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Wednesday, 23 December 2015

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 7

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 7



Buffer Action:

The ability of a buffer to resist change in its PH, even after the addition of small amount of strong acid or base is known as buffer action.
Now the question arises that how a buffer solution resists change in its PH, even with the addition of small amount of a strong acid or base.
It can be explained as: consider the CH3COOH / CH3COOna solution: if we acid small amount of a strong acid like HCL the PH doesn’t change. It is between of the fact that in CH3COOH / CH3COONa : CH3 COOH ionizes to large extent : so there are CH3COO- in large excess which come from CH3COONa . Here common ion effect of CH3COOH and HCL occurs because both produce H+. As a result the solubility of CH3COOH is further decreased. Thus most of H+ ion of HCL ( which are responsible for change of PH ) are converted into in weak acid, CH3COOH which is already in excess In buffer solution.
Thus the PH of buffer remains constant because H+ ions of HCL don’t remain in solution.

Buffer Capacity:

The extent till which a buffer can resist change in its PH is known as buffer capacity. A buffer can resist change in its Ph if we add particular amount of strong acid or base. If we exceed the amount of strong acid or base them the ph of ph will certainly change. So we can say  that there is certain limit till which a buffer can resist change in its ph. This limit is known as buffer  as buffer capacity.

Calculation of PH of Buffer Solution:

The ph of a buffer solution is calculated by an equation known as "Hander son Herelback equation" This equation is
Ph = Pka + log [ base ]-------------[acid ]

Application of Buffer Solutions:

  1. Buffer solutions have tremendous applications. Some important applications of Buffer Solutions are .The human blood is a buffer whose then serious problems may happen. In spite of having so many acidic and basic substances: the ph doesn’t changes.
  2. Buffer solutions are mostly used in Analytical laboratories for carrying out particular reactions at particular ph.
  3. Buffer solutions are used are used in various industries.
  4. Buffer solutions are used in Biological laboratories etc.

Written by: Asad Hussain

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 6

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 6


Auto-Ionization or Self-Ionization of Water:

Pure water is a very weak conclusion of electricity but the fact is that , water can conduct electric current, so there should be changed particles ( ions) in water.
The question arises that where do these ions come from?
This problem can be solved with the help of Lowery-Brorsted concept. Actually a water molecule releases a hydrogen ion (4+) ie proton and acid as an acid while this H+ is accepted by another H2O ion and thus this H2O molecule acts as a base (proton acceptor) this it is clear that H2O acts as an acid as well as, as a base, therefore we say H2O is amphoteric in nature

PH Scale:

In 1909, Soreson constructed a scale. Which is known as PH scale with the help of ph Scale the strength an acid or base is determined.
The term ph is refered as power of the hydrogen ions. The values of ph are from "O" to "14" The ph of acids is below "7". The stronger the acid the smaller will be its ph values and vice versa. Similarly the larger the value of ph a base the stronger will be that base and vice versa. The strength of bases can be determined by the tern POH. The values of POH also range from O to 14 but are in the reverse direction of values of ph

PH = 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14
POH 14 13 12 11 10 9 8 7 6 5 4 3 2 1 0

The strength of an acid is determined by determining its ph. The ph of an acid is the negative log of hydrogen ion (H+) concentration Ie
PH = -log [H+]

The strength of a base id determined by determining its POH. The POH of base is the negative log of  hydroxide ion concentration ie
Ph = -log [H+]

It is important to note that sum of the pH and pOH of a solution is always equal to 14
Ie pH + pOH = 14

The above relationship can be proved as consider the auto- ionization of water ie
2H2O ---------------> H3O + OH(+)
Here, the ionization of H2O is smaller that no considerable change occurs in the concentration of water and it still remains constant.

Solubility Product:

The term solubility product is used for the sparingly soluble salts in water. It can also be called as solubility product constant.
It can be defined as:
The solubility product or solubility product constant of sparingly soluble salt is defined as the product of molar concentration of its ion raised to the power of their stoichiometric co-efficient in the equilibrium equation:

Buffer Solutions:

The solution which resists change in its PH, even with the addition of small amount of a strong acid or a strong base is known as buffer solution.
A buffer solution is made as
1). When a weak acid ( like CH3COOH ) is mixed with its salt made by it with a strong base ( ie:  CH3COONa). They are mixed in a particular ration. This mixture is known as buffer and the solution of this mixture is known as buffer solution. Such a buffer is known as Acidic buffer. Ie CH3COOH/ CH3COONa buffer
2). When a weak base ( like NH4OH ) is mixed with its salt made by it with a strong acid ( ie NH4CL ) in affixed ratio and then dissolved in water. We get a buffer solution such a buffer is known as basic buffer
ie NH4OH / NH4CL . buffer The buffer solution has a particular PH range and this range remains constant even for long time or even after the addition of small amount of strong and or strong base.

