Solution chemistry

S H A R D A P U B L I C S C H O O L , A L M O RA 1

SOLUTION CHEMISTRY

CLASS XII, UNIT:2 SUBJECT EXPERT

DR. TANUJA NAUTIYAL

SHARDA PUBLIC SCHOOLALMORA

S H A R D A P U B L I C S C H O O L , A L M O RA

Solutions are homogeneous mixtures of two or more pure substances.

In a solution, the solute is dispersed uniformly throughout the solvent.Table gives examples of each type.

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Solution


How Does a Solution Form?1. Solvent molecules attracted to surface ions.2. Each ion is surrounded by solvent molecules.3. Enthalpy (DH) changes with each interaction

broken or formed. If the solvent is water, the ions are hydrated. The

intermolecular force here is ion-dipole. The ions are solvated (surrounded by solvent)

.

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How Does a Solution Form?


To determine the enthalpy change, we divide the process into 3 steps.

1. Separation of solute particles.

2. Separation of solvent particles to make ‘holes’.

3. Formation of new interactions between solute and solvent.

ENERGY CHANGES IN SOLUTION

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ENERGY CHANGES IN SOLUTION

DH1Separation of solute molecules

DH2 Separation of solvent molecules

DH3 Formation of solute-solvent interactions


ENTHALPY CHANGES IN SOLUTION

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The intermolecular attractions that hold molecules together in liquids and solids also play a central role in the formation of solutions. When one substance (the solute) dissolves in another (the solvent), particles of the solute disperse throughout the solvent.

The solute particles occupy positions that are normally taken by solvent molecules. The ease with which a solute particle replaces a solvent molecule depends on the relative strengths of three types of interactions:solvent-solvent interaction, solute-solute

interaction, solvent-solute interaction


ENTHALPY CHANGES DURING DISSOLUTIONDHSoln = DH1+ DH2+ DH3

This equation is an application of Hess’s law.

The enthalpy of solution, DHsoln, can be either positive or negative.

DHsoln (MgSO4) = -91.2 kJ/mol --> Exothermic

DHsoln (NH4NO3) = 26.4 kJ/mol --> Endothermic

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If the solute-solvent attraction is stronger than the solvent-solvent attraction and solute-solute attraction, the solution process is favourable, or exothermic (ΔHsoln < 0).

If the solute-solvent interaction is weaker than the solvent-solvent and solute-solute interactions, then the solution process is endothermic (ΔHsoln > 0).


A saturated solution contains the maximum amount of a solute that will dissolve in a given solvent at a specific temperature.

An unsaturated solution contains less solute than it

has the capacity to dissolve.

A third type, a supersaturated solution, contains more

solute than is present in a saturated solution. Supersaturated solutions are not very stable.

Types of Solution


Factors Affecting Solubility

Polar substances tend to dissolve in polar solvents.

Nonpolar substances tend to dissolve in nonpolar solvents.

Example: Vitamin A is soluble in nonpolar compounds (like fats). Vitamin C Vitamin C is soluble in water.

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Nature of Solute and solvent: “like dissolves like”


Factors Affecting Solubility

Glucose (which has hydrogen bonding) is very soluble in water.

Cyclohexane (which only has dispersion forces) is not water-soluble.

Intermolecular forces = H-bonds; dipole-dipole; dispersion Ions in water also have ion-dipole forces soluble in water.

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The stronger the intermolecular attractions between solute and solvent, the more likely the solute will

dissolve.


Gases in Solution: the solubility of gases in water increases with increasing mass. The reason is Larger molecules have stronger dispersion forces.

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Solubility of some gases in water at 200C, with 1 atm Gas Pressure

N2Solubility

(M) 0.69 X 10-3

COSolubility

(M) 1.04 X 10-3

O2Solubility

(M) 1.38 X 10-3

ArSolubility

(M) 1.50 X 10-3

KrSolubility

(M)2.79 X 10-3


Effect of pressure on solubility:The solubility of liquids and solids does not change appreciably with pressure. But, the solubility of a gas in a liquid is directly proportional to its pressure.

Henry’s Law:

Sg is the solubility of the gas;k is the Henry’s law constant for that gas in that solvent;

Pg is the partial pressure of the gas above the liquid.

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Sg = kPg


Effect of temperature:Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

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Effect of temperature:The opposite is true of gases. Higher temperature drives gases out of solution.Carbonated soft drinks are more “bubbly” if

stored in the refrigerator.Warm lakes have less O2 dissolved in them

than cool lakes.

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Ways of Expressing Concentrations of Solutions Mass percentage: = (mass of A in solution / total mass of solution)

X 100

Parts per Million (ppm): = (mass of A in solution/total mass of

solution) X106

Parts per Billion (ppb) = (mass of A in solution/total mass of

solution) X109

Mole Fraction (X): The mole fraction of a component of a solution, say, component A, is written XA

XA = (moles of A / total moles in solution)

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Ways of Expressing Concentrations of SolutionsMolarity (M):the number of moles of solute in 1 L of solution;that is,

(moles of solute / L of solution)

Molality (m): Molality is the number of moles of solute dissolved in 1 kg (1000 g) of solvent—that is, (moles of solute / kg of solvent)

Normality (N): the gram equivalents in 1 L of solution; (gram equivalents of solute / L of solution)

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COLLIGATIVE PROPERTIESColligative properties depend only on the number of solute particles present, not on the identity of the solute particles. Among colligative properties are

Vapor pressure lowering Boiling point elevation

Melting point depressionOsmotic pressure

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VAPOUR PRESSUREAs solute molecules are added to a solution, the solvent become less volatile (=decreased vapor pressure).

