Relative lowering of vapour pressure

Relative lowering of vapour pressure

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• Vapour pressure of a solvent present in solution is less than the vapour pressure of the pure solvent.
• According to Raoult lowering of vapour pressure of a solution depends only on the concentration of the solute particles and remains independent of their identity.
• p1 =x1 p10
• The reduction in the vapour pressure of solvent (p1) is given as:

As x2=1-x1 the above equation reduces to Δp1= x2p10

• Lowering of the vapour pressure of a solution with several non-volatile solutes depends upon the sum of the mole fraction of different solutes which makes the above equation as-

Where n1 and n2 are the number of moles of solvent and solute respectively present in the solution.

• The above equation can be written as:

• For dilute solutions n2 << n1, hence ignoring n2 in the denominator we get

w1 and M1 = Masses and molar masses of solvent.

w2 and M2 = Masses and molar masses of solute.

COLLIGATIVE PROPERTIES

 COLLIGATIVE PROPERTIES



The properties that depend on the number of solute particles irrespective of their nature relative to the total number of particles present in the solution are called colligative properties. 

There are four colligative properties: 
1. Relative Lowering of vapour Pressure 
2. Elevation in Boiling Point 
3. Depression in freezing point 
4. Osmotic pressure

Henry’s law

  Henry’s law

Henry’s law states that at a constant temperature, the solubility of a gas in a liquid is directly proportional to the pressure of the gas.

The most commonly used form of Henry’s law states that

“the partial pressure of the gas in vapour phase (p) is proportional to the mole fraction of the gas (x) in the solution”.

This is expressed as:

 p = K_H x Here K_H is the Henry’s law constant.

Characteristics of K_H : 

K_H is a function of the nature of the gas.

Higher the value of K_H at a given pressure, the lower is the solubility of the gas in the liquid. 

K_H values increase with increase of temperature indicating that the solubility of gases increases with decrease of temp.



Applications of Henry’s law

1. In the production of carbonated beverages.

2. In the deep sea diving.

3. For climbers or people at high altitudes. 

Raoult’s Law as a special case of Henry’s Law 

According to Raoult’s law,

 p_i = x_i   p_i⁰  

In the solution of a gas in a liquid, one of the components is so volatile that it exists as a gas.
 Its solubility according to Henry’s law,

 p = K_H x. 

Thus, Raoult’s law becomes a special case of Henry’s law in which K_H becomes equal to p_i⁰ .

Henry’s law

• Henry’s Law establishes aquantitative relation betweenpressure and solubility of a gasin a solvent.
• This law is for gas-liquid solution.
• According to the law at a constant temperature, thesolubility of a gas in a liquid is directly proportional to thepressure of the gas.
• It can also be stated as - The partial pressure of the gas in vapour phase (p) is proportional to the mole fraction of the gas (x) in the solution.
• Mathematically,

p = KH x where KH = Henry’s law constant.

• At same temperature different gasses have different KH
• It is a function of the nature of the gas. At a given pressure increasing value of KH implies lower solubility of the gas in the liquid.
• Value of KH increases with the increase in temperature therefore solubility of gases increases with decreasing temperature. Due to this reason cold water is more sustainable for aquatic life than warm water.

 Problem: H2S, a toxic gas with rotten egg like smell, is used for the qualitative analysis. If the solubility of H2S in water at STP is 0.195 m, calculate Henry's law constant.

Solution:  

It is given that the solubility of H2S in water at STP is 0.195 m, i.e., 0.195 mol of H2S is dissolved in 1000 g of water.

Moles of water = 1000g / 18g mol-1 = 55.56 mol 

∴Mole fraction of H2S, x =  Moles of H2S / Moles of H2S+Moles of water 0.195 / (0.195+55.56)= 0.0035 

At STP, pressure (p) = 0.987 bar

According to Henry's law:

p = KH x  ⇒ KH = p / x

= 0.0987 / 0.0035 bar

= 282 bar

SOLUBILITY

SOLUBILITY

 Solubility of a solid in liquid Solubility of a substance is its maximum amount that can be dissolved in a specified amount of solvent.

