(courtesy of L. Scheffler, Lincoln High School, 2010)
Acids and
Bases
1
Acids
•
React with certain metals to produce
hydrogen gas.
React with carbonates and bicarbonates
to produce carbon dioxide gas
•
Bases
•
•
•
Have a bitter taste
Feel slippery.
Many soaps contain bases.
2
Properties of Acids
 Produce H+ (as H3O+) ions in water (the hydronium ion
is a hydrogen ion attached to a water molecule)
 Taste sour
 Corrode metals
 Good Electrolytes
 React with bases to form a salt and water
 pH is less than 7
 Turns blue litmus paper to red “Blue to Red A-CID”
3
Some Common Acids
HC2H3O2
acetic acid
in vinegar
HCl
hydrochloric acid
stomach acid
H3C6H5O7
citric acid
fruits
H2CO2
carbonic acid
soft drinks
H32PO4
phosphoric acid
soft drinks
Properties of Bases
 Generally produce OH- ions in water
 Taste bitter, chalky
 Are electrolytes
 Feel soapy, slippery
 React with acids to form salts and water
 pH greater than 7
 Turns red litmus paper to blue
“Basic Blue”
5
Some Common Bases
NaOH
sodium hydroxide
lye
KOH
potassium hydroxide
liquid soap
Ba(OH)2
barium hydroxide
stabilizer for plastics
Mg(OH)2
magnesium hydroxide “MOM” Milk of magnesia
Al(OH)3
aluminum hydroxide
Maalox (antacid)
Arrhenius Definition
Arrhenius
Acid - Substances in water that increase the
concentration of hydrogen ions ((H+ or
hydronium ions H3O+).
Base - Substances in water that increase
concentration of hydroxide ions (OH-).
Categorical definition – easy to sort substances
into acids and bases
Problem – many bases do not actually contain
hydroxides
7
Bronsted-Lowry Definition
Acid - substance that donates a proton.
Base - substance that accepts a proton.
HA + B
Ex
HCl + H2O
Acid


Base


HB+
+
H3O+ +
Conj Acid
A-
ClConj Base
A “proton” is really just a hydrogen atom that has lost
it’s electron!
The classification depends on how the
substance behaves in a chemical reaction
8
Conjugate Acid Base Pairs
Conjugate Base - The species remaining after an acid
has transferred its proton.
Conjugate Acid - The species produced after base
has accepted a proton.
HA & A- - conjugate acid/base pair
A- - conjugate base of acid HA
B & HB+ - conjugate acid/base pair
HB+ - conjugate acid of base :B
9
A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
base
acid
conjugate
acid
conjugate
base
Examples of BronstedLowry Acid Base Systems
Note: Water can act as acid or base
Acid
Conjugate Acid
Base
Conjugate Base
+
H2O

H3O+ +
Cl-
H2PO4- +
H2O

H3O+
+
HPO42-
NH4+
H2O 
H3O+
+
NH3
HCl
+

11
Lewis Definition
Lewis
Acid - an electron pair acceptor
Base - an electron pair donor
12
Brønsted-Lowry vs.
Lewis

All B/L bases are Lewis bases BUT,
by definition, a B/L base cannot
donate its electrons to anything but a
proton (H+)

While B/L is most useful for our
purposes, Lewis allows us to treat a
wider variety of reactions (even if no
H+ transfer occurs) as A/B reactions
Acid Strength
Strong Acid
- Transfers all of its protons to water;
- Completely ionized;
- Strong electrolyte;
- The conjugate base is weaker and has a
negligible tendency to be protonated.
Weak Acid
- Transfers only a fraction of its protons to
water;
- Partly ionized;
- Weak electrolyte;
- The conjugate base is stronger, readily
accepting protons from water
 As acid strength decreases, base strength increases.
 The stronger the acid, the weaker its conjugate base
 The weaker the acid, the stronger its conjugate base
14
Acid Dissociation Constants
Dissociation constants for some weak acids
15
Base Strength
Strong Base - all molecules accept a proton;
- completely ionizes;
- strong electrolyte;
- conjugate acid is very weak, negligible
tendency to donate protons.
Weak Base
- fraction of molecules accept proton;
- partly ionized;
- weak electrolyte;
- the conjugate acid is stronger. It more
readily donates protons.
 As base strength decreases, acid strength increases.
 The stronger the base, the weaker its conjugate acid.
 The weaker the base the stronger its conjugate acid.
16
Common Strong Acids/Bases
Strong Acids
Strong Bases
Hydrochloric Acid
Sodium Hydroxide
Nitric Acid
Potassium Hydroxide
Sulfuric Acid
*Barium Hydroxide
Perchloric Acid
*Calcium Hydroxide
*While strong bases they are
not very soluble
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A/B Behavior & Chemical
Structure
Binary Acids
1.
•
Hydrogen and another element
Polyprotic Acids
2.
•
Have more than 1 Hydrogen to give away
Oxyacids
3.
•
have O in compound
Carboxylic Acids
4.
•
have –COOH in compound
Wait, water can go both
ways?
amphoteric substances can behave
as either an acid or base depending
on what they react with.
 water and anions with protons (H+)
attached are the most common
amphoteric substances