Solute-solvent interactions contribute to this effect.Therefore, the vapor pressure of a solution is lower than that of the pure solvent.

If a solute is non volatile (that is, it does not have a measurable vapor pressure), the vapor pressure of its solution is always less than that of the pure solvent.

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RAOULT’S LAW Thus, the relationship between solution vapor pressure and

solvent vapor pressure depends on the concentration of the solute in the solution. This relationship is expressed by Raoult’s law,

Raoult’s Law: It states that the vapor pressure of a solvent over a solution, PA is given by the vapor pressure of the pure solvent, P°A, times the mole fraction of the solvent in the solution, = XA

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PA = XAPA


RAOULT’S LAWThere are two cases we can consider.

1. Volatile solute – both solvent and solute are found in the vapor above the solution. A solution of ethanol in is an example.

2. Non-volatile solute – only the solvent has a vapor pressure. The solute does not contribute to the pressure so there is a “vapor pressure lowering”.

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TOTAL PRESSURE FOR IDEAL SOLUTION

P1 = X1P01 Ptotal = P1 + P2

P1 = X1P01 P2 = X2P02

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FOR AN IDEAL SOLUTIONThe total vapor pressure over an ideal solution is given by

Ptotal = P1 + P2 In a solution containing only one solute, X1 = 1 – X2, where X2 is the

molefraction of the solute.

= X1P1 + X2P2

= (1 - X2)P1 + X2P2

= P1+ X2 (P2- P1) A plot of the total pressure has the form of a straight line.

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BOILING POINT ELEVATIONSolute-solvent interactions also cause solutions to have higher boiling points

and lower freezing points than the pure solvent.

The boiling point elevation (DTb ) is defined as the boiling point of the solution (Tb) minus the boiling point of the pure solvent (T°b):

DTb = Tb - T°bThe value of DTb is proportional to the vapor-pressure lowering, and so is also

proportional to the concentration (molality) of the solution. That is, D Tb α mThe change in boiling point is proportional to the molality of the solution:

where Kb is the molal boiling point elevation constant, a property of thesolvent.

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DTb = Kb m



FREEZING POINT DEPRESSION A non scientist may remain forever unaware of the boiling-

point elevation phenomenon, but a careful observer living in a cold climate is familiar with freezing-point depression. Ice on frozen roads and sidewalks melts when sprinkled with salts such as NaCl or CaCl2. This method of thawing succeeds because it depresses the freezing point of water.

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FREEZING POINT DEPRESSIONThe freezing point depression (DTf) is defined as the freezing point of

the puresolvent (T °f) minus the freezing point of the solution (T f):

DTf = Tf - T °f

Because T°f . Tf, DT°f is a positive quantity. Again, DTf is proportional to the

concentration of the solution:DTf α mDTf = Kf m

where m is the concentration of the solute in molality units, and Kf is the molal

freezing-point depression constant. We use molality since the solvent properties

do matter and therefore, we would like to calculate the result in relation to a fixed

amount of solvent (1 kg). The constant Kf depends on the solvent and is a

function of the enthalpy of freezing and the molar mass. 29


FREEZING POINT DEPRESSION

DTb = Kb m

DTf = Kf mHere Kf is the molal freezing point depression constant of the solvent.

The similarity between freezing point depression and boiling point elevation is not accidental.

They both depend on the shift in the equilibrium caused when a solution is formed and changes the equilibrium of the liquid with respect to the solid or vapor.

In both equations, DT does not depend on what the solute is, but only on how many particles are dissolved.

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MOLAL BOILING-POINT ELEVATION AND FREEZING-POINT DEPRESSION CONSTANTS OF SEVERAL COMMON LIQUIDS

Solvent Freezing Kf Boiling Kb(°C)* (°C/m) (°C)*

(°C/m)

Water 0 1 .86 100 0.52Benzene 5.5 5.12 80.1 2.53Ethanol −117.3 1.99 78.4 1.22Acetic acid 16.6 3.90 117.9 2.93Cyclohexane 6.6 20.0 80.7 2.79

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COLLIGATIVE PROPERTIES OF ELECTROLYTES The study of colligative properties of electrolytes requires a

slightly different approach than the one used for the colligative properties of nonelectrolytes. The reason is that electrolytes dissociate into ions in solution, and so one unit of an electrolyte compound separates into two or more particles when it dissolves. Because these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) show greater changes than those of nonelectrolytes. e.g. NaCl dissociates to form 2 ion particles; its limiting Van’t Hoff factor is 2.

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VAN’T - HOFF FACTOR:

One mole of NaCl in water does not really give rise to two moles of ions.

