Unit 16: Acids and Bases, and pH acidic solution basic solution Arrhenius model

1. Unit 16: Acids and Bases, and pH acidic solution basic solution Arrhenius model Brønsted-Lowry model conjugate acid conjugate base conjugate acid-base pair amphoteric Lewis model strong acid weak acid acid ionization constant strong base weak base base ionization constant strong acid weak acid acid ionization constant strong base weak base base ionization constant acid-base indicator end point salt hydrolysis buffer buffer capacity

2. Acids and Bases: An Introduction Aqueous Solutions = acid + base hydrogen ions (H+) + hydroxide ions (OH-) hydronium ion (H3O+) is a hydrated hydrogen ion. The acidic symbols (H+) and (H3O+) can be used interchangeably in chemical equations.

3. All aqueous solutions contain hydrogen ions (H+) and hydroxide ions (OH-). An acidic solution contains more H+ ions than OH- ions, whereas a basic solution contains more OH- ions than H+ ions A neutral solution contains equal concentrations of H+ ions and OH- ions

4. Macroscopic Properties of Acids and Bases Taste and feel: Acids taste sour (lemon juice, vinegar)Bases taste bitterBases are slippery (soap)

5. Acids react with bases – the reaction of acids and bases are central to the chemistry of living systems, the environment, and many important industrial processes

7. Litmus test:Indicators- • change colors in the presence of an acid or a base • Litmus: • base = blue • acid = red

8. Submicroscopic Behavior of Acids:Hydronium ion formationHCl + H2O H3O+ + Cl-HNO3 + H2O H3O+ + NO3-

9. The Arrhenius Model • Acids – produce hydrogen ions in aqueous solution HCl H+ + Cl- • Bases – produce hydroxide ions in aqueous solution NaOH Na+ + OH-

10. The Bronsted-Lowry Model • Arrhenius concept is limited because it only allows for one kind of base. It expresses the substance. • Acid – proton donor • Base – proton acceptor

11. General reaction for a Bronsted-Lowry acid dissolving in water: • Conjugate acid-base pair: two substances related to each other by the donating and accepting of a single proton.

12. Brønsted-Lowry Model • HCl+ H2O H3O+ + Cl- • acid conjugate base • HCl + H2OH3O+ + Cl- • base conjugate acid • HNO3 + H2O H3O+ + NO3- • acid conjugate base • HNO3 + H2OH3O++ NO3- • base conjugate acid

13. Identify the acid, base, conjugate acid, and conjugate base in the following: • HCO3-1 + H2O CO3-2 + H3O +1

14. Water as an Acid and a Base • Amphoteric substance – can behave either as an acid or as a base • Ionization of water:

15. Ionization Example Boric Acid, H3BO4 H3BO4(aq) + H2O(l) ↔ H2BO4-(aq) + H3O+(l) H2BO4-(aq) + H2O(l) ↔ HBO42-(aq) + H3O+(l) HBO42-(aq)) + H2O(l) ↔ BO43-(aq) + H3O+(l)

17. Monoprotic and Polyprotic Acids • Hydrogen atoms that are bonded to electronegative elements are ionizable • Monoprotic acid can donate only one hydrogen Example- HCl • Polyprotic acids can donate more than one hydrogen atom • diprotic acid has two ionizable hydrogens Example: sulfuric acid • triprotic acids has three ionizable hydrogens Example: Boric acid

18. Strengths of Acids and Bases • Strong acids and bases ionize completely • Weak acids and bases ionize only partially

20. A strong acid contains a relatively weak conjugate base, one that has a low attraction for protons • A weak acid contains a relatively strong conjugate base

21. Hydrogen and Hydroxide Ions and pH • Acidity or Basicity of a substance is related to the concentration of hydrogen and hydroxide ions in that substance

22. The product of [H+] and [OH-] is always constant • Kw = ion product constant for water • No matter what the solution contains, Kw will always equal 1.0 x 10-14

24. Calculate the [H+] or [OH−] and state whether the solution is neutral, acidic, or basic • 1.0 x 10-5 M OH- • 1.0 x 10-7 M OH- • 10.0 M H+

25. Determining the Acidity of a Solution • The pH Scale: 0 – 14 • <7 = acidic • 7 = neutral • >7 = basic

27. A mathematical scale in which the concentration of H+ ions in a solution is expressed as a number from 0 – 14 • pH = −log [H+]

28. What is the pH of solutions having the following ion concentrations? • [H+] = 1.0 x 10-2 M • [H+] = 3.0 x 10-6 M

29. Because the pH scale is a log scale based on 10, the pH changes by 1 for every power of 10 change in the [H+]

30. Log scales similar to the pH scale are used for representing other quantities: pOH = −log [OH−]

31. What is the pOH of a solution having the following ion concentration? • [OH-] = 1.0 x 10-6 M • [OH-] = 6.5 x 10-4 M

32. pH + pOH = 14 • What is the pOH of a solution whose pH is 5? • What is the pH of a solution whose [OH-] = 4.0 x 10-3 M

33. Calculate the pH and pOH of the following solutions: • [H+] = 0.000033 M • [OH-] = 0.0095 M

34. It is also possible to find the [H+] or [OH-] from the pH or pOH by undoing the log operation • [H+]= 10-pH • [OH-] = 10-pOH

35. The pH of a human blood sample was measured to be 7.41. What is the [H+] and [OH-]in this blood?

36. Neautralization Reactions • Acid + Base salt + water • HCl + NaOHNaCl + H2O • Ionic: • Net ionic:

37. Titration • Method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration. • If unknown is acid, known must be base

38. Titration procedure • Measured volume of acid or base of unknown concentration is placed in a flask and initial pH is recorded • Buret is filled with the solution of known concentration (standard solution/titrant) • Standard solution added slowly until neutral pH is reached (equivalence point) [H+] = [OH-]

40. End point of a titration can be measured using a pH meter or an indicator • Indicator changes color at different pH values

41. Titration curve (pH curve) – plot of pH vs volume of titrant added

42. A volume of 18.28 mL of a standard soluiton of 0.1000M NaOH was required to neutralize 25.00 mL of a solution of nitric acid. What is the concentration of the nitric acid?

43. Buffered Solutions • Solutions that resist changes in pH when acids or bases are added • Weak acid + conjugate base • HF + NaF