Properties, Naming, and Definitions of Acids and Bases

1. Chapter 9Acids and Bases CHE 101 Sleevi

2. Properties of Acids • Taste sour • Produce H+ in water (H3O+) • React with metals to produce H2 (g) • pH <7 • Dissolve ionic compounds insoluble in water (e.g., CaCO3)

3. Properties of Bases • Solutions containing bases feel soapy or slippery • pH >7 • React readily with acids • React with fats and oils and convert them to smaller, soluble molecules • Used as part of cleaning solutions • Often contain hydroxide ion

4. Naming Acids • Binary acids: • Use prefix hydro- and derivative of element name with suffix -ic • HCl = hydrochloric • Acids containing polyatomic ions: • Ion ends in –ate, acid name ends in –ic • Ion ends in –ite, acid name ends in –ous H2SO4 = sulfuric acid H2SO3 = sulfurous acid

5. Naming Acids • H2Se • HClO3 • H3PO4 • HNO2 • HBr • HC2H3O2

6. Definition of Acids & BasesArrhenius Theory • Acid: • A substance that forms H+ in solution HCl(aq) H+(aq) + Cl-(aq) • Base: • A substance that forms OH- in solution KOH(aq) K+(aq) + OH-(aq)

7. Arrhenius Acids Write dissociation equations for the following acids to illustrate their behavior as Arrhenius Acids HBr  HCN  HClO4 

8. Arrhenius Bases Write dissociation equations for bases the Arrhenius bases. ID those that are not Arrhenius bases LiOH C2H5NH2 Ca(OH)2 NH3

9. Definition of Acids & BasesBrǾnsted-Lowry Theory • Acid (H-A) • donates a proton to another substance • Base (B:) • accepts a proton from another substance • must contain a lone pair of electrons that can be used to form new bond to the proton • NH3, H2O, OH-, Cl-

10. Proton TransferThe Reaction of a Brønsted–Lowry Acid witha Brønsted–Lowry Base This e− pair forms a new bond to H+ This e− pair stays on A gain of H+ H A + A− + H B+ B acid base loss of H+

11. Definition of Acids & BasesBrǾnsted-Lowry Theory • When a species gains a proton (H+), it gains a +1 charge • When a species loses a proton (H+), it effectively gains a -1 charge

12. Definition of Acids & BasesBrǾnsted-Lowry Theory • Conjugate Acid: • species formed when base gains a proton • Conjugate Base: • species formed when acid loses a proton H—A(aq) + B:(l) H—B+ + A- acid base conjugate conjugate acid base H—Cl(aq) + H2O(l) H3O + + Cl-

13. Conjugate Acids and Bases H-Br + H2O Br- + H3O+ • Conjugate acid-base pairs HBr and Br- H2O and H3O+ Equation must be mass and charge balanced!

14. Identifying Acids, Bases and Conjugates HNO3(aq) + H2O (l) H3O+ (aq) + NO3- (aq) NH3(aq) + H2O (l) OH- (aq) + NH4+ (aq) SO32-(aq) + H2O (l) HSO3- (aq) + OH- (aq)

15. Identifying Acids, Bases and Conjugates Write balanced equations: HClO3(aq) + NH3 (aq) H2SO4(aq) + H2O (l)

16. Types of Acids • monoprotic acid • gives up only one proton per molecule when dissolved (HCl) • diprotic acid • gives up two protons per molecule when dissolved (H2SO4) • triprotic acid • gives up three protons per molecule when dissolved (H3PO4)

17. Acid Behavior HCl + NaOH  NaCl + H2O HCl is monoprotic acid 1 mole of HCl reacts with 1 mole of NaOH

18. Acid Behavior H2SO4 + 2NaOH  Na2SO4 + 2H2O H2SO4 is a diprotic acid 1 mole of H2SO4 reacts with 2 moles of NaOH

19. Amphoteric Compounds • Compounds that can be either an acid or base • Contain both H and lone pair • Examples: • H2O • NH3 • HSO3-

20. The pH Scale

21. Analyzing Acids and Bases • Determining pH range of an acid or base: • pH meter • acid-base indicator • universal indicator • red, orange, green, blue, purple • litmus paper • blue  pink, pink  blue • phenolphthalein • (acid/neutral  colorless; base  pink)

22. pH pH = - log [H+] [H+] = 1 x 10-pH • Each unit on pH scale is a factor of 10 different from the next lower or higher number

23. Examples

24. Acid and Base Strength • Strong acid, strong base • fully dissociated in water NaOH HCl

25. Acid and Base Strength • Weak acid, weak base • partially dissociated in water • HC2H3O2, NH4+, H2CO3, citric acid, NH3

26. Acid and Base Strength • Strong acid forms weak conjugate base • Strong base forms weak conjugate acid • Concentration of acids and bases measured in molarity (moles/L) • Dilute vs weak – a strong acid may be dilute, a weak acid may be concentrated – pH does not differentiate

27. Acid and Base StrengthPredicting the Direction of Equilibrium • When the stronger acid and base are the reactants • on the left side, the reaction readily occurs and • the reaction proceeds to the right. H A + B: A- + H B+ stronger acid stronger base weaker base weaker acid • A larger forward arrow means that products are • favored.

