Introduction/ Concept of acid and base, Importance of acids and bases in Pharmacy, storage condition. Official acids: Phosphoric acid (Conc/dil), HCl (Conc/dil), Boric acid. Official Bases: NaOH, KOH, Ca (OH)2, dil. and strong NH3, Na2CO3, Acidosis and Alkalosis.

1. Unit 2: Acid, Base, Buffers and
Water (9 hrs)
-Kabin Maleku

2. Acid and base (3 hrs)
• Introduction/ Concept of acid and base,
Importance of acids and bases in Pharmacy,
storage condition. Official acids: Phosphoric
acid (Conc/dil), HCl (Conc/dil), Boric acid.
Official Bases: NaOH, KOH, Ca (OH)2, dil. and
strong NH3, Na2CO3, Acidosis and Alkalosis.

3. Definitions
• Acid
• Base

4. Acid
• An acid is a substance which, when dissolved in
water, releases protons.
• The extent of dissociation, that is, the amount of
protons released compared to the total amount
of compound, is a measure of the strength of the
acid.
• For example, HCl (hydrochloric acid) is a strong
acid, because it dissociates completely in water,
generating free [H+] and [Cl-].
• Acidity can be measured on a scale called pH
(more scarily, “the negative logarithm of the
hydrogen ion concentration”).

5. An acid is a compound that in
aqueous solution will readily:
A.Shed a proton.
B.Shed an electron.
C.Gain a proton.
D.Gain an electron.
E.None of the above.

6. 6
• Most living cells have a very
narrow range of tolerance for
pH, i.e. [H+].
• The [H+] concentration will be
important (either explicitly or
implicitly) for many other
topics in biology.
• [H+] is controlled in all
biological organisms, and in
virtually all biochemical
experiments.
• Each pH unit represents a
factor of 10 difference in [H+].
The pH scale goes from 0 to 14—because [H+][OH-] = 10-14
pH
10-2
10-3
10-5
10-4
10-8
10-7
10-6
[H+] M
10-10
10-9
10-11
10-12
10-13
10-14
10-1
100 A strong acid
A strong base
SOURCE: http://en.wikipedia.org/wiki/Image:PH_scale.png#file

7. CQ#2: In an aqueous solution where the
H+ concentration is 1 x 10-6 M, the OH-
concentration must be:
7
A. 14 x 10-6 M
B. 1 x 10-6 M
C. 1 x 10-7 M
D. 1 x 10-8 M
E. 14 x 10-8 M

8. The Conceptual Problem with pH
• Because it’s a logarithmic scale, it doesn’t make
“sense” to our brains.
• But Paul explains it well—every factor of 10
difference in [H+] represents 1.0 pH units, and
• Every factor of 2 difference in [H+] represents 0.3
pH units.
• Therefore, even numerically small differences in
pH, can have profound biological effects…
8

9. How Can You Actually
Determine the pH of a Solution?
• Use a pH meter—read the number.
• Use pH paper (color patterns indicate pH).
• Titrate the solution with precise amounts of
base or acid in conjunction with a soluble
dye, like phenolphthalein, whose color
changes when a specific pH is reached.
8

10. ACID
• 1. Lavoisier's knowledge of strong acids
was mainly restricted to oxyacids, which
tend to contain central atoms in high
oxidation states surrounded by oxygen,
such as HNO3 and H2SO4
• has at least one hydrogen atom bound to
oxygen; and forms an ion by the loss of
one or more protons

11. ACID
• According to Liebig, an acid is a hydrogen-
containing substance in which the
hydrogen could be replaced by a metal.
Liebig's definition, while completely
empirical, remained in use for almost 50
years until the adoption of the Arrhenius
definition.

12. ACID-BASE
• The Arrhenius definition of acid-base
reactions is a more simplified acid-base
concept devised by Svante Arrhenius
• As defined at the time of discovery, acid-base
reactions are characterized by Arrhenius
acids, which dissociate in aqueous solution
form hydrogen or the later-termed oxonium
(H3O+) ions, and Arrhenius bases which form
hydroxide (OH-) ions.
• 2NaOH + H2SO4 → 2 H2O + Na2SO4

13. ACID-BASE
• The Brønsted-Lowry definition, formulated independently by
its two proponents Johannes Nicolaus Brønsted and Martin
Lowry in 1923 is based upon the idea of protonation of bases
through the de-protonation of acids -- more commonly
referred to as the ability of acids to "donate" hydrogen ions
(H+) or protons to bases, which "accept" them.
• CH3COOH + H2O === CH3COO- + H3O+
• NH3 + H2O === NH4
+ + OH-
• CH3COO- + H2O === CH3COOH + OH-
• NH4
+ + H2O === NH3 + H3O+

14. Model Definition of
Acid
Definition of
Base
Arrhenius H+ producer OH- producer
Bronsted-Lowry H+ donor H+ acceptor
Lewis Electron-pair
acceptor
Electron-pair
donor

15. Examples
• Strong Acid
• Weak acid
• Weak base
• Strong Base

16. Importance
• pH in the Digestive System
• Mouth-pH around 7. Saliva contains
amylase, an enzyme which begins to break
carbohydrates into sugars.
• Stomach- pH around 2. Proteins are broken
down into amino acids by the enzyme pepsin.
• Small intestine-pH around 8. Most digestion
ends. Small molecules move to bloodstream
toward cells that use them

17. Storage
• Segregate acids from bases
• Segregate incompatible classes of dangerous
goods
• Store acids and bases in a compliant corrosive
resistant storage cabinet
• Protective equipment (goggles, gloves, lab
coat or apron, chemically-resistant clothing)

18. Buffers
• Types
• Properties
• pH of buffers

19. - Decide on the Buffer Properties
• Before making a buffer you must know;
1. what molarity you want it to be
2. what volume to make
3. what the desired pH is.
• Most buffers work best at concentrations
between 0.1 M and 10 M.
• The pH should be within 1 pH unit of the acid/
conjugate base pKa.

20. Requirements of buffers
• Solubility
• Permeability
• Ionic strength
• Dependence of the pKa value
• Inert substances
• UV absorption
• Purity – simple method of manufacture
• Costs