1. Acids and Bases
Chapters 20-21

2. Properties of Acids Sour taste
Change the color of dyes known as acid-base indicators (turns litmus red)
React with bases to produce a salt (ionic compound) and water
Electrolytes (aqueous solutions will conduct electric current)
Can react with active metals to give off hydrogen gas (H2)

3. Definitions of Acids
Arrhenius (traditional): compound that contains hydrogen and ionizes in solution to form hydrogen ions (H+)
Bronsted-Lowry: molecule or ion that is a proton (H+) donor
Lewis: atom or molecule that is an electron-pair acceptor

4. Number of Protons in Acids
Monoprotic acids can donate one H+ ion per molecule
Examples: HCl, HNO3, HBr, HClO4
Diprotic acids can donate two H+ ions per molecule
Examples: H2SO4, H2Se
Triprotic acids can donate three H+ ions per molecule
Example: H3PO4
Polyprotic acids can donate more than one H+ ion per molecule
H2SO4       H+ + HSO4-
HSO4-       H+ + SO42-

5. Hydronium Ion
Hydrogen ion (H+) in an aqueous solution (H2O) becomes H3O+
H3O+ is known as a hydronium ion
The terms hydrogen ion (H+), proton (p+), and hydronium ion (H3O+) are used interchangeably

6. Naming Acids (Review) Binary Acids contain two elements
(a cation, H+, and a monatomic anion)
hydro______ic acid
Oxyacids
contains more than two elements (H, O, and
3rd element)
(a cation, H+, and a polyatomic anion with O)
________ic acid (-ate ion goes with –ic ending)
________ous acid (-ite ion goes with –ous ending)

7. Oxyacid Naming Series per______ic acid most oxygen’s
______ic acid goes with –ate anion
______ous acid goes with –ite anion
hypo______ous acid least oxygen’s

8. Properties of Bases Bitter taste Change the color of dyes known
as acid-base indicators (turns litmus blue)
React with acids to produce a salt (ionic compound) and water
Electrolytes (aqueous solutions will conduct electric current)
Feel slippery to the skin
Basic substances are referred to being “alkaline."

9. Definitions of Bases
Arrhenius (traditional): compound that contains hydroxide and dissociates in solution to form hydroxide ions (OH-)
Bronsted-Lowry: molecule or ion that is a proton (H+) acceptor
Lewis: atom or molecule that is an electron-pair donor

10. Naming Bases (Review) Name as ionic compound
Cation named first, anion named second
Potassium Chloride KCl
All transition metals and Pb and Sn need Roman numerals to denote their charge. (Zn, Ag, and Cd don’t need Roman numerals.)
Change ending of all nonmetals to –ide to name their monatomic anion.
Polyatomic ions must be memorized! Review them if you have forgotten.

11. Name the Following: H3PO4 Ba(OH)2 HClO2 H3N Al(OH)3 HIO3 H2SO3 HBr
Cu(OH)2
H2S
phosphoric acid
barium hydroxide
chlorous acid
hydronitric acid
aluminum hydroxide
iodic acid
sulfurous acid
hydrobromic acid
copper (II) hydroxide
hydrosulfuric acid

12. Neutralization Acid + Base  Ionic Compound + H2O
The ionic compound formed from the reaction of an acid and a base is known as a salt. It is formed from the cation of the base and the anion of the acid.
The cation from the acid (H+) and the anion from the base (OH-) form the water
Hydrolysis: reaction between water and ions of a dissolved salt; causes water to dissociate into H+ and OH- ions

13. Write balanced equations:
barium hydroxide solution mixed with phosphoric acid
aqueous sodium hydroxide neutralized with hydrochloric acid
aluminum hydroxide solution mixed with chloric acid
sulfuric acid reacted with aqueous magnesium hydroxide

14. Answers: 3 Ba(OH)2 (aq) + 2 H3PO4 (aq)  Ba3(PO4)2 (s) + 6 H2O (l)
NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (l)
Al(OH)3 (aq) + 3 HClO3 (aq)  Al(ClO3)3 (aq) + 3 H2O (l)
H2SO4 (aq) + Mg(OH)2 (aq)  MgSO4 (aq) + 2 H2O (l)

15. Info for Hydrolysis Worksheet
strong acids (ionize 100% in solution):
HClO4, HNO3, H2SO4, HI, HBr, HCl
strong bases (ionize 100% in solution):
Hydroxides of Groups 1 and 2 on Periodic Table
If they’re not strong, we assume they are weak!

16. Bronsted-Lowry Conjugate Acid-Base Pairs
Remember, B-L Acid: H+ donor
B-L Base: H+ acceptor
When an acid is dissolved in water, the acid (HA) donates a proton to water to form a new acid (conjugate acid) and a new base (conjugate base)
HA (aq) + H2O (l)  H3O+ (aq) + A- (aq)
acid base conjugate acid conjugate base H+ H+

17. Bronsted-Lowry Acid Conjugate BaseHA
HCl
HNO3
H2SO4
HSO4-
H3O+
H2O A-
Cl-
NO3-
HSO4-
SO4-2
H2O
OH-

18. A few more items:
Amphoteric: any species that reacts as either an acid or a base
strong acids have weak conjugate bases
weak acids have strong conjugate bases
strong acids (ionize 100% in solution):
HClO4, HNO3, H2SO4, HI, HBr, HCl
strong bases (ionize 100% in solution):
Hydroxides of Groups 1 and 2 on Periodic Table

19. pH notes The pH scale ranges from 0 to 14.
Any substance with a pH below 7 is classified as an acid
Any substance with a pH above 7 is classified as a base.
Substances with a pH of 7 are said to be neutral.
(Pure water has a pH of 7.)
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| | | | | | | | | | | | | | |
strong acid weak acid neutral weak base strong base

20. Calculating pH
See typed notes

21. Concentration Comparisons
[H+] > [OH-]
[H+] < [OH-]
[H+] = [OH-]
acidic
neutral
basic

22. Titration
Titration is the method used to determine the concentration of a solution (usually an acid or base).
A solution of known concentration (the standard) is added to a measured amount of the solution of unknown concentration until an indicator signals the endpoint.
The endpoint occurs when equivalent amounts of H+ and OH- have reacted in a titration, thus neutralizing the resultant solution.

23. Indicators
Acid-base indicators are dyes used in titrations whose colors are sensitive to changes in pH, or hydronium ion concentration.
There are many indicators, each indicative of a different pH range. For most titrations, a neutralization reaction is desired.
For a strong acid and strong base, the chosen indicator will change near the neutral point at a pH of 7.
Examples: phenolphthalein (changes from clear in acid to pink in base) and bromothymol blue (changes from yellow in acid to blue in base)

24. Sample Titration 50. mL of 0.1 M HCl when titrated with
0.1 M NaOH should take just over 50. mL to cause a color change
because it is a monoprotic strong acid being neutralized by a group I hydroxide strong base.
(For this titration, we would choose an indicator that changed color at a pH of 7 or neutral to indicate that the acid had been fully neutralized by the added base.)
0.1 M H2SO4 substituted in the above problem would have a different effect since the H+ ion concentration would change for a diprotic acid

25. Titration Equation MaVa = MbVb nbMaVa = MbVbna
In titrations between acids and bases, neutralization is required so the [H+] = [OH-]
In order to take this into account, our equation becomes
nbMaVa = MbVbna
# of H+ ions in base
Molarity of acid
volume of acid
Molarity of base
volumeof base
# of OH- ions in acid

26. Functions as an Base or Acid
Amphoteric
Functions as an Base or Acid