1. Principles of Reactivity: Chemistry of Acids and Bases
Chapter 16

2. Learning Objectives Students understand
The similarities and differences between Brønsted-Lowry acid-base and Lewis acid-base theories.
The influence of structure and bonding on acid-base properties.

3. Learning Objectives Students will be able to
Use the Brønsted-Lowry and Lewis theories of acids and bases
Apply the principles of chemical equilibrium to acids and bases in aqueous solution
Understand how Brønsted-Lowry theory is used to predict the outcome of reactions of acids and bases
Calculate pH
Given Ka and Kb, calculate concentration

4. 16.1 The Brønsted-Lowry Concept of Acids and Bases
Acids are proton (H+) donors
Bases are proton acceptors
Acids capable of donating one proton and so are called monoprotic acids.
Other acids, called polyprotic acids, are capable of donating two or more protons.
Polyprotic bases can accept more than one proton.

5. The Brønsted-Lowry Concept of Acids and Bases
Some molecules and ions can behave either as Brønsted acids or bases and are called amphiprotic.
The reaction of a Brønsted acid and base produces a new acid and base.
Conjugate acid-base pairs differ from each other by the presence of one hydrogen ion

6. Practice Problem
Write a balanced equation for the reaction that occurs when H3PO4, phosphoric acid, donates a proton to water to form the dihydrogen phosphate ion. Is the dihydrogen phosphate ion an acid, a base, or amphiprotic?
Write a balanced equation for the reaction that occurs when the cyanide ion, CN-, accepts a proton from water to form HCN. Is CN- a Bronsted acid or base?

7. Practice Problem
In the following reaction, identify the acid on the left and its conjugate base on with right. Similarly, identify the base on the left and its conjugate acid on the right.
HNO3 + NH3  NH4+ + NO3-

8. 16.2 Water and the pH Scale
Two water molecules interact with each other to produce a hydronium ion and a hydroxide ion.
This property is called autoionization. Water will still conduct a small amount of electricity since it contains low concentration of hydronium and hydroxide.

9. Water Ionization Constant
Kw = [H3O+][OH-] = 1.0 x at 25 oC
In pure water the two ion concentrations are equal and the water is said to be neutral.
Adding acid increases the hydronium concentration, adding base increases the hydroxide. Both additions would disturb the equilibrium.

10. Practice Problem
A solution of the strong acid HCl has [HCl] = 4.7 x 10-3 M. What are the concentrations of H3O+ and OH- in this solution? (remember that HCl is 100% ionized in water)

11. pH Scale
pH = -log[H3O+]
pOH = -log[OH-]
pKw = = pH + pOH

12. Indicators
Approximate pH of solutions can be determined using an acid-base indicator.
Substances that change color in a known pH range are Bronsted acids or bases for which the acid and its conjugate base have different colors.
Modern pH meters are preferable!

13. Practice Problem What is the pH of a 0.0012 M NaOH solution?
The pH of a diet soda is 4.32 at 25oC. What are the hydronium and hydroxide ion concentrations in the soda?
If the pH of a solution of the strong base Sr(OH)2 is 10.46, what is the concentration of the Sr(OH)2 in mol/L?

14. Homework
After reading sections , you should be able to do the following…
P. 629 (4-14)

15. 16.3 Equilibrium Constants
The lower the pH the stronger the acid.
For a strong base/acid, the [OH-]/[H3O+] concentration is equal to the original base/acid concentration.
For a weak base/acid, the ion concentration is much less than the original acid concentration.

16. Equilibrium Constants
For the general acid HA, we can write
Ka = [H3O+][A-]/[HA]
where the Ka is the equilibrium constant for an acid in water.
K is less than 1 for weak acids. Ka increases as acid strength increases.

17. Equilibrium Constants
For a weak base B in water…
Kb = [BH+][OH-]/[B]
Ionization Constants for Some Acids and their Conjugate Bases are on page 595.
The weaker the acid, the stronger its conjugate base.

18. Practice Problem Which is the stronger acid, H2SO4 or H2SO3?
Is benzoic acid, C6H5CO2H, stronger or weaker than acetic acid?
Which has the stronger conjugate base, acetic acid or boric acid?
Which is the stronger base, ammonia or the acetate ion?
Which has the stronger conjugate acid, ammonia or the acetate ion?

