 Acids are all around us.
– Foods like lemons, limes, and oranges contain acidic compounds.
– When wine is oxidized, it becomes acetic acid.
– Our muscles produce lactic acid.
– Our stomach uses hydrochloric acid to break down food.

 Arrhenius definition – an acid is a substance that produces
H + ion when added to water.
HCl
(aq)

H +
(aq)
+ Cl -
(aq)
 Acids taste sour, are electrolytes, and neutralize bases.

 Arrhenius definition – a base is a substance that produces
OH when added to water.
NaOH
(aq)

Na +
(aq)
+ OH -
(aq)
 Bases taste bitter, have a slippery, soapy feel, and neutralize acids.

 Expanded definition of acids & bases.
– Acid = a proton (H + ) donor
– Base = a proton acceptor
 Important note about the use of H + in equations.
 For weak acids and bases, the reactions are reversible, equilibria.

 HC
2
H
3
O
2
+ H
2
O

H
3
O + + C
2
H
3
O
2
-
 NH
3
+ H
2
O

NH
4
+ + OH -

 Write the conjugate base of an acid.
– Remove an H +
1. HBr – H + = Br
−
2. H
2
S – H + = HS −
 Write the conjugate acid of a base.
– Add an H +
1. NO
2
−
+ H + = HNO
2
2. NH
3
+ H + =
NH
4
+

1. The conjugate base of HCO
3
− is a. CO
3
2− b. HCO
3
− c. H
2
CO
3
2. The conjugate acid of HCO
3
− is a. CO
3
2− b. HCO
3
− c. H
2
CO
3
3. The conjugate base of H
2
O is a. OH
− b. H
2
O c. H
3
O +
4. The conjugate acid of H
2
O is a. OH
− b. H
2
O c. H
3
O +

 Strong acids completely ionize
(dissociate) in water.
 Strong acids are also strong electrolytes.
 There are six: HClO
4
, H
2
SO
4
, HI, HBr,
HCl, and HNO
3
.

 Weak acids only partially dissociate to produce ions in solution.
 Weak acids are weak electrolytes.
 Too many to list, but some common ones are: H
3
PO
4
, HC
2
H
3
O
2
, HF, and
HC
6
H
7
O
6
.

 Strong bases completely ionize in water and are also strong electrolytes.
 Only group 1A and 2A hydroxides are strong bases.
 All other metal hydroxides are insoluble in water.

 Weak bases only partially react with water (accepting a proton) and are, thus, weak electrolytes.
 Ammonia, NH
3 weak base.
, is the most common
 Lone pair on N group will accept a proton.
 Other organic weak bases.

 Water can act as both an acid and a base.
 Any substance that does this is called amphoteric.
 Pure water – two water molecules will occasionally react with each other where one is an acid and one is a base.

 H
2
O + H
2
O

H
3
O + + OH -
 [ X ] = symbol for molarity.
 In pure water, [H
3
O + ] = [OH ].
 K w
= [H
3
O + ] x [OH ]; where K w
= 1 E-14.
 Thus, [H
3
O + ] = [OH ] = ____________

3
+
-
 If we know [H
3
O + ], then [OH ] =
 If we know [OH ], then [H
3
O + ] =
 Ex) [H
3
O + ] = 2.5 E-5
 Ex) [OH ] = 4.8 E-3

 One convenient method for measuring the acidity or basicity of a solution is to use the pH scale.
 pH is a logarithmic (log) scale and equal to: pH = -log[H
3
O + ].
 pOH = -log[OH ].
 pH + pOH = 14.

 A word about significant figures.
 2.4
x 10 -3 M pH = 2 .
62
 Red numbers are the significant digits.
 Blue numbers are exact numbers.


1.
2.
3.
4.
Basic Calculators
Enter concentration
Press “log” key
Change the sign
Record answer to proper s.f.’s
3.

1.
2.
4.
TI-83 or TI-89
Press negative sign
Press “log” key
(select in catalog)
Enter concentration, close parenthesis
Enter key and round to proper s.f.’s

 To convert a pH back to a concentration, you will use the “antilog” key = 10 x .
 [H
3
O + ] = 10 -pH
 [OH ] = 10 -pOH
 Can also use the universal power (^) key.

 Can be done with a meter, pH paper, or an indicator.

[H
3
O + ] [OH ]
1.8 E-5 pH
3.72
3.4 E-2 pOH A/B
5.48

 An acid will react with most metals.
– Mg
(s)
+ 2 HCl
(aq)

MgCl
2(aq)
+ H
2(g)

 Acids react with any carbonate (CO
3
-2 ) and bicarbonate (HCO
3
) to generate CO
2
.

 Normally, rain is slightly acidic – pH of
5.5 to 6.2 – due to dissolved CO
2
.
 Burning fossil fuels, which contain small amounts of Sulfur, produces SO
3
.


This is converted to H
2
SO
4
.
H
2
SO
4
+ CaCO
3

CaSO
4
+ H
2
O + CO
2




Acids neutralize bases and vice versa.
Acid + Base

Water + Salt
HCl + NaOH

H
2
O + NaCl
 Balancing more complex acid-base neutralization.
– Each H + needs one OH .
– Each H + and OH makes one H
2
O.

 Acid-base titration – can use a known acid or base solution to determine an unknown counterpart.
 Endpoint – when all of the unknown acid or base has reacted.
 Indicator – a substance that changes color when it changes pH.

 Requires precise glassware to deliver the known solution = buret.
 Stopcock allows for delivery drop-by-drop.
 Volumes can be read to nearest 0.05mL.

 When a small amount of acid or base is added to pure water, the pH swings drastically.
 Some solutions, though, will resist these wild swings in pH and are called buffer solutions.
 Made from a ___________ and a salt containing the ______________.

 1.0L of pure water + 0.0100moles
(0.365g) of HCl.


[H
3
O + ] = 0.010 mol / 1.0L = 0.010M
pH =
 1.0L of pure water + 0.0100moles
(0.400g) of NaOH.
 [OH ] = 0.010 mol / 1.0L = 0.010M
 pH =

 Buffer of HF and NaF (note: Na + is a spectator ion).
 Ideal buffer would contain a 50 / 50 mixture of each.
 1.0L of 0.10 moles HF (2.0g) and 0.10 moles of NaF (4.2g) will have a pH of
3.17.

 HF + H
2
O

H
3
O + + F -
50% 50% (from NaF)
 Addition of strong acid
– reacts with F ion to generate more HF
 Addition of strong base
– reacts with HF to produce water plus more
F -

 Buffer plus 0.010 moles of HCl.
 pH = 3.08 (from 3.17).
 Buffer plus 0.010 moles of NaOH.
 pH = 3.26 (from 3.17).




1.
2.
Normal blood pH is 7.35 to 7.45.
Outside this range, cells cannot function properly.
Two buffer systems are present to maintain this pH.
H
2
CO
3
H
2
PO
4
-
/ HCO
3
-
/ HPO
4
-2