Factors affecting the solubility of a solid in liquid :
1. Nature of solute and solvent 
2 Like dissolves like. 
For example, While sodium chloride and sugar dissolve readily in water, naphthalene and anthracene do not. 
On the other hand, naphthalene and anthracene dissolve readily in benzene but sodium chloride and sugar do not. 

2. Temperature :
 In a nearly saturated solution, 
If (Δ_solH > 0), the solubility increases with rise in temperature and If (Δ_solH < 0) the solubility decreases with rise in temperature. 

Effect of pressure : 
Does not have any significant effect as solids and liquids are highly incompressible.

Azeotropes

 Azeotropes 

A mixture of two liquids which has a constant boiling point and composition throughout distillation.

1. Azeotropes are binary mixtures having the same composition in liquid and vapour phase and boil at a constant temperature.
2.  Minimum boiling azeotrope The solutions which show a large positive deviation from Raoult’s law form minimum boiling azeotrope at a specific composition.
3.  For example, ethanol-water mixture containing approximately 95% of ethanol forms an azeotrope with boiling point 351.15 K.
   

Maximum boiling azeotrope :

1.  The solutions that show large negative deviation from Raoult’s law form maximum boiling azeotrope at a specific composition.
2. Nitric acid and water mixture containing 68% nitric acid forms an azeotrope with a boiling point of 393.5K 

An azeotrope  or a constant boiling point mixture is a mixture of two or more liquids whose proportions cannot be altered or changed by simple distillation. This happens because when an azeotrope is boiled, the vapour has the same proportions of constituents as the unboiled mixture.

Ideal and Non-ideal solutions

 Ideal and Non-ideal solutions

Ideal solutions :What is an ideal solution in chemistry?

(In chemistry, an ideal solution or ideal mixture is a solution with thermodynamic properties analogous to those of a mixture of ideal gases. ... The vapor pressure of the solution obeys Raoult's law, and the activity coefficient of each component (which measures deviation from ideality) is equal to one)
An ideal solution is the solution in which each component obeys Raoult’s law under all conditions of temperatures and concentrations.

Properties of Ideal solutions :

ΔH_MIXING = 0

ΔV_MIXING = 0

Intermolecular attractive forces between the A-A and B-B are nearly equal to those between A-B. 
Eg. solution of benzene and toluene, solution of n-hexane and n-heptane

Non – ideal solutions :What is an Non- ideal solution in chemistry?

(A non-ideal solution is a solution whose properties are generally not very predictable on account of the intermolecular forces between the molecules. None.Non-ideal solutions by definition cannot be dealt with through Raoult's Law. Raoult's Law is strictly forideal solutions only. A non-ideal solution)
 When a solution does not obey Raoult’s law over the entire range of concentration, then it is called non-ideal solution.

Solutions showing positive deviation from Raoult’s Law :

Solvent-Solute(A-B) type of force is weaker than Solute-Solute (B-B) &  Solvent-Solvent(A-A) forces.

The vapour pressure is higher than predicted by the law

ΔH_MIXING > 0

ΔV_MIXING > 0

Eg. ethanol and acetone, carbon disulphide and acetone 

Solutions showing negative deviations from Raoult’s law :

1. Solvent-Solute(A-B) type of force is stronger than the other two.
2. The vapour pressure is lower than predicted by the law.
3. ΔH_MIXING < 0
4. ΔV_MIXING < 0

For example,phenol and aniline, chloroform and acetone etc 

What is meant by azeotropic mixture?

An azeotrope (UK /əˈziːəˌtrəʊp/, US /əˈziəˌtroʊp/) or a constant boiling pointmixture is a mixture of two or more liquids whose proportions cannot be altered or changed by simple distillation. This happens because when an azeotrope is boiled, the vapour has the same proportions of constituents as the unboiled mixture.

Raoult’s Law

 Vapour pressures of solutions of solids in liquids and Raoult’s Law

A law stating that the freezing and boiling points of an ideal solution are respectively depressed and elevated relative to that of the pure solvent by an amount proportional to the mole fraction of solute.