Written by: Asad Hussain

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 5

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 5



Applications of Le- Chatelier's Principle : (Effect of Change in Conditions upon Equilibrium)

Le – Chatelier's principle is used to predict, how the variable like concentration. Temp, pressure etc, affect the position of equilibrium.
This principle also has a valuable application in prediction the condition for maximum yield of a particular in a reversible reaction.
1) Effect of Concentration:
According to Le – Chatelier's principle an increase in the concentration of any the reactants, shifts the equilibrium to the right i.e forward reaction increases similarly an increase in the concentration of any of the products shifts the equilibrium to the left i.e reverse reaction increases Same examples are as follow

N2 + 3H2--------------------2NH3
The addition of H2 and N2 would shift the equilibrium to the right. Thus the use of excess of N2 yield of NH3

FeCl3+n3NH4CNS --------------------Fe ( CNS)3 + 3NH4C ff2
Yellow colorless blood red colorless The addition of Fecl3 , shifts the equilibrium to the right & thus more Fe (CNS)3 (blood red) is formed , forward reaction increases. Similarly the addition of NH4Cl , shift the equilibrium to the left resulting in the formation of more NH4CNs & FeCl3 i.e reverse reaction is increased.

2) Effect OF Pressure:
In case of reactions ( reversible ) in which the total volume of reactants is equal to the total volume of the products , the equilibrium is not affected by change in pressure.
e.g

1. CO(g) + H2O(g) ---------------- CO2(g) + H2(g)
    1 vol       1 vol                                 1vol        1vol
Total     2 vol                                             2 vol

2. N2 (g) + O2(g) --------------- 2NO(g)
     1vol      1 vol                             2 vol
Total   2 vol                                    2 vol
Change of pressure has no effect upon the equilibrium in the above mentioned equilibrium system between total volume of reactants is equal to total volume of products. In case of reactions ( reversible ) where the total volume of reactants is not equal to the volume of the products, with an increase of pressure the equilibrium system shifts towards the smaller volume.
In case of these reversible reactions in which the volume of reactants is greater than that of the products, the equilibrium shifts to the right side i.e more products are formed by increasing the pressure of the system.
In case of those reversible reactions in which the volume of products is greater than that of reactants increases with increase of pressure.

3) Effect of Temperature:
There are two types of reversible reactions i.e endothermic and exothermic.
In case of endothermic reversible reactions, heat is absorbed in the forward direction and the same is evolved in the reverse direction and the same is absorbed in reverse direction. With increase of temperature of an reversible reaction ( endothermic or exothermic ) at equilibrium, the equilibrium always shifts in that direction where the temperature decreases. This in case of endothermic reaction (reversible) the equilibrium with crease of temperature . shifts to the forward direction
Examples :
Reversible endothermic Reactions
N2+O2+Heat ………………..2NO
H2S+O2+Heat ………………….H2 + S
As these are endothermic breakfronts thus in decease of temperature Shifts The equilibrium in forward direction . As the reaction are exothermic thus an increase in temperature will shift the equilibrium in
reverse direction. Thus more amount of NH3& SO2 can be formed by lowering the temperature.

4) Effect of Catalyst:
According to Le – Chatelier's principle, catalyst has no effect upon the equilibrium except the equilibrium reaches earlier in the presence of a catalyst. Actually the catalyst increases the speed of both forward & reverse reactions of a reversible reaction to the some extent, therefore the equilibrium remains unchanged.  However it is reached earlier.

