 Water molecules can dissociate and ionize when a hydrogen atom shared by
two water molecules in a hydrogen bond shifts from one molecule to the other.
The electron of the hydrogen atom is not transferred, only a single proton. This
single proton is referred to as a hydrogen ion (H+), with a charge of +1. the
water molecule that lost the proton becomes a hydroxide ion (OH-) with a
charge of -1. the water molecule bonding to the hydrogen ion becomes a
hydronium ion (H3O+) with a charge of +1. The chemical equation is written:
2 H2O
H3O+ + OH –
To simplify the equation, by convention we write:
H2O
H + + OH -
2
 A solution that has equal concentrations of H+ and OH- is said to be
neutral.
 An acid is a substance that dissociates when dissolved in water,
producing H+. Hydrochloric acid is a strong acid.
HCl
H+ + Cl-
 A base is a substance that, when dissolved in water, reduces the
hydrogen ion concentration. Some bases do this by donating OH- that
interact with H+ to form water. Sodium hydroxide is an example of a
strong base.

NaOH
Na+ + OH-
3
 A measure of the power of the hydrogen ion concentration in a solution
is known as pH.
pH= - log₁₀ [H+]
 The pH scale ranges from 0-14 for biological systems. At neutrality, the
hydrogen and hydroxide ion concentration is 10⁻⁷ M. Therefore, at
neutrality pH is equal to 7. There is an inverse relationship between
pH and hydrogen ion concentration; that is, as the hydrogen ion
concentration increases, the pH decreases, and vice versa.
4
 Base: dissociates hydroxide ions into
solution or accepts hydrogen ions
pH above 7
 Acid: dissociates hydrogen ions into
solution
pH below 7
5
 It is very important for living organisms to maintain a pH around 7.2-
7.4.
 Buffers are special solutions that can help maintain pH by resisting any
changes to pH caused by the addition of an acid/base.
 Buffers are solutions of weak acid/weak base. An example is carbonic
acid/bicarbonate in your blood.
H₂CO3 + H₂O ↔ H3O⁺ + HCO3⁻
 H2CO3 (carbonic acid) is able to act as an acid by donating H+ when a
base is added and HCO3 (bicarbonate) is able to act as a base by
accepting H+ when acid is added.
 The more the resistance to a change in pH the better the buffering
capacity of the solution.
6
 Anthocyanins are plant pigments that can be used as a pH indicator
because they change color in response to changes in pH. These
pigments are responsible for red, blue, and purple colors that are found
in flowers, fruits, and autumn leaves.
 This is an example of the colors obtained for a few pHs.
pH 2
pH3
pH5
pH7
pH8 pH12
7
pH reading
 It has a greater accuracy.
 Before a pH meter can be used, it must be calibrated.
 For best results, calibration should be performed with both
pH 7 and pH 4 and it should be done every couple of hours.
 The pHs obtained in class were as follows:
Solution
pH
Distilled water
6.53
Apple juice
4.56
Vinegar
2.90
Soft drink
3.4
Coffee
5.02
Detergent water
8.10
Milk
6.72
8
 The pH paper is orange in color
before use and changes color
depending on the pH being
tested.
The pHs obtained in class were as follows:
Solution
pH
Distilled water
7
Apple juice
5
Vinegar
3
Soft drink
3
Coffee
5
Detergent water
8
Milk
7
9
 Determines how good a buffer is in the presence of an
acid or a base.
 You will record the pH of two solutions (A and B) as
their buffering capacities are tested with the addition
of acid and then base. By examining your results you
will determine which of the two solutions is the betterbuffered solution.
 Initial pH of solution A= 7
 Initial pH of solution B = 7
10
pH of Solution A
1
2
3
HCl
7
6.9
6.87 6.82 6.8
6.75 6.59 6.47 6.2
NaOH
7
7.01
7.03 7.04 7.10
7.30 7.35
ml of acid/base added
4
5
6
7
8
9
10
6
7.60 7.90 8
pH of Solution B
ml of acid/base added
HCl
NaOH
1
2
2.39 2.35
10.2
3
4
5
2.29 2.20 2.01
10.31 10.4
10.5
10.6
6
7
8
9
10
1.99
1.96
1.92
1.80
1.73
10.7
10.8
10.9
11.01 11.3
11
pH
0
End
12