. Practice Problem • Ex. Find Kb value for F-

36. Acids & Bases The Relationship Between Ka and pKa Acid Name (Formula) Ka at 25oC pKa 1.0x10-2 Hydrogen Sulfate ion (HSO4-) 1.99 Nitrous Acid (HNO2) 7.1x10-4 3.15 Acid Strength Acetic Acid (CH3COOH) 1.8x10-5 4.74 Hypobromous Acid (HBrO) 2.3x10-9 8.64 Phenol (C6H5OH) 1.0x10-10 10.00

37. Acids & Bases The pH Scale

38. Acids & Bases • Strong Acids & Bases • Strong acids and bases dissociate completely in aqueous solutions Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4 Strong bases: all Group I & II Hydroxides NaOH, KOH, Ca(OH)2, Mg(OH)2 • The term strong has nothing to do with the concentration of the acid or base, but rather the extent of dissociation • In the case of HCl, the dissociation lies very far to the right because the conjugate base (Cl-) is extremely weak, much weaker than water HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) • H2O is a much stronger base than Cl- and “out competes it” for available protons

39. Acids & Bases Activity Series Strengths of acids and bases Strength determined by K, not concentration

40. Practice Problem A strong monoprotic acid is dissolved in a beaker of water. Which of the following pictures best represents the acid solution. Ans: B A solution of a strong acid contains equal amounts of H3O+ and the acid anion, such as a “Cl-” The monoprotic acid is completely dissociated A B C

41. Hydrohalic Acids – Weak vs Strong • The reason Hydrofluoric acid is weak relative to the other Hydrohallic Acids (I, Br, Cl) is quite complicated • The material being somewhat beyond the scope of the discussions we have in the chem 211 & chem 212 courses • It starts out with the short length of the H - F bond and the very high electronegativity of the Fluoride ion • When you take into account the Enthalpy, Entropy, and Free Energy properties (Chapter 19), the numbers clearly show a distinct difference between HF and the other Hydrohalic acids

42. Hydrohalic Acids – Weak vs Strong • These properties can be summarized as follows: • Very strong Hydrogen Bonding between the un-ionized Hydrogen Fluoride (HF) molecules and water molecules • This energy required to break the H-F bond, is much greater then for the bonds of the other Hydrohalic Acid bonds • A large decrease in Entropy when the Hydrogen Fluoride molecules react with water

43. Hydrohalic Acids – Weak vs Strong • This is particularly noticeable with hydrogen fluoride because the attractiveness of the very small fluoride ions produced imposes a lot of order on the surrounding water molecules, and also on nearby hydroxonium ions • The effect falls as the halide ions get bigger

44. Hydrohalic Acids – Weak vs Strong • Note the values below. Hydrofluoric acid clearly stands out • The Ka values clearly suggest a weak acid H TS G Ka (kJ mol-1) (kJ mol-1) (kJ mol-1) (mol dm-3) HF -13 -29 +16 1.6 x 10-3 HCl -59 -13 -46 1.2 x 10+8 HBr -63 -4 -59 2.2 x 10+10 HI -57 +4 -61 5.0 x 10+10

45. Practice Problem What are the concentrations of Hydronium Ions (H3O+) and Hydroxide (OH-) ions in 1.2 M HBr? HBr + H2O  H3O+ + Br- HBr is a strong acid and is completely dissociated Autoionization of Water

46. Practice Problem What are the concentrations of Hydronium Ions (H3O+) and Hydroxide (OH-) ions in 0.085 M Ca(OH)2? Ca(OH)2  Ca+2 + 2OH- Ca(OH)2 is a strong base and is completely dissociated Autoionization of Water

47. Practice Problem What is the pH of the following solution:? [OH-] = 5.6 x 10-10 M

48. Practice Problem What is the pOH of the following solution:? [H3O+] = 9.3 x 10-4 M

49. Practice Problem What is the pH of a solution containing 0.00256 M HCl?

50. Practice Problem What is the pH of a solution containing 0.00813 M Ca(OH)2?