1. Acids and bases
Chapter 19

2. i. Arrhenius definition
Acids increase the concentration of hydrogen ions in aqueous solution
Acids have more H+ ions than OH- ions
Formulas for acid begin with H
Bases increase concentration of hydroxide ions (OH-) in aqueous solutions
Bases have more hydroxide ions than hydrogen ions
Formulas for bases end with –OH-

3. ii. Properties
Acids
Bases
Taste sour
Taste bitter, feel slippery
Reacts with metals to form H2(g)
Litmus – blue
Corrosive
Phenolphthalein - pink
Litmus – red
Phenolphthalein – colorless

4. iii. Bronsted-lowry acids/bases
Acids are hydrogen ion donors
Bases accept hydrogen ions
Conjugate acid – species formed after the base has taken a H+
Conjugate base – species formed after the acid has lost a H+
A base is a proton (hydrogen ion) acceptor, so the conjugate acid is the acid in the reverse reaction/process
Example: Base = NO3-, Conjugate Acid = HNO3
Example = HCl, Conjugate Base = Cl-
Polyprotic acids – can donate more than one H+
Example: H3PO4 (in reality, only one H is given up)

5. iv. Strong vs. Weak acids/bases
Strong acids: completely ionize in water, strong electrolyte
Strong Bases: same definition as strong acids
Examples: HNO3, HCl, HI, HBr, H2SO4
Group 1 and 2 hydroxides (NaOH, KOH, Sr(OH)2, etc)
Weak acids: do not ionize completely, some whole molecules remain, weak electrolytes
Weak bases: same as weak acids
NH3
Examples: HC2H3O2, HF

6. v. Neutralization reactions
Acid + Base  Salt + water
Salt – ionic compounds (except oxides)
Salts can be
Neutral – when both the acid and base are strong
Acidic – when a strong acid reacts with a weak base
Basic – when a strong base reacts with a weak acid

7. pH – measures the concentration of H+ ions in solution
vi. Ph and poh
pH – measures the concentration of H+ ions in solution
Calculate: pH = -log[H+]
pOH – measures the concentration of OH- ions in solution
Calculate: pOH = -log[OH-]

8. c) Self-ionization of water
vi. Ph and poh
c) Self-ionization of water
Amphoterism – water can act as either an acid or a base
H2O + H2O  H3O+ + OH-
H2O + NH3
H2O + HCl
d) In one liter of pure water at 25oC
i. the concentration of H+ = 1x10-7
ii. the concentration of OH- = 1x10-7

9. vi. pH and poh
e) Formulas useful for pH and pOH problems (these are in your reference tables) i. pH = -log[H+] ii. pOH = -log[OH-] iii. pH + pOH = 14 iv. [H+] = 10-pH vi.[OH-] = 10-pOH vii. Kw = [H+][OH-] = 1x10-14

10. vi. Ph and poh
f) Below 7 = acidic – more H+ ions exactly 7 = neutral – equal H+ and OH- Above 7 = basic – more OH- ions

11. standard solution
unknown solution
VII. Titration
Concept: analytical method used to determine the concentration of an unknown substance. Often an acid-base reaction.
Equivalence point: point at which moles of H+ = moles of OH-; can use this fact (and the volumes involved) to calculate the molarity of the unknown (similar to dilution formula)
The equivalence point can be determined by measuring the pH at regular intervals and graphing the data or
Estimate the equivalence point by using an indicator

12. c) End Point: the point at which an indicator changes color
vii. titration
c) End Point: the point at which an indicator changes color
This will be close to (but not exactly) the equivalence point. For example – phenolphthalein changes from colorless to pink at a pH of about 8. If you are doing a strong acid/strong base titration, the equivalence point is at a pH of 7. The difference is only a mL or so, so it gives a very close approximation of the equivalence point

13. vii. titration

16. VII. titration
d) Titration Calculation Formula MaVana = MbVbnb M = molarity V = volume n = number of moles of OH- (for the base) or H+ (for the acid) -use the subscript a refers to the acid, b refers to the base

17. vii. titration
d. Titration Calculation i. example: 42.5mL of 1.3M KOH are required to neutralize 50.0mL of H2SO4. Find the molarity of the sulfuric acid