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solution
a homogenous mixture of 2+ substances
characteristics of solutions
single phase system
doesnt allow a beam of light to scatter
stable & cannot be separated through filtration
density, BP, RI etc. are uniform throughout
particle size of solute - 10^-7 - 10^-8
classification of solutions on the basis of solvent
aqueous: solute is dissolved in water
non-aqueous: solute is dissolved in something else
classification of solutions on the basis of solute
concentrated: large amt. of solute
diluted: small amt. of solute
classification of solutions on the basis of no. of components
binary: 2 components
tertiary: 3 components
mass/mass %
mass of component/ total mass x 100
vol/vol%
vol. of component/ total vol. x 100
mass/vol %
mass of component/ total vol. x 100
ppm
no. of parts of component/ total no. of parts x 10^6
mole fraction
no. of moles of component/ total no. of moles
molarity
no. of moles of solute/ vol. of sol (L)
(note: it varies with temperature)
molality
no. of moles of solute/ mass of solvent (kg)
(note: it is independent of temperature
solubility of solid in liquid
it is the maximum weight of solute that can be dissolved in 100g of solvent.
it depends on nature & temperature
concept of solubility of solid in liquid
when solid is added in liquid, the solid keeps dissolving and the concentration increases (dissolution)
but some of the particles already in the solution collide with the ones coming in and get separated out (crystallisation)
dynamic equilibrium: when the no. of particles going in is equal to the no. of particles separating out. at this point, the concentration of solution is max.
how nature of the solid & liquid affects the solubility
like dissolves like
polar solids dissolve in polar liquids (ex: NaCl & H2O)
non-polar solids dissolve in non-polar liquids (ex: some covalent compounds)
how temperature affects the solubility of solids in liquids
the affect of temperature follows the le-chatelier’s principle;
\Delta solH = --ve
dissolution is exothermic
solubility decreases as temperature increases
\Delta solH = +ve
dissolution is endothermic
solubility increases as temperature increases
affect of pressure on solubility of solids in liquids
no effect as solids and liquids are highly incompressible in nature
solubility of gas in liquid
it is the volume of gas that can be dissolved in 1cc of liquid
almost all gases are soluble in water, but at different extents
how nature of gas and liquid affects the solubility
gases which are easily liquified are more soluble in liquids
ex: NH3, HCl, SO2
how temperature affects solubility of gases in liquids
solubility decreases with increase in temperature
this is because molecular motion increases, resulting in them escaping out of the solution
also, dissolution of gas is exothermic
\Delta solH = —ve
increase in temperature → decrease in solubility
affect of pressure on solubility of gases in liquids
pressure increases → solubility increases
when gas is put in liquid,
some are dissolved in the liquid
some are above the liquid in gaseous state
if we increase the pressure by compressing the gas, more particles enter the liquid
the solubility would then keep increasing until a new equilibrium is achieved
henry’s law
at constant temperature, the partial pressure of the gas in vapour phase is directly proportional to the mole fraction (solubility) of the gas in solution
p = KhX
p-X graph → straight line
key points on Kh
it depends on the nature of gas
it is different in different solvents
Kh \alpha T
Kh \alpha 1/ solubility
henry’s law in relation to deep sea diving
scuba divers breathe high pressure air underwater. this high pressure increases the solubility of gases in blood
when the diver comes back up, the pressure decreases and the dissolved gases are released, which may block the capillaries
to prevent this, they wear tanks filled with air diluted with He
henry’s law in relation to high altitudes
at high altitudes, the pressure of O2 is low
hence, solubility in blood and tissues is also low, which may make us feel weak
