1. Acids, Bases, and Salts

2. Objectives Know the fundamental properties of acids and bases.
Know the ions associated with acids and bases.
Know how acids and bases are produced.

3. Properties of Acids contain H (HCl, H2SO4, etc.)
produce hydronium ions (H3O+) in water
electrolytes
pH < 7
sour taste
react with metals to make H2 gas
Zn + 2HCl → H2 + ZnCl2
made from non-metal oxide and water: SO3 + H2O →H2SO4

4. Properties of Bases often contain OH (NaOH, Ca(OH)2, etc.)
bases produce hydroxide ions (OH–) in water
electrolytes
pH > 7
bitter taste
feel slippery
made from metal oxide and water: ZnO + H2O →Zn(OH)2
Acids and bases
“neutralize” each other!
HCl + NaOH → NaCl + H2O

5. Objectives Be able to name acids.
Be able to identify Brønsted-Lowry acids, bases, conjugate acids, and conjugate bases.
Understand and correctly apply the meaning of the term amphoteric.

6. Acid Nomenclature USE YOUR YELLOW SHEET!
use the stem and ending of the anion name
-ide hydro-stem-ic acid
-ate stem-ic acid
-ite stem-ous acid
HCl = H+ + Cl– (chlor-ide) = hydrochloric acid
HNO3 = H+ + NO3– (nitr-ate) = nitric acid
HNO2 = H+ + NO2– (nitr-ite) = nitrous acid
Common exceptions:
sulfuric (H2SO4) and phosphoric (H3PO4)

7. Original Definitions of A/B
Acids contain H, dissolve to form H+
Bases dissolve to form OH-
HCl (g) → H+ (aq) + Cl− (aq)
NaOH (s) → Na+ (aq) + OH− (aq)
…but this is over-simplified.

8. Brønsted-Lowry Acids and Bases
acid: proton (H+) donor
base: proton (H+) acceptor
HCl(g) + H2O(l) ↔ Cl−(aq) + H3O+(aq)
hydronium ion
acid
base
conjugate
base
conjugate
acid
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH−(aq)
conjugate
acid
conjugate
base
base
acid
HNO3(aq) + NH3(aq) ↔ NO3−(aq) + NH4+(aq)
acid
base
conj base
conj acid

9. Water: Acid and Base!
amphoteric: a substance that can act as either an acid or a base (such as water)
hydronium ion
H3O+ = +
ACID +
BASE
conj. acid
BASE +
hydroxide ion
OH- = −
ACID
conj. base

10. Objectives Understand the process of self-ionization.
Understand how the concentrations of hydronium and hydroxide ion can vary in water.
Understand the concept of pH.
Be able to make pH calculations using the log and 10x functions on a calculator.

11. Self-Ionization of Water
H2O + H2O ↔ OH− + H3O+ (reactant strongly favored)
[OH− ] = 10-7 M and [H3O+] = 10-7 M
KW = [OH−]·[H3O+]
water: Kw = [10-7]·[10-7] = 10-14
acidic: Kw = [10-9]·[10-5] = 10-14
basic: Kw = [10-3]·[10-11] = 10-14
[OH−] and [H3O+] are inversely proportional, KW = 10-14

12. [H3O+] ACIDS [H3O+] > 10-7 M memorize this!
HNO3 (g) + H2O (l) →H3O+ (aq) + NO3− (aq)
[H3O+] = 10-6 M, 10-5 M, 10-4 M, … 10-1 M or more
BASES [H3O+] < 10-7 M memorize this!
NaOH(s) → Na+(aq) + OH−(aq) OH− removes H3O+
[H3O+] = 10-8 M, 10-9 M, M, … M or less

13. pH Scale pH = −log[H3O+] [H3O+] = 10-3 M, pH = 3 (acidic)
[H3O+] = 10-7 M, pH = 7 (neutral)
[H3O+] = M, pH = 11 (basic)
Calculating pH?
[H3O+] = 5.7 x 10-2 M pH = −log(5.7E-2) = 1.2
Calculating [H3O+]? Use [H3O+] = 10-pH
If pH = [H3O+] = = 1.6 x 10-4 M

14. Objectives
Understand the distinction between natural acid rain and human-caused (anthropogenic) acid rain.
Know the types of pollution that result in acid rain formation.
Understand the effects of acid rain and how they can be reduced.

