1. Unit 12: Acids, Bases, and Salts

2. Properties of Acids Taste sour
Are electrolytes: substances that produce ions in solution and as a result conduct electricity
Strong acid  strong electrolyte
Weak acid  weak electrolyte
Will react with active metals (above H2) to produce H2(g)

3. Review of Naming Rules: from Unit 5!
Binary Acids:
hydro_______ic acid
Ternary Acids (use Table E)
ate ic
ite ous

4. Properties of Bases
Strong bases consist of a Group 1 or Group 2 metal w/ hydroxide ion
Taste bitter, feel slippery
Also electrolytes
Strong bases  strong electrolytes
Weak bases  weak electrolytes
Will NOT REACT with active metals

5. Look on RT for common acids and bases

6. Electrolytes Form ions in water Conduct electricity
Salts
Ionic Compounds
M to NM
Acids
pH < 7
H+ > OH-
Bases
pH > 7
OH- > H+
Molecules (NM to NM) are NOT electrolytes…do NOT produce ions in solution!!

7. Indicators- Table M Below the range  the color is given on the left
Above the range  the color is given on the right.

8. Arrhenius Theory of Acids and Bases
Arrhenius Acid
Arrhenius Base
A substance, that when dissolved in water, produces H+ ions as the only positive ions in solution
A substance, that when dissolved in water, produces OH- ions as the only negative ions in solution
Monoprotic, diprotic, triprotic
Hydronium ion
HCl (aq)  H+ (aq)+ Cl- (aq)
NaOH (aq)  Na+ (aq) + OH- (aq)
known as a hydrogen ion
known as a hydroxide ion

9. Bronsted-Lowry Theory of Acids and Bases ***based on the proton (H+)
Bronsted Acid
Bronsted Base
A proton donor
A proton acceptor
Must release an H+ (aq) ion
Does not need to have the OH-(aq) ion, but MUST have a lone pair of electrons.
Ex) HCl (aq)
Ex) NH3 (aq)

10. Reversible Acid-Base Rxns
Acid-Base reactions can be reversible! H2SO4(aq) + NH3(aq) HSO4-(aq) + NH4+(aq) B.A. B.B. B.B. B.A.
A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion (H+)
Strong acids have weak conjugate bases
Strong bases have weak conjugate acids
Pg.9

11. Amphoteric Substances
A substance that can act as both an acid and a base is said to be amphoteric
Ex) HCO3- (aq) + H2O (l)  CO32- (aq) + H3O+(aq)
**Here, HCO3- acts as an acid
Ex) HCO3- (aq) + H2O (l)  H2CO3 (aq) + OH- (aq)
**Here, HCO3- acts as a base
Water is a common amphoteric substance- can act as both an acid and base

12. Lewis Definition of Acids and Bases
Lewis Acid
Lewis Base
an electron pair ACCEPTOR
An electron pair DONOR
Ex) NH3 + HCl  NH4+ + Cl-

13. Strengths of Acids and Bases
STRONG acids/bases dissociate completely in solution
HCl(aq) H+(aq) + Cl-(aq)
Ka = [H+] [Cl-] = very large
[HCl]
STRONG acids have large Ka’s

14. Strength of acids and bases
CH3COOH(aq) CH3COO-(aq) + H+(aq)
Ka = [CH3COO-] [H+] = very small (less than 1)
[CH3COOH]
Not a lot of H+, therefore, weak acid
weak acids have small Ka’s

15. MEMORIZE STRONG ACIDS Weak acidsHI
HBr
HCl
HNO3
H2SO4
HClO4
CH3COOH (vinegar)
H3PO4 HF
Strong acids have weak conjugate bases

16. MEMORIZE STRONG BASES Weak bases
Any group 1 or group 2 metal with OH-
NH3
Strong bases have weak conjugate acids

17. For diprotic and triprotic acids, the H+’s come off in steps
For diprotic and triprotic acids, the H+’s come off in steps. It is more difficult to remove subsequent H+’s.
Ex Ka
H3PO4  H+ + H2PO
H2PO4-  H+ + HPO
HPO42-  H+ + PO

18. Water as an Acid/Base Water molecules undergo self-ionization
H2O(l) + H2O (l) H3O+ (aq) + OH- (aq)
In pure water, the equilibrium concentration of [H3O+] = [OH-] = 1 x 10-7 M
In other words, a neutral solution has the same number of moles of H+ ions as OH- ions.