Some Na+ and Cl− re associate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl. Some Na+ and Cl− re associate as hydrated ion pairs, so the true concentration of particles is somewhat less than two times the concentration of NaCl.

Re association is more likely at higher concentration. Therefore, the number of particles present is concentration

dependent.

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THE VAN’T HOFF FACTORWe modify the previous equations by multiplying by the van’t

Hoff factor, ii = actual number of particles in solution after dissociation /number of

formulaunits initially dissolved in solution.

i = 1 for nonelectrolytes

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DTf = Kf .m.i


OSMOSIS Semipermeable membranes allow some particles to pass

through while blocking others. In biological systems, most semipermeable membranes (such as cell walls) allow water to pass through, but block solutes.

In osmosis, there is net movement of solvent from the area of higher solvent concentration (lower solute concentration) to the are of lower solvent concentration (higher solute concentration.

Water tries to equalize the concentration on both sides until pressure is too high.

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OSMOSIS

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MOLAR MASS FROM COLLIGATIVE PROPERTIESWe can use the effects of a colligative property such as osmotic

pressure todetermine the molar mass of a compound. The Vant Hoff’s equation can be modified to form used for the

determination ofmolar mass by osmometry.

Here we related to the concentration C in mol/lt to the concentration w in gms/lt

and the molar mass Mm in grams/mole. The experimental configuration uses the measurement of height as an

estimate of the osmotic pressure. The equation = ρ gh is used (h = / ρ g).

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OSMOTIC PRESSUREThe pressure required to stop osmosis, known as osmotic pressure, , is

where M is the molarity, n expresses the number of moles of solute, an d n/V, of the solution, R is the gas constant (0.0821 L.atm/K.mol),. This equation is called the Van't Hoff equation for osmotic pressure. If the osmotic pressure is the same on both sides of a membrane

(i.e., the concentrations are the same), the solutions are isotonic. If the solute concentration outside the cell is greater than that inside

the cell, the solution is hypertonic.

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= MRT


OSMOSIS IN CELLS If the solute concentration outside the cell is less than that

inside the cell, the solution is hypotonic.

Water will flow into the cell, and hemolysis results.

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COLLOIDSSuspensions of particles larger than individual ions or

molecules, but too small to be settled out by gravity.

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Among the most important colloids are those in which the dispersing medium is water. Such colloids are divided into two categories called hydrophilic, or water-loving, and hydrophobic, or water-fearing. Hydrophilic colloids are usually solutions containing extremely large molecules such as proteins. In the aqueous phase, a protein like haemoglobin folds in such a way that the hydrophilic parts of the molecule, the parts that can interact favourably with water molecules by ion-dipole forces or hydrogen-bond formation, are on the outside surface

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TYNDALL EFFECTColloidal suspensions can scatter rays of light. This

phenomenon is known as the Tyndall Effect When a beam of light passes through a colloid, it is scattered by the

dispersed phase No such scattering is observed with ordinary solutions because the solute molecules are too small to interact with visible light. Another demonstration of the Tyndall effect is the scattering of sunlight by dust or smoke in the air.

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COLLOIDS IN BIOLOGICAL SYSTEMSSome molecules have a polar, hydrophilic (water-loving) end

and a nonpolar, hydrophobic (water-hating) end.

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USING COLLIGATIVE PROPERTIES TO DETERMINE MOLAR MASS

The colligative properties of nonelectrolyte solutions provide a means of determining the molar mass of a solute. Theoretically, any of the four colligative properties is suitable for this purpose.

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DTb = Kb . m

DTf = Kf . m

= MRT

PA = XAPA


SUMMARY Solutions are homogeneous mixtures of two or more

substances, which may be solids, liquids, or gases.

The ease of dissolving a solute in a solvent is governed by intermolecular forces. Energy and the disorder that results when molecules of the solute and solvent mix to form a solution are the forces driving the solution process.

The concentration of a solution can be expressed as percent by mass, mole fraction, molarity, and molality.

Increasing temperature usually increases the solubility of solid and liquid substances and usually decreases the solubility of gases in water.

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SUMMARY

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According to Henry’s law, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas over the solution.

Raoult’s law states that the partial pressure of a substanceA over a solution is equal to the mole fraction (XA) of A times the vapor pressure (P°A) of pure A. An ideal solution obeys Raoult’s law over the entire range of concentration. In practice, very few solutions exhibit ideal behaviour.

Vapor-pressure lowering, boiling-point elevation, freezing point

depression, and osmotic pressure are colligative properties of solutions; that is, they depend only on the number of solute particles that are present and not on their nature.


SUMMARY

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In electrolyte solutions, the interaction between ions leads to the formation of ion pairs.

The Van’t Hoff factor provides a measure of the extent of dissociation of electrolytes in solution.

A colloid is a dispersion of particles (about 1 X103 pm to 1 X 106 pm) of one substance in another substance.

A colloid is distinguished from a regular solution by the Tyndall effect, which is the scattering of visible light by colloidal particles.

Colloids in water are classified as hydrophilic colloids and hydrophobic colloids.


SHARDA PUBLIC SCHOOL

Presentation by Dr. Tanuja Nautiyal

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Solution chemistry

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