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58 Terms

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solution

a homogenous mixture of 2+ substances

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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

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classification of solutions on the basis of solvent

  • aqueous: solute is dissolved in water

  • non-aqueous: solute is dissolved in something else

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classification of solutions on the basis of solute

  • concentrated: large amt. of solute

  • diluted: small amt. of solute

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classification of solutions on the basis of no. of components

  • binary: 2 components

  • tertiary: 3 components

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mass/mass %

mass of component/ total mass x 100

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vol/vol%

vol. of component/ total vol. x 100

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mass/vol %

mass of component/ total vol. x 100

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ppm

no. of parts of component/ total no. of parts x 10^6

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mole fraction

no. of moles of component/ total no. of moles

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molarity

no. of moles of solute/ vol. of sol (L)

(note: it varies with temperature)

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molality

no. of moles of solute/ mass of solvent (kg)

(note: it is independent of temperature

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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

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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.

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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)

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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

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affect of pressure on solubility of solids in liquids

no effect as solids and liquids are highly incompressible in nature

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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

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how nature of gas and liquid affects the solubility

gases which are easily liquified are more soluble in liquids

ex: NH3, HCl, SO2

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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

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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

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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

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key points on Kh

  • it depends on the nature of gas

  • it is different in different solvents

  • Kh \alpha T

  • Kh \alpha 1/ solubility

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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

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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

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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

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vapor pressure

it is the pressure exerted by the vapors, when the system is in equilibrium

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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

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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

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how temperature affects the v.p of liquid

v.p increases with increase in temperature

this is because the kinetic energy increases

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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

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Ptotal from raoult’s law

P = P1 + P2

= P1*X1 + P2*X2

but, X1 + X2 = 1

thus,

P = P1* + (P2* - P1*)X2

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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*)]

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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*

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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

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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

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