19. Ka and Kb Values for Polyprotic Acids
Ka values for each successive ionization step become smaller because it is more difficult to remove H+ from a negatively charged ion
Successive Kb values for each ionization become smaller as well

20. pKa
The negative log of the Ka value (pKa) becomes smaller as the acid strength increases.
pKa = -logKa

21. Practice Problem What is the pKa value for benzoic acid, C6H5CO2H?
Is chloroacetic acid (ClCH2CO2H), pKa = 2.87, a stronger or weaker acid than benzoic acid?
What is the pKa for the conjugate acid of ammonia? Is this acid stronger or weaker than acetic acid?

22. Relating Ionization Constants
As Ka decreases, Kb increases. The product of the two is equal to the autoionization constant for water.
Kw = KaKb

23. Practice Problem
Ka for lactic acid, CH3CHOHCO2H, is 1.4 x What is Kb for the conjugate base of this acid, CH3CHOHCO2-? Where does this base fit in Table 16.2?

24. 16.4 Acid-Base Properties of Salts
Anions that are conjugate bases of strong acids are such weak bases that they have no effect on solution pH. (Cl-, NO3-)
There are numerous basic anions; all are the conjugate bases of weak acids. (C2H3O2-)
Acidic anions arise from polyprotic acids, and are amphiprotic as well. (HCO3-)
Alkali metal and alkaline earth cations have no measurable effect on solution pH.
All metal cations are hydrated in water. Only when the cation is a +2 or +3 does the ion act as an acid. ([Al(H2O)6]3+ is acidic)

25. Practice Problem
For each of the following salts in water, predict whether the pH will be greater than, less than, or equal to 7.
KBr
NH4NO3
AlCl3
Na2HPO4

26. 16.5 Predicting the Direction of Acid-Base Reactions
All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base. (Equilibrium favors the weaker acid and base.)

27. Practice Problem
Which is the stronger Brønsted acid, HCO3- or NH4+? Which has the stronger conjugate base?
Is a reaction between HCO3- ions and NH3 product- or reactant-favored?
HCO3-(aq) + NH3(aq)  CO32-(aq) + NH4+(aq)

28. Practice Problem
Write the net ionic equation for the possible reaction between acetic acid and sodium hydrogen sulfate, NaHSO4. Does the equilibrium lie to the left or right?

29. Homework
After reading sections , you should be able to do the following…
P. 629b (19-37 odd)

30. 16.6 Types of Acid-Base Reactions
Strong Acid with a Strong Base
mixing equal quantities will produce a neutral solution
Weak Acid with a Strong Base
mixing equal quantities produces a basic solution; pH depends on Kb for anion produced

31. Types of Acid-Base Rxns
Strong Acid with a Weak Base
mixing equal quantities produces an acidic solution; pH depends on Ka for the cation produced
Weak Acid with a Weak Base
mixing equal quantities produces a solution in which the pH depends on the Ka and Kb for the cations and anions produced

32. Practice Problem
Equal molar quantities of HCl and NaCN are mixed. Is the resulting solution acidic, basic, or neutral?
Equal molar quantities of acetic acid and sodium sulfite, Na2SO3, are mixed. Is the resulting solution acidic, basic, or neutral?

33. 16.7 Calculations
You can use an ICE table to calculate K from initial concentrations and measured pH.

34. Strategy
Write the K expression, set up an ICE table, and convert pH to [H3O+].
Enter initial concentration.
Assign “x” to represent changes, based on reaction stoichiometry and enter into ICE.
Recognize that [H3O+] = x
Enter expressions for equilibrium concentrations into ICE
Solve for Ka

35. Practice Problem
A solution prepared from mol of butanoic acid dissolved in sufficient water to give 1.0L of solution has a pH of Determine Ka for butanoic acid. The acid ionizes according to the balanced equation
CH3CH2CH2CO2H + H2O  H3O+ + CH3CH2CH2CO2-

36. Equilibrium Constants
Due to the fact that very little ionization occurs in a weak acid, we can assume that the acid concentration at equilibrium is basically the same as the initial acid concentration.
This is valid whenever [HA]0 is greater than or equal to 100Ka.