Written by: Asad Hussain

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 4

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 4



Prediction of Direction of Reaction:

Let we have a reversible reaction for which value of Kc is 4 the reaction is
A+B ------------------ C+D
Let at a particular time interval, their molar concentrations are (a), (b), (c) and (d) respectively. Here we are not sure whether these concentrations are the equilibrium concentration. If the provided  concentrations of reactants and products are b/f equilibrium, then the reaction may be in forward or in reverse direction.
The direction of reaction can decided easily by taking the ration of product of concentration of products and product of concentration of reactants
Now
If Qc = Kc
Then the reaction is at equilibrium

If Qc > Kc
Then the reaction is in the reverse direction.

If Q> Kc
Then the reaction is in the forward direction.

Prediction of Extent of Reaction.

If we have a reversible reaction i.e
A+B -----------------------C+D
Let their equilibrium concentrations be (a), (b), (c) and (d). Then equilibrium constant Kc will be
Kc = (c) (d)

Now there are several possibilities :
  1. If the equilibrium concentration of C& D are very large as compared to those of A & B, them the value of Kc will be large enough i.e Kc >> 1. It shows that the reactants, react very well and thus produce large quantity of produces. But on the other hand, the products do not react well. Thus at equilibrium most of the concentration of reactants is converted into the products.
  2. If the concentration of C & D is very small as compared to that of reactants A & B, at equilibrium, then the value of Kc will very very small, i.e Kc<< 1 If shows that the reactants do not react well to form the products in enough quantity but the products react well to reform the reactants.
  3. If at equilibrium, the concentration of the reactants (A,B) and products (C,D) are almost equal, then value of kc will be nearly or exactly 1. If shows that both the reactants and products react well.

Calculation of Equilibrium Concentration:

Le-chatelier's  Principle:

To the study disturbance of a system at equilibrium, by changing its conditions like temperature,
pressure , concentration or presence of catalyst, a French chemist Henry Le Chatelier (1888)
presented a principle which is known as Le-chatelier's . It states that : If a system at equilibrium is
subjected t disturbance by changing it conditions like concentration the system shifts in such a direction to re-establish a new equilibrium.

Written by: Asad Hussain

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 3

FSc Notes Chemistry Part 1 Chapter 8 Chemical Equilibrium Lecture 3


Law of  Mass Action & Equilibrium Constant Expression:

In 1864, two Norwegian chemist, Goldberg and Peter, presented a law for the rate of a chemical reaction which is known as law of Mass Action.
This law states that the rate of a chemical reaction is directly proportional to the product of the concentration (or active masses) of reading substances.
The term active mass represents the concentration (in mole* dm-3) of reactants and products. Let we have a reversible reaction, i.e
A+B----------> kr C+D
kr
The equilibrium concentration of A, B, C and D are represented in square brackets i.e (A), (B), (C) and (D) are expressed in moles xdm-3.
According to law of Mass Action, the rate of forward reaction, is proportional to the product of molar  concentration of A and B.
Rate of forward reaction (rf) * (A) (B)
Or rf  = kf (A) (B) --------I
Here Kf is the constant of proportionality and is known as “rate constant for forward reaction”
Similarly the rate of reverse reaction is,
Rate of reverse reaction (*r)----- (C) (D)
Or Rr = Kr (C) (D)-------II
Here Kr is the constant of proportionality and is known as “rate constant for reverse reaction”
As the reaction , Under consideration is a reversible reaction, therefore at equilibrium Rf = Rr.

Kc is known as equilibrium constant in terms of concentration of reactants and products.


Equilibrium Content Expression. For Heterogeneous Equilibrium:

A heterogeneous equilibrium is an equilibrium in which the reactants and products are in more than one physical state. While writing the equilibrium constant expression for a heterogeneous equilibrium, the concentration for pure solids and liquids are neglected i.e. they are not indicated, It is because of the fact, that , the concentration of a pure solid or pure liquid is constant at a constant temperature and does not  depend upon the quantity of the substance.
For example the molar concentration of copper (at 20 c) is the same, whether we have 1g or 1 ton of copper.
3Fe (s) + 4H2O(g)---------------> Fe3O4(s) + 4H2(g)
The equilibrium constant expression for the above system is ,
Kc = Concentration of Fe & Fe3 O4 are neglected because they are in their pure solid state and their concentrations do not undergo appreciable change.


Significance or Applications of Kc:

Kc ( equilibrium constant ) has the following applications,
  1. Prediction of Direction of Reaction :
  2. Prediction of extent of reaction.
  3. Calculation of equilibrium concentration.

Written by: Asad Hussain