Relative Strengths of Acids

1. Relative Strengths of Acids • The relative strength of various acids can be compared using Ka. • The larger the value of Ka, the stronger the acid • The structure of an acid plays an important role in determining the strength of an acid.

2. Relative Strengths of Acids • For an acid with the general formula, HX, the strength of the acid depends on: • the polarity of the H - X bond H X • the strength of the H - X bond • the stability of the conjugate base, X-

3. Relative Strengths of Acids • Using information about these properties, several trends related to the relative strengths of acids can be identified for binary acids and oxyacids. • Binary acid: • an acid with the general formula, HX, which is generally composed of two elements (one of which is hydrogen) • Oxyacid: • an acid that contains oxygen

4. Relative Strengths of Acids Relative Strengths of Binary Acids: • Within a group, the strength of an acid increases moving down the group • HCl is stronger than HF • Within the same period, the strength increases as the electronegativity of the element X increases (i.e. left to right) • HCl is stronger than H2S

5. Relative Strengths of Acids Increasing acid strength For binary acids: Periodic Table Increasing acid strength

6. Relative Strengths of Acids • Relative Strengths of Oxyacids: • For oxyacids with the same central atom, acid strength increases as the number of oxygen atoms attached to the central atom increases. • HClO3 is stronger than HClO

7. Relative Strengths of Acids • Relative Strengths of Oxyacids: • For oxyacids with the same number of O atoms, acid strength increases as the electronegativity of the central atom increases • HClO is stronger than HBrO In general, electronegativity increases toward the top within a group and from the left toward the halogens within a period.

8. Periodic Table Halogens Relative Strengths of Acids For oxyacids with the same # of oxygens: Increasing acid strength Increasing acid strength

9. Relative Strengths of Acids Example: Identify the stronger acid in each of the following pairs: HClO2 vs. HIO2 H2SeO3 vs. H2SeO4 H2SeO3 vs. HBrO3 H2Ovs. HF

10. Buffers • Before measuring the pH of an aqueous solution using a pH meter, chemists must standardize the pH meter. • Adjust the reading of the pH meter to the correct value using standardbuffers. • Buffer: • A solution that resists a change in pH when small amounts of acid or base are added

11. Buffers • Buffers resist changes in pH because they contain both: • An acid • Neutralizes any OH- ions added • A base • Neutralizes any H+ ions added • At the same time, the acidic and basic components of a buffer must not react with each other. • Generally use a weak conjugate acid-base pair

12. Buffers • Examples of buffers: • NaC2H3O2 + HC2H3O2 • NH4Cl + NH3 • Lactic acid + sodium lactate • Two important properties of a buffer: • pH • Buffer capacity

13. Buffers • Buffer capacity • the amount of acid or base a buffer can neutralize before the pH begins to change appreciably • The buffer capacity depends on the amount of acid and base from which the buffer is prepared.

14. Buffers • The pH of a buffer depends on the pKa of the acids and on the relative concentrations of the acid and base in the buffer. pH = pKa + log10 [base] [acid] • pKa = - log Ka Henderson- Hasselbalch Equation

15. Buffers • The Henderson-Hasselbalch equation can be used to: • determine the pH of a particular buffer • determine the ratio of base to acid needed to prepare a buffer with a specific pH. OR

16. Buffers Example: What is the pH of a buffer containing 0.12 M benzoic acid and 0.20 M sodium benzoate? Ka = 6.3 x 10-5.

17. Buffers • First, assume that the ionization of the acid and base in the buffer is negligible. • Use the initial concentrations of the acid and base in the H-H equation. pH = pKa + log [base] [acid]

18. Buffers Example: What base to acid ratio is needed to make a pH 4.95 buffer using benzoic acid and sodium benzoate? Ka = 6.3 x 10-5.

19. Buffers

20. Acid-Base Titrations • An acid-base titration can be used to determine the concentration of an acid or base solution • titration: • a technique for determining the concentration of an unknown solution using astandard solution • a solution with a known concentration

21. Acid-Base Titrations • The equivalence point in a titration can be determined using either a pH indicator (ex. phenolphthalein) or a pH meter. • Equivalence point: • the point in the titration where stoichiometrically equivalent amounts of base have been added to the acid (or vice versa) • the base added has completely reacted with all available protons (H+)

22. Equivalence point Region of high buffer capacity pKa Acid-Base Titrations • The general shape for a titration curve for a monoprotic acid: pH mL base added

23. Acid-Base Titrations • General Shape for a Diprotic Acid Titration Curve:

24. Acid-Base Titrations • General Shape for a Triprotic Acid Titration Curve:

25. Titration of Polyprotic Acids Example: The titration curve on the next slide was obtained by titrating 0.250 g of a polyprotic acid with 0.100 M NaOH. Answer the following questions. • How many equivalence points are there? • Is the acid monoprotic, diprotic, or triprotic? • What are the values for each pKa? • What volume of NaOH was needed to neutralize the acid? • What is the molar mass of the unknown acid?

26. Titration of Polyprotic acids

28. Lewis Acids and Bases • In order for a substance to be a Bronsted-Lowry base (i.e. a proton acceptor), it must have an unshared pair of electrons to bind the proton. • Similarly, a Bronsted-Lowry acid (i.e. a H+ ion) always gains a pair of electrons during an acid-base reaction.

29. Lewis Acids and Bases • Lewis noticed this trend and proposed a new definition of acids and bases: • Lewis Acid: • An electron pair acceptor • Lewis Base: • An electron pair donor.

30. Lewis Acids and Bases • Examples of Lewis Acids: • H+ • BF3 • Fe3+ • CO2 • Examples of Lewis Bases: • OH- • CH3NH2

31. Lewis Acids and Bases Example: Identify the Lewis acid and the Lewis base in the following reactions.