1. Chapter 16: Acids and Bases

2. Students will learn… 3 definitions of Acids and Bases
Acid Strength – pH scale
Water as acid and base
Calculating pH of strong acids and bases
Buffered solutions
*Review naming acids, Pg 132 & 133

3. Three Definitions of Acids-Bases
Arrhenius
BrØnsted-Lowry
Lewis

4. 3 Definitions of Acids and Bases
Arrhenius
Produce H+ ion* (= proton)
Produce OH– ion
BrØnsted-Lowry
Donate H+ ions
Accept H+ ions
Lewis
Accept an electron-pair
Donate an electron-pair
*H+ ion always combines with water molecule, forming hydronium ion, H3O+

5. Arrhenius Acids: HCl, HC2H3O2, HCN Arrhenius Bases: NaOH, Ca(OH)2
BrØnsted-Lowry Acids and Bases: can only be determined from chemical equations
H2O + HCl → H3O+ + Cl-
base acid conjugate conjugate
acid base
conjugate pair
conjugate pair

6. Conjugate Acid-Base Pair
An acid and a base that only differ by one H+ ion
Always made of one reactant and one product
Often in reversable reactions
(Ex) HSO4- + H2O ↔ SO42- + H3O+
acid base base acid
HSO4- + H2O ↔ H2SO4 + OH-
base acid acid base

7. Examples
1. Which of the following represent conjugate acid–base pairs? a) HCl, HNO3 b) H3O+, OH– c) H2SO4, SO42– d) HCN, CN–

8. 2. Write the conjugate base for each of the following.
HClO4
H3PO4
CH3NH3+

9. Strong vs. Weak Strong: 100% dissociation of H+ or OH–
Strong electrolytes (conduct electricity very well)
Weak: Less than 100% dissociation of H+ or OH–
Weak electrolytes (conduct electricity somewhat)

10. Common Strong and Weak Acids/Bases
Strong acids: H+7A, H2SO4, HClO4, HNO3
Weak acids: HCN, HC2H3O2
Strong bases: 1A+OH, Sr(OH)2, Ba(OH)2
Weak bases: NH3 (ammonia) in water = NH4OH (ammonium hydroxide)

11. Relative Strength of Acids and Bases

12. Example
Acetic acid (HC2H3O2) and HCN are both weak acids. Acetic acid is a stronger acid than HCN. Arrange these bases from weakest to strongest : Cl– CN– C2H3O2–

13. Terms related Acids
Amphoteric: Ions or molecules that can become an acid and a base
(Ex) HSO4- + H2O → SO42- + H3O+
acid base
HSO4- + H2O → H2SO4 + OH-
base acid
(Ex) H2O + H2O → H3O+ + OH-
Oxyacids: acids containing oxygen
(Ex) HNO3, H2SO4, HClO

14. Monoprotic Acids Diprotic Acids Triprotic Acids
Have one proton (H+) to donate
(Ex) HCl, HNO3, HI, HClO3
Diprotic Acids
Have more than two protons to donate
(Ex) H2SO4
Triprotic Acids
Have three protons to donate
(Ex) H3PO4

15. Self-ionization of Water
In neutral water,
H2O + H2O ↔ H3O+ + OH–
[H3O+] = [OH–] = 1×10-7 mol/L
[H3O+] · [OH–] = (1×10-7)·(1×10-7)=1× @ 25 °= Kw
Kw = ion-product constant for water
*The constant changes to the temperature

16. Acidic or Basic Solutions
Kw =1×10-14 always maintained
[H3O+]
[OH-]
pure water
[H3O+]=1×10-7 mol/L
[OH-]= 1×10-7 mol/L
acidic solution
[H3O+]>1×10-7 mol/L
[OH-]< 1×10-7 mol/L
basic solution
[H3O+]<1×10-7 mol/L
[OH-]> 1×10-7 mol/L

17. Common Misunderstanding
An acidic solution has only H+ ions
A basic solution has only OH‒ ions

18. Example
(1) In an acidic aqueous solution, which statement below is correct?
a) [H+] < 1.0 × 10–7 M
b) [H+] > 1.0 × 10–7 M
c) [OH–] > 1.0 × 10–7 M
d) [H+] < [OH–]

19. (2) In an aqueous solution in which [OH–] = 2
(2) In an aqueous solution in which [OH–] = 2.0 x 10–10 M, the [H+] = ________ M, and the solution is ________.
a) 2.0 × 10–10 M ; basic
b) 1.0 × 10–14 M ; acidic
c) 5.0 × 10–5 M ; acidic
5.0 × 10–5 M ; basic

20. pH and pOH Scale pH = –log[H+] [H+] = 10-pH pOH = ‒log [OH‒]
A compact way to represent solution acidity
pH decreases as [H+] increases
[H+] = 10-pH
pOH = ‒log [OH‒]
pOH decreases as [OH‒] increases
[OH‒] = 10-pOH
pH + pOH = 14

21. pH = 7; neutral
pH > 7; basic
Higher the pH, more basic.
pH < 7; acidic
Lower the pH, more acidic.

22. Acidity and pH Scale
high pH = weak acid
high pOH = weak base

23. Examples (1) Calculate the pH for each of the following solutions.
1.0 × 10–4 M H+
b) M OH–

24. (2) The pH of a solution is 5.85. What is the [H+] for this solution?

25. (3) Consider an aqueous solution of 2. 0 × 10–3 M HCl. What is the pH
(3) Consider an aqueous solution of 2.0 × 10–3 M HCl. What is the pH? *If the solution was 2.0 × 10–3 M HNO2, a weak acid, you can’t take the same approach. Why?

26. Buffered Solutions
resists a change in its pH when either an acid or a base has been added.
Presence of a weak acid and its conjugate base buffers the solution.

27. Characteristics of Buffered Solutions
The solution contains a weak acid HA and its conjugate base A–.
The buffer resists changes in pH by reacting with any added H+ or OH– so that these ions do not accumulate.
Any added H+ reacts with the base A–.
H+(aq) + A–(aq) → HA(aq)
4. Any added OH– reacts with the weak acid HA.
OH–(aq) + HA(aq) → H2O(l) + A–(aq)

28. Example
If a solution is buffered with NH3 to which has been added NH4Cl, what reaction will occur if a strong base such as NaOH is added?
a) NaOH + NH3 → NaNH4O
b) NaOH + NH4+ → Na+ + NH3 + H2O
c) NaOH + Cl– → NaCl + OH–
d) NaOH + H2O → NaH3O2