15. Acid-Base Indicators ↔ compounds that change color at a specific pH
phenolphthalein is a typical example…
indicator
anion
indicator
anion H+ ↔
H3O+ +
+ H2O
ACID = clear
CONJ BASE = pink
add base (removes H3O+) = pink in high pH
add acid (add H3O+) = clear in low pH
transition occurs around pH = 9

16. Acid-Base Indicators
many plants contain indicators: hydrangea, red cabbage
universal indicator: mixture w/ wide pH range

17. Plant Dyes and pH
serviceberry, willow bark, & Oregon grape root are local plants that contain indicators
used as natural dyes for skins, feathers, etc.
Coastal Salish

18. Objectives
Understand the concept of KA and how it relates to strong and weak acids.
Be able to calculate the KA of an acid solution if given the initial molarity and the pH of the solution.

19. Strengths of Acids strong acid: completely ionizes in water
products favored (hydronium ions and acid anions)
HNO3 (g) + H2O (l) → H3O+(aq) + NO3−(aq)
weak acid: partially ionizes in water
reactants favored (molecular form of acid)
HC2H3O2(l) + H2O (l) ↔ H3O+(aq) + C2H3O2−(aq)

20. Acid Dissociation Constant (KA)
HA + H2O ↔ H3O+ + A−
(A is the acid anion)
Strong acids—high KA ( > 1, products favored)
Weak acids—low KA ( < 1, reactants favored)
Note that [H3O+] = [A−] * use [H+] = 10-pH
[HA] = C – [H+]

21. Calculating KA
The initial concentration of an HNO2 solution is M.
What is the KA of HNO2 if the pH of the solution is 1.93?
Determine [H3O+] (same value as [A-] )
[H3O+] = 10-pH = = M
Determine [HA]
[HA] = C – [H3O+] = M – M = M
Calculate KA
KA < 1, weak acid

22. Objectives
Be able to explain the distinction between strong and weak acids versus concentrated and dilute solutions.
Understand the concept of acid neutralization and be able to determine the products of an acid-base neutralization reaction.
Be able to calculate either acid or base concentration using data from an acid-base titration.

23. Strength vs. Concentration
strength relates to degree of ionization (KA)
concentration relates to amount of solute (molarity)
strong = product favored (H+, A-)
weak = reactant favored (HA)
concentrated = lots of solute
dilute = not much solute

24. Neutralization acid + base → “salt” + water H+ + OH− → H2O (neutral)
salt: ionic compound consisting of a base cation and an acid anion
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l)
Try this one…
HNO3 (aq) + Ca(OH)2 (aq) → ? + ?
2HNO3 (aq) + Ca(OH)2 (aq) → Ca(NO3)2 (aq) + 2H2O (l)

25. Acid-Base Titration
standard solution (known concentration) is added to an unknown solution until pH = 7
the concentration of the unknown can be calculated

26. Titration Calculation
What is the concentration of H2SO4 if 10.0 mL is
completely neutralized by 14.2 mL of 1.0 M NaOH?

27. Buffers
buffer: a solution in which the pH remains relatively constant when a small amount of acid or base is added
consists of weak acid (or base) and one of its salts
Example: Your blood pH (= 7.2) is maintained by H2CO3/HCO3− buffer
Add acid: H+ + HCO3− → H2CO3
Add base: H2CO3 + OH− → HCO3− + H2O

28. Acid Rain normal rainfall is slightly acidic (pH = 5-6)
CO2 + H2O → H2CO3
acid rain: precipitation with a pH < 5
burning “high-sulfur” coal produces SO2 & SO3
SO2 + H2O → H2SO3
cars make NOX that form HNO2 & HNO3

29. Impacts of Acid Rain destroys decomposes forests limestone
kills aquatic life
corrodes metal

30. Acid Rain in the USA

31. Neutralizing Acid Rain
Limestone bedrock neutralizes acid, reducing environmental damage.
Granite does not. Bases such as CaO or CaCO3 must be used to neutralize acids.
H2SO4 + CaCO3 → CaSO4 + H2O + CO2