1. Acids and Bases

2. Acids and Bases
Acid: when dissolved in water, increases the concentration of H+ (arrenhius acid)
HCl H+ + Cl-
HCl + H2O H3O+ + Cl-
Strong acid: an acid that completely ionizes/dissociates in water
HNO3 + H2O H3O+ + NO3-
Strong electrolyte
H2O

3. Bronsted-Lowry definition

5. Strong Acids: Hydrochloric Acid HCl Nitric Acid HNO3
Sulfuric Acid H2SO4
Perchloric Acid HClO4
Hydrobromic Acid HBr
Hydroiodic Acid HI

7. Weak Acid: an acid that only partially ionizes/dissociates in water
CH3COOH(aq) CH3COO-(aq) + H3O+(aq)
Weak electrolytes
Weak Acids:
Phosphoric acid H3PO4
Acetic Acid CH3COOH
Carbonic Acid H2CO3
Hydrocyanic Acid HCN
Benzoic Acid C6H5COOH

9. Polyprotic Acids
Polyprotic acids: acids that can release more than one H+
Sulfuric Acid
H2SO4(aq) HSO4-(aq) + H3O+(aq)
HSO4-(aq) SO42-(aq) + H3O+(aq)

10. Bases
Base: a substance that, when put in water, increases the concentration of OH- ions or a substance that accepts H+ ions
NaOH(aq)  Na+(aq) + OH-(aq)
Strong Bases: bases that completely ionize in water

13. Weak Bases: bases that only partially ionize in water
Ammonia: NH3
Pyridine: C5H5N

16. Each acid has a conjugate base and every base has a conjugate acid
Congugate Acid-Base Pairs
Each acid has a conjugate base and every base has a conjugate acid
conjugate acid-base pair 1
HA + B A− + BH+
conjugate acid-base pair 2

17. Water is amphoteric – can act as an acid or a base

19. Neutralization Reactions
When strong acids and bases in aqueous solution react with each other, they form water and a salt.
HX(aq) + MOH(aq)  HOH(l) + MX(aq)
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
Water
Salt

20. Acid Ionization Constant
Acid Ionization Constant (Ka): the equilibrium constant for the ionization reaction of an acid with water
HA + H2O A- + H3O+
Large Ka = Strong acid
Small Ka = Weak acid

22. Base Ionization Constant
Base Ionization Constant (Kb): the equilibrium constant for the ionization reaction of a base with water
B + H2O OH- + BH+
Large Kb = Strong base
Small Kb = Weak base

26. Autoionization of Water

27. Autoionization of Water
Also called “Self Ionization”
About 1 out of every 10 million water molecules form ions through self ionization
H2O Û H+ + OH–
H2O + H2O Û H3O+ + OH–
All aqueous solutions contain both H3O+ and OH–

28. Ion Product Constant for Water
Ion Product Constant for Water (Kw):
the numerical value obtained by multiplying the molar concentrations for hydronium and hydroxide ions present in pure water at 25°C
Kw = [H3O+][OH-] = 1.00 x at 25 oC
the concentration of H3O+ and OH– are equal in pure water
[H3O+] = [OH–] = 25°C

29. Ion Product of Water
the product of the H3O+ and OH– concentrations is always the same number
Kw =[H3O+][OH–] = 1.00 x 25°C
if you measure one of the concentrations, you can calculate the other
as [H3O+] increases the [OH–] must decrease so the product stays constant
inversely proportional

30. H+ OH- [H+] vs. [OH-] Acid Base
Even though it may look like it, neither H+ nor OH- will ever be 0

31. Acidic and Basic Solutions
Neutral solutions have equal [H3O+] and [OH–]
[H3O+] = [OH–] = 1 x 10-7
acidic solutions have a larger [H3O+] than [OH–]
[H3O+] > [OH–]
[H3O+] > 1 x 10-7; [OH–] < 1 x 10-7
basic solutions have a larger [OH–] than [H3O+]
[H3O+] < [OH–]
[H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

32. pH and pOH Acidic solutions Basic solutions Neutral solutions

34. pH + pOH = 14 pH = -log [H+] OR pH = -log [H3O+] [H3O+] = 10-pH
pH is a measure of the concentration of H3O+ in solution
pH = -log [H+] OR pH = -log [H3O+]
[H3O+] = 10-pH
pOH is a measure of the concentration of OH- in solution
pOH = -log [OH-]
[OH-] = 10-pOH
pH + pOH = 14

35. Classification of Water Soluble Substances
Electrolytes: solutes that separate into ions when dissolved in water (they’re soluble)
Have the ability to conduct electricity
2 types
Strong electrolytes
Weak electrolytes

36. Strong electrolytes: solutes that completely dissociates into ions when dissolved in water
Examples: NaCl, MgBr2, HCl
Strong electrical conductors
Strong electrolyte(aq or s) → Cation+(aq) + Anion-(aq)
Example: NaCl(s) → Na+(aq) + Cl-(aq)

37. Examples: HF, NH3, acetic acid Weak electrical conductors
Weak Electrolytes: solutes that, when dissolved in water, only partially dissociates into ions
Examples: HF, NH3, acetic acid
Weak electrical conductors
Weak electrolyte(aq) ↔ Cation+(aq) + Anion-(aq)
Example: HF(aq) ↔ H+(aq) + F-(aq)

38. Nonelectrolyte (s or l) → Nonelectrolyte(aq)
Nonelectrolytes: solutes that dissolve in water without separating into ions
Examples: sucrose, ethanol
Do not conduct electricity
Nonelectrolyte (s or l) → Nonelectrolyte(aq)
Example: C12H22O11(s) → C12H22O11(aq)

39. Titration
Titration: a procedure for the quantitative analysis of a substance of unknown concentration whereby a measured quantity of another substance, of know concentration, is completely reacted with the with the original substance.
Often used to determine the concentration of acids and bases

41. Equivalence point: the point in a titration at which one reactant has been exactly consumed by the addition of another reactant
Midpoint of vertical rise
Occurs at pH = 7 in a strong acid-strong base titration
[H3O+] = [OH-]

42. Indicators
Acid-Base Indicator: a chemical that changes color with a change in pH
Added to solutions in small amounts in order to determine to solution’s pH visually
Usually organic compounds
Weak acid or base
establishes an equilibrium with the H2O and H3O+ in the solution 42

43. HInd(aq) + H2O(l)  Ind(aq) + H3O+(aq)

44. Phenolphthalein 44 44

45. Bromocresol Green
Yellow
Green
Blue

46. Methyl Red 46