limitations of henry’s law
if the pressure is too high or temperature is too low, the law becomes too inaccurate
the gas cant undergo any chemical change
it cant undergo association or dissolution in the solution
vapor pressure
it is the pressure exerted by the vapors, when the system is in equilibrium
how a liquid achieves equilibrium
at a certain temperature, when a liquid is in a closed vessel, some of the liquid gets converted into vapor to fill the empty space
more the evaporation, more the molecules in the vapor
these molecules move randomly, and some strike the water and get condensed
when the rate of evaporation and condensation is equal, it is in equilibrium
how nature of liquid affects the v.p
weaker the intermolecular forces, greater the v.p
this is because more molecules can enter the vapor
how temperature affects the v.p of liquid
v.p increases with increase in temperature
this is because the kinetic energy increases
raoult’s law
at a given temperature, in a liquid-liquid solution, the partial v.p of each component is directly proportional to its mole fraction in the solution
p1 \alpha X1 → p1 = p1*X1
p2 \alpha X2 → p2 = p2*X2
Ptotal from raoult’s law
P = P1 + P2
= P1*X1 + P2*X2
but, X1 + X2 = 1
thus,
P = P1* + (P2* - P1*)X2
raoult’s law with respect to henry’s law
we know,
p = p*X for any volatile component
if a gas is volatile, then
p = KhX
hence,
p* = Kh
v.p of solids in liquids
solids are non-volatile
hence, the v.p only depends on the solvent, which would be less than that of pure solvent
this is because the surface is covered by both solute and solvent molecules, reducing the no. of evaporated solvent molecules
ideal solutions
they are solutions which obey raoult’s law over the entire range of concentration
enthalpy of mixing, \Delta mix H = 0
volume of mixing, \Delta mix V = 0
this implies that no heat is evolved during mixing, and total volume remains same
explanation of ideal solutions
consider solute A and solvent B
in pure components, the intermolecular interactions will be A—A and B—B
in solution, there will be A—A, B—B and A—B
an ideal solution is one where
A—A = B—B = A—B
non-ideal solutions
they are solutions which do not obey raoult’s law over the entire range of concentration
enthalpy of mixing, \Delta mix H =/ 0
volume of mixing, \Delta mix V =/ 0
the v.p is either higher or lower as predicted by raoult’s law
if higher → positive deviation
if lower → negative deviation
explanation of positive deviation in non-ideal solutions
in this case, A—A, B—B are stronger than A—B
this means the molecules find it easier to escape in solution form (A—B), increasing the v.p
here,
Pa > Pa*Xa
Pb > Pb*Xb
\Delta mixH > 0 (+ve)
\Delta mixV > 0 (+ve)
here, heat is absorbed
explanation of negative deviation in non-ideal solutions
here, A—B is stronger than A—A and B—B
this results in molecules finding it harder to escape in solution form (A—B), decreasing the v.p
here,
Pa < Pa*Xa
Pb < Pb*Xb
\Delta mixH < 0 (-ve)
\Delta mixV < 0 (-ve)
here, heat is evolved
azeotropes
they are liquid mixtures having same composition in both liquid and vapor phase
they boil at constant temperature
it is not possible to separate these via fractional distillation
minimum boiling azeotropes
they are non-ideal solutions with large +ve deviation
here, there is an intermediate composition where
v.p is highest
b.p is lowest
maximum boiling azeotropes
non-ideal solutions with large -ve deviation
here, there is an intermediate composition where
b.p is highest
v.p is lowest
colligative properties
they are properties of solutions which depend only on the number of solute particles, irrespective of their nature
there are 4 main properties
relative lowering of vapor pressure
elevation in boiling point
depression in freezing point
osmotic pressure of solution
relative lowering of vapor pressure
we know, addition of non-volatile solute decreases the v.p
let X1 be mole fraction of solvent, X2 be that of solute, p1* the v.p of pure solvent and p be v.p of solution
thus, from raoult’s law, p = p1 = p1*X1
