1. INTRODUCTION TO
PHARMACY AND
CHEMISTRY OF
INORGANIC MEDICINALS
WITH QUALITATIVE
CHEMISTRYperiodic table.jpg

2. What is chemical kinetics?
• is the study of reaction rates and the
mechanism or step sequence under
which reactions occur.
• The rate can be expressed as:
rate = change in substance/time
for change to occur

3. Chemical Kinetics
• The rate of a reaction is the amount of
chemical change, which takes place in a given
interval of time.
• It is measured in terms of products formed
per unit time at a given temperature
• It may be measured in terms of disappearance
of reactants per unit time at a given
temperature.

4. Chemical Kinetics
• Few chemical reactions occur in a single step.
There are usually one or more intermediate
products, which form, and the stepwise
sequence of reactions is known as the
reaction mechanism.

5. Chemical Kinetics
• Reaction rates are of great importance. The
feasibility of a chemical reaction, particularly a
commercial chemical process, often depends
on the reaction rate. If the process is too slow,
it may not be suitable for commercial
exploitation.

6. MOLECULAR COLLISION THEORY
• A chemical change takes place as a result of
collision of molecules. The greater the number
of collision per unit time, the greater the
conversion of initial substances into products
per unit time -- that is, the greater the speed
of reaction.
• However, not all collisions of molecules
necessarily result in chemical change.

7. with concentration phenomenon explained
by collision theory

8. Two conditions must be met in order to
have an effective collision.
• The orientation of the colliding molecules
must be favorable for the making and
breaking of bonds, and

9. • Before molecules can react, they must possess
a certain minimum energy, termed activation
energy, Only molecules with the necessary
activation energy will undergo reaction.

10. FACTORS WHICH INFLUENCE THE
SPEED OF CHEMICAL REACTIONS
1. Nature of the Reacting Substances
2. Physical state
3. Concentration
4. Temperature
5. Catalysts
6. Pressure

11. Nature of the Reacting Substances
• Substances differing activity and hence in the
speed 'with which they react with other
substances.
• The active metals displace hydrogen
vigorously and rapidly from acids, while the
less active metals act slowly, if at all. Metals
differ in their rates of corrosion because of
differences in speed of combination with
oxygen and other elements.

12. • Depending upon what substances are
reacting, the reaction rate varies. Acid/base
reactions, the formation of salts, and ion
exchange are fast reactions.
• When covalent bond formation takes place
between the molecules and when large
molecules are formed, the reactions tend
to be very slow. Nature and strength of
bonds in reactant molecules greatly
influence the rate of its transformation into
products.

13. Physical state
• The physical state (solid, liquid, or gas) of a
reactant is also an important factor of the
rate of change.
• This means that the more finely divided a
solid or liquid reactant the greater
its surface area per unit volume and the
more contact it makes with the other
reactant, thus the faster the reaction.

14. Concentration
• Concentration plays a very important role
in reactions, because, according to
the collision theory of chemical reactions,
molecules must collide in order to react
together.
• As the concentration of the reactants
increases, the frequency of the molecules
colliding increases, striking each other
more frequently by being in closer contact
at any given point in time.

15. Temperature
 Reacting molecules or atoms must collide before
they can react.
 At a high temperature, the molecules or atoms
move faster and therefore should come in
contact more often. This is especially true in
gaseous reactions. In general, the speed of a
chemical change is approximately doubled or
tripled for every ten degree rise in temperature
regardless of whether the reaction is exothermic
or endothermic.

16. Catalysts
• A catalyst is a substance that alters the rate of
a chemical reaction without being used up in
the reaction. the catalyst may be recovered
unchanged at the conclusion of the process.

17. Catalysts
• Two types of catalysts:
–accelerators
–inhibitors

18. Catalysts
• Accelerators or positive catalysts speed up a
chemical reaction
• Inhibitors or negative catalysts slow down a
reaction and are used to control undesired
reactions.

19. CHEMICAL REACTION
It entails the removal of valence
electrons, adding electrons to a partly
filled valence shell, or sharing a pair of
electrons between two atoms.

20. Chemical Equilibrium
• Chemical equilibrium applies to reactions
that can occur in both directions. In a
reaction such as:
CH4(g) + H2O(g) <--> CO(g) + 3H2(g)
• The reaction can happen both ways. So
after some of the products are created the
products begin to react to form the
reactants

21. • When the net change of the products and
reactants is zero the reaction has reached
equilibrium.
Equilibrium Constant
• To determine the amount of each compound
that will be present at equilibrium you must
know the equilibrium constant. Consider the
generic equation:
aA + bB <--> cC + dD
•

22. • The upper case letters are the molar
concentrations of the reactants and products.
The lower case letters are the coefficients that
balance the equation.
• Use the following equation to determine the
equilibrium constant (Kc).

23. Law of Mass Action
• The law stating that the rate of any given
chemical reaction is proportional to the
product of the activities (or concentrations) of
the reactants.

24. • A reversible reaction is a chemical
reaction where the reactants form products
that, in turn, react together to give the
reactants back.
Reversible reactions will reach an equilibrium
point where the concentrations of the
reactants and products will no longer change.

25. • A reversible reaction is denoted by a
double arrow pointing both directions in a
chemical equation. For example, a two
reagent, two product equation would be
written as
A + B ↔ C + D

26. Le Chatelier’s principle
• If a chemical system at equilibrium
experiences a change
in concentration, temperature, vol
ume, or partial pressure, then the
equilibrium shifts to counteract the
imposed change and a new equilibrium is
established.

27. pH and pOH
• The symbol “p” means “the negative of the
logarithm of”
• Water undergoes auto- or self-ionization as
shown in the following equation
2H2O H3O + OH
An equilibrium is established between the ions
produced and the unionized water.
Kw = [H3O ] [OH ]
+ -
+ -

28. • Kw is used to represent equilibrium constant
for the ionization of water. It is a special
notation of the general representation Kc.
• The value of Kw when measured at 25˚C has
been determined to be 1x10
pH + pOH = pKw
pH + pOH = 14
-14

29. • pH is a measure of the hydrogen ion
concentration, [H+]
• pH is calculated using the following formula:
pH = -log10[H+]
• pOH is a measure of the hydroxide ion
concentration, [OH-]
pOH is calculated using the following formula:
pOH = -log10[OH-]

30. pKa and pKb
pKa = -log (Ka)
Ka is the equilibrium constant for the ionization of
acids.
pKb = -log(Kb)
Kb is the equilibrium constant for the ionization of
bases.

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