19. “See-Saw Diagram”
Ion-Product Constant for Water (for aqueous solutions) Kw= [H+] x [OH-] = 1.0 x 10-14
** exponents always add to 14

20. The pH scale
Scale is based on the concentration of H+ ions in aqueous solution
pH = - log [H+]
pH < 7, acidic
pH = 7, neutral
pH > 7, basic

21. Determining pH given concentrations
[H+] = 1.0 x 10-6 M
pH = 6
pH + pOH ALWAYS equals 14
[OH-] = 1.0 x 10-8 M
pOH = 8

22. The pH scale is logarithmic
An increase in one unit of the pH scale represents a ten-fold decrease (x 10) in the H+ ion concentration

23. Practice pH 6 to pH 4 pH 5 to pH 6 [H+] increases by a factor of 100
[H+] decreases by a factor of 10
pH 2 to pH 4
[H+] decreases by factor of 100

24. pH of weak acids Q. What is the pH of 0.10 M CH3COOH?
CH3COOH(aq)  H+(aq) + CH3COO-(aq)
0.10 M – x x x
Ka= [H+] [CH3COO-] = X = 1.8 x 10-5
[CH3COOH] x
x = 1.8 x 10-5
0.10-x
x2= 1.8 x 10-5(0.10)
x2= 1.8 x 10-6
x= 1.3 x 10-3
pH= -log(1.3 x 10-3)= 2.9
Very small. Ignore if 5% or less dissociation
0.10 M HCl pH=1
0.10 M CH3COOH pH=2.9

25. Neutralization Reactions
Neutralization occurs when an acid and a base react to form a salt and water
Acid + Base  salt + water
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Acid Base Salt Water
Net Rxn: H+ (aq) + OH- (aq)  H2O (l)

26. Titrations
Used to determine the concentration of an acid or base
The equivalence point is reached when the moles H+ = moles OH-
Titration Equation:
MaVa = MbVb
** Can use any unit of volume as long as you are consistent

27. unknown volume, unknown concentration
Known volume, known concentration

28. Indicators for Titrations
Phenolphthalein is used to approximate the endpoint of a titration.
Colorless indicates an acidic solution and pink indicates a basic solution
When a pale pink color persists, the endpoint has been reached
The endpoint, or equivalence point, is when moles H+ = moles OH- (pH =7)

29. Warm-Up
What is the pH of a M solution of Ca(OH)2? What is the hydroxide ion concentration of a M solution of HNO3 (aq)? What is the pH of a M solution of hydrofluoric acid? (Assume 3% dissociation)

30. Hydrolysis (reverse of neutralization)
The process in which ions of a salt react with water to produce an acidic or basic solution is called hydrolysis

31. Examples Example Parent Base Parent Acid NaCl NaOH HCl
Neutral  strong strong
NaCH3COO NaOH CH3COOH
Basic  strong weak
NH4NO NH4OH HNO3
Acidic weak strong
NH4NO NH4OH HNO2
pH? Depends on Ka/Kb values weak weak

32. Salt formed from pH
SB SA
neutral
SB WA
More than 7
WB SA
Less than 7
WB WA
? Need Ka/Kb

33. Basic Anhydrides
Group I and II metal oxides that can react with water to form basic solutions.
Ex. Na2O(s) + H2O(l)  2NaOH(aq)
Basic anhydride

34. Acid Anhydrides
Non-metal oxides that can react with water to form acidic solutions
Produce acid rain
Ex. SO3(g) + H2O(l)  H2SO4(aq)
Ex. N2O5(g) + H2O(l)  2HNO3(aq)
Ex. CO2(g) + H2O(l)  H2CO3(aq)

35. Chemistry of Dyeing Eggs
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