1. Acids, Bases, and Salts

2. Objectives Know the fundamental properties of acids and bases.
Be able to identify an Arrhenius acid.
Be able to write a dissociation equation for an Arrhenius acid.
Be able to identify monoprotic, diprotic, and triprotic acids.

3. Properties of Acids ACIDS pH < 7 sour taste electrolytes
react with metals to make hydrogen gas
Zn + 2HCl → H2 + ZnCl2
often formed from non-metal oxide and water
SO3 + H2O →H2SO4

4. Properties of Bases BASES pH > 7 bitter taste electrolytes
feel slippery
often formed from metal oxide and water
ZnO + H2O →Zn(OH)2
can form from group 1 or 2 metal and water
Ca + H2O →Ca(OH)2
Acids and bases
“neutralize” each other!
HCl + NaOH → NaCl + H2O

5. Arrhenius Acids What explains acidic properties?
Arrhenius acid: contains hydrogen and produces H+ ions in water
Acids are molecular compounds that ionize in water.
HCl(g) →H+(aq) + Cl–(aq)
H2SO4(l) →H+(aq) + HSO4–(aq)
monoprotic: HC2H3O2
diprotic: H2CO3
triprotic: H3PO4
Svante Arrhenius

6. Objectives
Be able to properly name an acid when given the formula, and be able to write the proper formula when given the name on an acid.
Be able to identify Arrhenius bases.
Be able to write a dissociation equation for an Arrhenius base.

7. Acid Nomenclature 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)

8. Arrhenius Bases
Bases dissociate to form OH- (hydroxide) ions when aqueous. Phenolphthalein indicates
NaOH(s) → Na+(aq) + OH-(aq)
Mg(OH)2(s) → Mg2+(aq) + 2OH-(aq)
phenolphthalein indicates OH-
problem: why is ammonia basic?

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

10. 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

11. Water: Acid or Base!
amphoteric: a substance that can act as either an acid or a base (such as water)
H+ is really H3O+ because water bonds with H+ + = +
hydronium ion
base
conj. acid = +
hydroxide ion −
acid
conj. base

12. 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.

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

14. [H3O+] ACIDS [H3O+] > 10-7 M
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
NaOH(s) → Na+(aq) + OH−(aq) OH− reduces H3O+
[H3O+] = 10-8 M, 10-9 M, M, … M or less

15. 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

16. Objectives Understand how acid precipitation forms.
Understand the effects of acid precipitation and how they can be reduced.

17. Acid Rain, Acid Fog acid rain: precipitation with a low pH (< 5)
burning “high-sulfur” coal produce SO2 and SO3 that react w/ H2O to make H2SO3 and H2SO4
cars make NOX: reacts w/ H2O to make HNO2 and HNO3
corrodes metal
dangerous to
organisms
decomposes
limestone

18. Acid Rain in the USA

19. 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

20. Objectives Be able to identify strong and weak acids.
Understand how acid-base indicators work.

21. Strengths of Acids
strong acid: completely ionizes in water, products favored
HNO3 (g) + H2O (l) → H3O+(aq) + NO3−(aq)
weak acid: partially ionizes in water, reactants favored
HC2H3O2(l) + H2O (l) ↔ H3O+(aq) + C2H3O2−(aq)

22. Acid-Base Indicators ↔
compounds that respond to pH change by changing color
contain a “weak acid” that exists in a chemical equilibrium
indicator
anion
indicator
anion H+ ↔
H3O+ +
+ H2O
ACID=clear
CONJ BASE=pink
add base (removes H3O+) = solution turns pink in high pH
add acid (add H3O+) = solution turns clear in low pH
Universal indicators are mixtures that show a range
of pH

23. Plant Dyes and pH
serviceberry, willow bark, Oregon grape root, contain indicators. These plants have traditionally been used as natural dyes for skins, feathers, and so on.

24. 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.

25. Acid Dissociation Constant (KA)
HA ↔ H+ + A−
Strong acids—high KA ( > 1, products favored)
Weak acids—low KA ( < 1, reactants favored)
Note that [H+] = [A− ] * use [H+] = 10-pH
[H+] = [H3O+]
[HA] = initial molarity – [H+]

26. 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 [H+] (same value as [A-] )
[H+] = 10-pH = = M
Determine [HA]
[HA] = initial – [H+] = M – M = M
Calculate KA
KA < 1, weak acid

27. Objectives
Be able to explain the distinction between strong and weak acids, as well as concentrated and dilute solutions.
Understand the concept of acid neutralization and be able to determine the products of an acid-base neutralization reaction.

28. Strength vs. Concentration
strength relates to degree of ionization (KA)
concentration relates to amount of solute (M)
strong
weak
concentrated
dilute

29. Objectives
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.

30. Neutralization acid + base → salt + water H+ + OH− → H2O
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)

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

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

33. 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