• a law stating that the vapour pressure of an ideal solution is proportional to the mole fraction of solvent.

(Raoult’s law for non volatile solutes)

Q-1 what happens when a non volatile solute is added to a solvent?

1.  If a non-volatile solute is added to a solvent to give a solution, the number of solvent molecules escaping from the surface is correspondingly reduced, thus, the vapour pressure is also reduced.

Q-2 Which factor effects the vapour pressure of solvent ?

1.  The decrease in the vapour pressure of solvent depends on the quantity of non-volatile solute present in the solution, irrespective of its nature.

Q-3 What does Raoult law in its general form can be stated?  

1.  Raoult’s law in its general form can be stated as, for any solution the partial vapour pressure of each volatile component in the solution is directly proportional to its mole fraction.

Q-4 How to Denote binary solution?

1.  In a binary solution, let us denote the solvent by 1 and solute by 2.

Q-5 What happens when the solute is non-volatile? 

1. When the solute is non-volatile, only the solvent molecules are present in vapour phase and contribute to vapour pressure.

Q-6 How to use  Raoult's law?

1.  Let p₁ be the vapour pressure of the solvent, x₁ be its mole fraction, p⁰₁ be its vapour pressure in the pure state. Then according to Raoult’s law
2. p₁ 𝞪 x₁ and p₁ = x₁  p⁰₁ = p_(total)

If a solution obeys Raoult’s law for all concentrations, 

its vapour pressure would vary linearly from zero to the vapour pressure of the pure solvent.

According to Raoults law, for any volatile component of the solution. Pa=P°a×Xa. Hence vapour pressure is directly proportional to the mole fraction of solute.

If gas is a solute and liquid is the solvent, then according to Henry law, Pa=KaXa. Hence partial pressure of volatile component is directly proportional to mole fractionof that component.

Hence both are identical with only different proportionality constant's.

1. Raoult's law  is a law of thermodynamics established by French chemist François-Marie Raoult in 1887. It states that the partial vapor pressure of each component of an ideal mixture of liquids is equal to the vapour pressure of the pure component multiplied by its mole fraction in the mixture.
2. Henry's law is one of the gas laws formulated by William Henry in 1803 and states: "At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid."
3. Henry's law is one of the gas laws, formulated by the British chemist, WilliamHenry, in 1803. It states that: At a constant temperature, the amount of a given gas dissolved in a given type and volume of liquid is directly proportional to the partial pressure of that gas in equilibrium with that liquid

What is the relationship between Henry's law and Raoult's law?

Both are ideal vapor-liquid equilibrium relationships that hold for very dilute binary component mixtures.
  Henry's Law tends to approximate real behavior for the solute (i.e., lower concentration) at a very low concentration while Raoult's Law tends to approximate real behavior for the solvent in a dilute solution.
  They are both limit cases of an underlying binary vapor-liquid equilibrium relationship. 
 Unifying them basically requires that you replace the pressure terms with fugacity terms to accurately represent real gas behavior.

1. If you're familiar with chemical equilibria, it might be useful to think of the Henry's law 'constant' (not really a constant as it is dependent on temperature and ionic strength of the solution) as the equilibrium 'constant' of the reaction X(g) <-> X(aq), where X is the gas in question and is dissolving in water. (http://www.chemguide.co.uk/physi...) I.e. Henry' law constant is is the ratio of aqueous concentration to gas phase concentration of the gas AT EQUILIBRIUM.
2.  Any deviation from this, means the gas and liquid phases are not at equilibrium with respect to the particular gas and there will be a net transfer of gas molecules one way or another to move towards equilibrium (Le Chatelier's principle).
3.  Henry's law only holds for dilute solutions in single solvents (normally water).
4.  Henry's LAW basically says that the position of this equilibrium, i.e. the value of the equilibrium constant is always the same under the same conditions.
5. 
   
