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
For binary acids:
Periodic
Table
Increasing acid strength
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. Relative Strengths of Acids
For oxyacids with the same # of oxygens:
Halogens
Periodic Table
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
H2O vs. 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 standard buffers.
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. Henderson- Hasselbalch Equation
Buffers
The pH of a buffer depends on the pKa of the acid 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 OR 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
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.

18. 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 a standard solution
a solution with a known concentration

19. 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+)

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

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

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

23. Titration of Polyprotic Acids
Example: The titration curve on the next slide was obtained by titrating g of a polyprotic acid with 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?

24. Titration of Polyprotic acids

26. 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.

27. 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

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

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