reduction in v.p, \Delta p1 = p1* - p1 = p1* - p(1 -X1)
but, X1 = 1 - X2
hence,
\Delta p1 = p1*X2
expression of relative lowering of v.p in terms of molar mass
let W1, W2 be masses and M1, M2 be molar masses of solvent and solute respectively
we know,
X2 = n2/ (n1 + n2)
= W2/M2 // (W1/M1) + (W2/M2)
but, for dilute solutions, n2 ««« n1
hence,
X2 = W2/M2 // W1/M1
substitute this in the previous conclusion, and we get
M2 = W2M1/ [W1(\Delta p1/ p1*)]
elevation in boiling point
the boiling point of a liquid is the temperature at which the vapor pressure and atmospheric pressure are equal
when a non-volatile solute is added, the b.p becomes higher than that of pure solvent
the difference in boiling point of solution Tb and pure solvent Tb* is called elevation in boiling point (\Delta Tb)
\Delta Tb = Tb - Tb*
boiling point elevation constant/ ebullioscopic constant
for dilute solutions, elevation in boiling point is directly proportional to the molal concentration of solute in the solution (molality)
\Delta Tb \alpha m
=> \Delta Tb = KbM
where Kb = boiling point elevation constant
expression of elevation in b.p in terms of molar maas
molality, m = n2 × 1000/ W1 (g)
but, n2 = W2/M2
thus,
m = W2 × 1000/ M2 x W1
but, \Delta Tb = Kbm
= Kb x W2 × 1000/ M2 x W1
thus,
M2 = Kb x W2 × 1000/ \Delta Tb x W1
depression in freezing point
the freezing point is the temp. at which the solid and liquid phases have same v.p
it decreases with the addition of non-volatile solute
this is because the v.p deceeases
if Tf* is f.p temp of pure solvent and Tf is the f.p temp of solution then
\Delta Tf = Tf* - Tf
freezing point depression constant/ cryoscopic constant
the depression in freezing point is directly proportional to the molality of the solute in solution
\Delta Tf\alpha m
=> \Delta Tf = Kf x m
relationship of Kf and Kb with enthalpy change
Kf = R x M1 x Tf²/ 1000 x \Delta fus H
Kb = R x M1 x Tb² / 1000 x \Delta vap H
where R = gas constant, M1 = m.m of solvent
osmosis
it is the phenomenon of the flow of solvent molecules through a semipermeable membrane from pure solvent to solution
if the osmosis takes place between solutions of different concentrations, then the solvent molecules will move from low to high solute concentration until equilibrium is attained
examples of osmosis
lot of salty food will lead to water retention in tissue cells and intercellular spaces because of osmosis
water movement from soil to plant
raw mango placed in concentrated salt solution will lose water via osmosis and shrivel into pickle
salting meats and veggies preserves them as through osmosis, the bacteria would lose water and die
osmotic pressure (\pi)
it is the excess pressure that must be applies to the solution side of membrane (as opposed to solvent), to prevent osmosis
it is a colligative property
osmotic pressure is proportional to the molarity (C) of solution at temperature T
\pi \alpha C
\pi = CRT
in terms of molar mass,
M2 = W2RT/V\pi
advantage of osmotic pressure to determine molar mass
it is done in room temperature
molarity is used instead of molality
magnitude is large even for very dilute solutions
types of solutions based on osmotic pressure
isotonic: two solutions having same osmotic pressure at a given temperature
no osmosis occurs as they have same molar concentration
eg: saline solution: 0.9% mass/vol NaCl solution which has same o.p as blood
hypertonic: a solution having more osmotic pressure than the other
solvent would move into the solution
eg: a sol. of more than 0.9% mass/vol NaCl would make the blood cells shrink as the water would come out of it
hypotonic: a solution having less osmotic pressure than the other
solvent would move out of the solution
eg: a sol. of less than 0.9% mass/vol NaCl would make the blood cells swell as water would enter it
reverse osmosis
if a pressure larger than the osmotic pressure is applied to the solution side, it would start flowing towards the solvent instead of the other way around. this is called reverse osmosis
it is used to desalinate sea water to get fresh water