28. Acid and Base StrengthPredicting the Direction of Equilibrium • If an acid–base reaction would form the stronger • acid and base, equilibrium favors the reactants • and little product forms. H A + B: A- + H B+ weaker acid weaker base stronger base stronger acid • A larger reverse arrow means that reactants are • favored.

29. Predict Direction of Equilibrium −CN(aq) HCN(g) −OH(aq) + H2O(l) + See Table 9.1 - Relative Strength of Acids and Their Conjugate Bases

30. Table 9.1

31. Equilibrium and Acid Dissociation Constants For the reaction where an acid (HA) dissolves in water, HA(g) + H2O(l) H3O+(aq) + A:-(aq) the following equilibrium constant can be written: [H3O+][A:-] K = [HA][H2O]

32. Equilibrium and Acid Dissociation Constants • Multiplying both sides by [H2O] forms a new constant, • called the acid dissociation constant, Ka. [H3O+][ A:- ] Ka = K[H2O] = [HA] acid dissociation constant • The stronger the acid, the larger the value of Ka. • Equilibrium, though, favors formation of the weaker acid—that is, the acid with the smaller value of Ka.

33. Equilibrium and Acid Dissociation Constants

34. − H O Dissociation of Water Water can behave as both a Brønsted–Lowry acid and a Brønsted–Lowry base. Thus, two water molecules can react together in an acid–base reaction: loss of H+ + H H H H O H H O H + + O acid base conjugate acid conjugate base gain of H+

35. Dissociation of Water • From the reaction of two water molecules, the • following equilibrium constant expression can be • written: [H3O+][−OH] K = [H2O]2 • Multiplying both sides by [H2O]2 yields Kw, the • ion-product constant for water. Kw = [H3O+][−OH] ion-product constant

36. Dissociation of Water • Experimentally it can be shown that [H3O+] = [−OH] = 1.0 x 10−7 M at 25 oC Kw = [H3O+] [−OH] Kw = (1.0 x 10−7) x (1.0 x 10−7) Kw = 1.0 x 10−14 • Kw is a constant, 1.0 x 10−14, for all aqueous • solutions at 25 oC.

37. Dissociation of Water To calculate [H3O+] when [−OH] is known: To calculate [−OH] when [H3O+] is known: Kw = [H3O+][−OH] Kw = [H3O+][−OH] Kw Kw [−OH] [H3O+] = = [−OH] [H3O+] 1 x 10−14 1 x 10−14 [−OH] [H3O+] = = [H3O+] [−OH]

38. Dissociation of Water If the [H3O+] in a cup of coffee is 1.0 x 10−5 M, then the [−OH] can be calculated as follows: Kw 1 x 10−14 [−OH] = = 1.0 x 10−9 M = 1 x 10−5 [H3O+] In this cup of coffee, therefore, [H3O+] > [–OH], and the solution is acidic overall.

39. Dissociation of Water

40. Chemical Equations and Solutions • Molecular equation • Each substance represented by its formula HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l) 2 Al (s) + 3 Cu(NO3)2 (aq)  2 Al(NO3)3(aq) + 3 Cu (s)

41. Chemical Equations and Solutions • Total Ionic Equation • All soluble ionic substances represented by the ions they form in solution • Solids, liquids, gases and aqueous solutions of molecular compounds do not dissociate HCl(aq) + NaOH(aq)  NaCl(aq) + H2O (l) =>H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l)

42. Spectator Ions • Ions that appear on both sides of the chemical equation (not changed in the chemical reaction)

43. Net Ionic Equation • Contains only unionized or insoluble materials and ions that undergo changes in the reaction • All spectator ions are eliminated H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l) H+(aq) + OH- (aq) H2O (l)

44. Net Ionic Equations 2 Al (s) + 3 Cu(NO3)2 (aq)  2 Al(NO3)3(aq) + 3 Cu (s) 2 Al (s) +3Cu2+ (aq) + 6NO31-(aq)  2 Al3+ (aq) + 6NO31-(aq) + 3 Cu (s) 2 Al (s) +3 Cu2+ (aq)  2 Al3+ (aq)+ 3 Cu (s) Net ionic equations must be mass and charge balanced

45. How to Write a Net Ionic Equation • Write balanced molecular equation • Write total ionic equation • Aqueous ionic compounds written as individual ions • Multiply by subscript and coefficient to balance mass and charge • Compounds that appear as solids, liquids, gases or aqueous solutions of molecular compounds are written in molecular form. • Write net ionic equation • Eliminate spectator ions • Include all solids, liquids, gases and non-spectator ions • Verify mass and charge balance

46. Analyzing Acids and Bases • Determine concentration of acid or base using neutralization reactions and titration • Equivalence point – the point at which the acid has exactly neutralized the base (neither is in excess)

47. Performing a Titration • Slowly add base from a buret to an acid in a receiving flask • Use phenolphthalein to indicate when endpoint is reached • Measure volumetric amount of base of known concentration • Calculate concentration of acid using solution stoichiometry

48. Performing a Titration

49. Calculating Unknown Concentration of Acid Solution • Write balanced equation for neutralization reaction • Titrate acid solution with known concentration of base solution (to phenolphthalein endpoint) • Determine accurate volume of base used in neutralization • Calculate concentration of acid solution using solution stoichiometry

50. Titration Videos https://www.youtube.com/watch?v=sFpFCPTDv2w https://www.youtube.com/watch?v=2z4mlE6MK0U