Autoionization of Water
H 2O + H 2O
OH- + H3O+
O H-
H3O+
@ 25 oC the concentrations for both
[H3O+] and [OH-] = 1.00 x 10-7 and
[H3O+] [OH-] = 1.00 x 10-14 = Kw
Since [H3O+] [OH-] = 1.00 x 10-14 = Kw
when [H3O+]=[OH-] the solution is neutral
 when [H3O+]>[OH-] the solution is acidic
 when [H3O+]<[OH-] the solution is basic

The pH scale is a way of
expressing the strength of
acids and bases. Instead of
using very small numbers,
we just use the NEGATIVE
power of 10 on the Molarity
of the H+ (or OH-) ion.
Under 7 = acid
7 = neutral
Over 7 = base
pH of Common Substances
pH calculations – Solving for H+
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both
sides and get
10-pH = [H+]
[H+] = 10-3.12 = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or “2nd function”
and then the log button
Calculating the pH
pH = - log [H+]
(Remember that the [ ] mean Molarity)
Example: If [H+] = 1 X 10-10
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example: If [H+] = 1.8 X 10-5
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74
pH calculations – Solving for H+

A solution has a pH of 8.5. What is the
Molarity of hydrogen ions in the solution?
pH = - log [H+]
8.5 = - log [H+]
-8.5 = log [H+]
Antilog -8.5 = antilog (log [H+])
10-8.5 = [H+]
3 X 10-9 = [H+]
pOH

Since acids and bases are opposites,
pH and pOH are opposites!

pOH does not really exist, but it is
useful for changing bases to pH.

pOH looks at the perspective of a base
pOH = - log [OH-]

Since pH and pOH are on opposite
ends
pH + pOH = 14
The pH Scale
pH
[H3O+ ]
[OH- ]
pOH
28
pH testing

There are several ways to test pH
 Blue litmus paper (red = acid)
 Red litmus paper (blue = basic)
 pH paper (multi-colored)
 pH meter (7 is neutral, <7 acid, >7
base)
 Universal indicator (multi-colored)
 Indicators like phenolphthalein
 Natural indicators like red cabbage,
radishes
pH indicators




Indicators are dyes that can be
added that will change color in the
presence of an acid or base.
Some indicators only work in a
specific range of pH
Once the drops are added, the
sample is ruined
Some dyes are natural, like radish
skin or red cabbage
Indicators
31
pH and acidity
The pH values of several
common substances are
shown at the right.
Many common foods are
weak acids
Some medicines and many
household cleaners are
bases.
32
Neutralization
An acid will neutralize a base,
giving a salt and water as products
Examples
Acid
HCl
H2SO4
H3PO4
2 HCl
Base
+ NaOH
+ 2 NaOH
+ 3 KOH
+ Ca(OH) 2
Salt




NaCl
Na2SO4
K3PO4
CaCl2
water
+ H2O
+ 2 H2O
+ 3 H2O
+ 2 H 2O
A salt is an ionic compound that is
formed from the positive ion
(cation) of the base and the
negative ion (anion) of the acid
34
Titration & Titration Curves

Titration: the adding of one solution of an
known concentration into another solution


standard solution: a solution with a known
concentration
Titration curve: a graph showing pH vs
volume of acid or base added


The pH shows a sudden change near the
equivalence point
The Equivalence point (a.k.a. stoichiometric
point) is the point at which the moles of OHare equal to the moles of H3O+
35
Buffer Solutions - Characteristics




A solution that resists a change in pH.
It is pH stable.
A weak acid and its conjugate base
form an acid buffer.
A weak base and its conjugate acid
form a base buffer.
We can make a buffer of any pH by
varying the concentrations of the
acid/base and its conjugate.
37