37. pH of Weak Acid or Base
We can use ICE tables to calculate the pH of a solution of a weak acid or base and using known equilibrium constants.

38. Practice Problem
What are the equilibrium concentrations of acetic acid, acetate ion, and H3O+ for a 0.10M solution of acetic acid (Ka = 1.8x10-5)? What is the pH of the solution?

39. Practice Problem
What are the equilibrium concentrations of HF, fluoride ion, and H3O+ when a M solution of HF is allowed to come to equilibrium? What is the pH of the solution?

40. Practice Problem
Sodium hypochlorite, NaOCl, is used as a disinfectant in swimming pools and water treatment plants. What are the concentrations of HOCl and OH- and the pH of a 0.015M solution of NaOCl?

41. pH After Acid-Base Reaction
In order to calculate pH in a resulting solution, you must write a balanced equation and decide whether the products are acid or basic.
You must then find initial concentrations and can calculate pH by solving an equilibrium problem.

42. Practice Problem
Calculate the pH after mixing 15mL of 0.12M acetic acid with 15mL of 0.12M NaOH. What are the major species in solution at equilibrium (besides water) and what are their concentrations?

43. Homework
After reading 16.6 and 16.7, you should be able to do the following…
P. 629 (41-42, 47-48, 55-58, 62-63)

44. 16.8 Polyprotic Acids and Bases
Acids that can donate more than one proton are polyprotic.
Each step has its own Ka, which becomes progressively smaller due to the increased energy required to remove a proton.
The pH of many polyprotic acids depends primarily on the hydronium ion generated in the first ionization step, hydronium produced in the second step can be neglected.

45. Practice Problem
What is the pH of a 0.10M solution of oxalic acid, H2C2O4? What are the concentrations of H3O+, HC2O4-, and the oxalate ion, C2O42-?

46. 16.9 Molecular Structure, Bonding, and Acids-Bases
Stronger acids have weak H-X bonds (such as HCl) and weaker acids have strong H-X bonds (such as HF).
Oxoacids, such as HNO3, contain an atom bonded to one or more oxygen atoms, some with hydrogen atoms attached.
Inductive effect: the attraction of electrons from adjacent bonds by more electronegative atoms

47. Bonding and Acids/Bases
Carboxylic acids, hydrated metal cations, and anions all act as Brønsted bases.
Organic amines, such as ammonia compounds, act as Brønsted bases.

48. Practice Problem Which is the stronger acid, H2SeO3 or H2SeO4?
Which is the stronger acid, Fe(H2O)62+ or Fe(H2O)63+?
Which is the stronger acid, HOCl or HOBr?

49. 16.10 Lewis Acids and Bases
This concept is based on the sharing of electron pairs between acid and base.
A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond.
A Lewis base is a substance that can donate a pair of electrons to another atom to form a new bond.

50. Lewis Acids and Bases
The product of an acid-base reaction in the Lewis sense is often called an acid-base adduct. This type of bond is called a coordinate covalent bond.
The formation of hydronium and ammonium are both examples.

51. Cationic Lewis Acids
All metal cations form hydrated cations in which the metal ion is surrounded by water molecules, such a [Fe(H2O)6]2+.
These structures are called complex ions, or coordination complexes.
Hydroxide ion is a Lewis base and binds readily to metal cations to form metal hydroxides. Metal hydroxides are usually amphoteric. See table on p. 625.

52. Molecular Lewis Acids
Acidic oxides such as carbon dioxide and sulfur dioxide.
Due to oxygen’s high electronegativity, electrons are polarized away from the other element which can therefore react with a Lewis base such as hydroxide ion.

53. Practice Problem
Describe each of the following as a Lewis acid or a Lewis base. (Draw the Lewis dot structure. Are there lone pairs on the central atom? If so, it may be a Lewis base. Does the central atom lack an electron pair? If so it can behave as a Lewis acid.
PH3, BCl3, H2S, HS-

54. Homework
After reading sections , you should be able to do the following…
P. 629d-e (67-85 odd)