6. Raoult's law, as stated by one of the other answers, concerns the bulk components of a solution made up of different solvents, and their relative contributions to the total vapour pressure. In short, the total vapour pressure is the sum of the partial pressures:
7. 
   
8. vapour pressure of mixture = partial pressure of component A + partial pressure of component B +....
9. 
   
10. and each partial pressure is the product of the vapour pressure of a pure solution of the compound and the mole fraction of the compound in the solvent mixture (i.e. what proportion, by moles rather than mass, of the total mixture is made up of each particular compound):
11. 
    
12. partial pressure of A = vapour pressure of pure solution of A x mole fraction of A in mixture.
13. 
    
14. Both Henry's law and Raoult's make important assumptions about ideality of gases and solutions respectively.
15. 
    

Both laws can get into quite details and more complicated calculation, my answer just provide the base of both law. 

1. With a little more detail: 
2. Henry's law states that the solubility of a gas in a liquid is proportional to the pressure of the gas over the solution.
3. The equation is c = kP ,
4.  where c is the molar concentration in mol/L of the dissolved gas,
5.  P is the pressure (in atm) of the gas over the solution at equilibrium and k is a constant that depends only on temperature for a given gas.
6. Which means that when you want to mix a gas and a liquid,
7.  the amount of gas that will actually dissolve in the liquid is proportional to two things:
8.  the pressure of that gas at equilibrium over the solution and a constant k that changes depending on the gas and the temperature.
9.  By using the equation up there, you will find the molar concentration (mol/L) of the gas that will dissolve in the solution. 

On the other hand,

Raoult's law is used in a case where the solute (the smallest component of the solution) is non-volatile. The law considers that the vapor pressure of the whole solution will always be less than that of the pure solvent. Therefore, the vapor pressure of the solution will depends on the concentration of the solute. 

Raoult's law equation is P1 = X1 P1* (they are multiplied), where P1 is the vapor pressure of the solvent over the solution, X1 is the mole faction of the solvent in the solution and P1* is the vapor pressure of the pure solvent (if it was alone in the solution). 

Overall, the difference is that Henry's law takes care of what happen IN the solution when you have gas over it, while Raoult's law looks at what is happening OVER the solution when you mix a non-volatile solute to a solvent that has a known vapor pressure when it's pure (e.g. water). 
Henry's law will give you the molar concentration of a dissolved gas in the solution, Raoult's law will give you a vapor pressure over a solution after you mixed a solvent with a non-volatile solute. 
Hope it helped!

Raoult’s law explains partial vapour pressure of volatile component of the solution.

According to Raoult’s law,partial vapour pressure of volatile component is equal to product of vapour pressure of pure component and mole fraction of that component in the solution.

P =P⁰ X

Henry’s law explains the solubility of gas in liquid.

According to Henry’s law,pressure is directly proportional to the mole fraction of gas.

P = KH X

When KH and P⁰ are same ,then Raoult’s law becomes special case of Henry’s law. This happens in case of few ideal solutions of gas in liquid.

On the other hand,

1. Raoult's  law is used in a case where the solute (the smallest component of the  solution) is non-volatile.
2. The law considers that the vapor pressure of  the whole solution will always be less than that of the pure solvent.
3. Therefore, the vapor pressure of the solution will depends on the  concentration of the solute. 
4. 
   
5. Raoult's law equation is P1 = X1  P1* (they are multiplied),
6.  where P1 is the vapor pressure of the solvent  over the solution,
7.  X1 is the mole faction of the solvent in the  solution
8.   P1* is the vapor pressure of the pure solvent (if it was  alone in the solution). 
9. 
   
10. Overall, the difference is that Henry's  law takes care of what happen IN the solution when you have gas over it, 
11.  while Raoult's law looks at what is happening OVER the solution when  you mix a non-volatile solute to a solvent that has a known vapor  pressure when it's pure (e.g. water). 
12. Henry's law will give you the  molar concentration of a dissolved gas in the solution,
13.  Raoult's law  will give you a vapor pressure over a solution after you mixed a solvent  with a non-volatile solute.