Chemistry-ch10_solutions And Solubility

10 Solutions and solubility

10.1 Introduction to solutions and solubility •

•

Solution Homogeneous mixture of two or more pure substance – May be gaseous, liquid or solid –

Solvent –

•

Liquid of a liquid solution

Solute –

Dissolved substance in liquid solution

10.1 Introduction to solutions and solubility •

•

Solubility Maximum amount of solute that dissolves completely in a given amount of solvent at a particular temperature, –T

Saturated solution –

•

Solution in which no more solute will dissolve

Dissolution –

Process of dissolving a solute in a solvent to give a homogeneous solution

10.2 Gaseous solutions • All gases mix completely with all other gases in all proportions • Gases mix spontaneously hence ΔG for mixing process negative

10.2 Gaseous solutions • •

Enthalpy change when two gases are mixed is usually small Large positive Δmix S term ensures Δmix G negative

Intermolecular forces between individual gas molecules are tiny, and gas molecules are far • apart •

No impediment to the mixing of gases

10.2 Gaseous solutions • The magnitude of interatomic and intermolecular attractions in the condensed phase can determine whether or not two substances can mix

10.3 Liquid solutions • Gas–liquid solutions

– For a gas to dissolve in a liquid, the gas molecules must be able to disperse themselves evenly throughout the solvent

10.3 Liquid solutions • Gas–liquid solutions – Intramolecular forces between the solvent molecules are not negligible – Neither therefore is the value of ΔsolH, the enthalpy of solution

10.3 Liquid solutions • Gas–liquid solutions

10.3 Liquid solutions • Gas–liquid solutions – Solubility of gases vary significantly with temperature – Solubility of gases vary significantly with pressure

10.3 Liquid solutions • Gas–liquid solutions – For gases that do not react with the solvent, Henry’s law gives the relationship between gas pressure and gas solubility – cgas

c gas = kHpgas( constant T) is the concentration of the gas

– pgas is the partial pressure of the gas above the solution – kH is called the Henry’s law constant and is unique to each gas

10.3 Liquid solutions • Gas–liquid solutions – Equation is only true at low concentrations and pressures – An alternative expression of Henry’s law is: c1 c 2 = p1 p2 – c1 and p1 refer to initial conditions – c2 and p2 refer to final conditions

10.3 Liquid solutions • Liquid–liquid solutions – Formation of a liquid–liquid solution requires that the attractive forces present between the molecules of the two pure liquids is overcome – Two substances are MISCIBLE when they mix completely in all proportions – Two substances are IMMISCIBLE when they form two layers upon the addition of one to the other

10.3 Liquid solutions • Liquid–liquid solutions H H

H

H

C

C

H

H

δ− δ+ O H

H

H C

C C

C

C C

H H

H

Ethanol

Benzene

hydrogen bond H H C2H5

O

O

H H

H

H

O

O

O

H C2H5

H

Like–dissolves–like

10.3 Liquid solutions • Liquid–solid solutions – Basic principles remain the same – Solvation is when a solute molecule is surrounded by solvent molecules – Hydration occurs when solutes become surrounded by water molecules

10.3 Liquid solutions • Liquid–solid solutions – Like-dissolves-like – When intermolecular attractive forces within solute and solvent are sufficiently different, the two do not form a solution

10.3 Liquid solutions • Liquid–solid solutions

10.3 Liquid solutions • Liquid–solid solutions

10.3 Liquid solutions • Liquid–solid solutions – Temperature can have a significant effect on the solubility of a solid solute in a liquid

10.4 Quantification of solubility: the solubility product • Ionic salts are generally classified as being either soluble or insoluble in water AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) AgCl(s) Ag+(aq) + Cl–(aq) Ksp = [Ag+][Cl–]

• Ksp is called the solubility product MaXb(s) aMc+ (aq) + bXd– (aq) Ksp = [Mc+ ]a[Xd– ]b

10.4 Quantification of solubility: the solubility product • The relationship between Ksp and solubility – Molar solubility (s) molar concentration of a salt in its saturated solution – Molar solubility can be used to calculate Ksp , assuming that all of the salt that dissolves is 100% dissociated into its ions

10.4 Quantification of solubility: the solubility product

The solubility of AgBr in water is 1.3 × 10–4 g L–1 at 25 °C. Calculate Ksp for AgBr at this temperature AgBr(s) Ag+(aq) + Br–(aq) Ksp = [Ag+][Br–] m 1.3 × 1 0− 4 g −7 n= = = 6 . 9 × 10 mol − 1 M 1 8 7.77 gmol

[Ag+] = [Br–] = 6.9 × 10–7 mol L–1 Ksp = [Ag+][Br–] = (6.9 × 10–7 mol L–1 )(6.9 × 10–7 mol L–1 ) Ksp = 4.8 × 10–13 at 25 °C

10.4 Quantification of solubility: the solubility product

Calculate the molar solubility of lead iodide, PbI2, given that Ksp (PbI2) = 7.9 × 10–9 PbI2(s) Pb2+ (aq) + 2I–(aq) Ksp = [Pb2+ ][I–]2 Define the molar solubility of PbI2(s) as s mol L–1 [Pb2+ ] = s mol L–1 [I–] = 2s mol L–1 Ksp = (s)(2s)2 = (s)(4s2) = 4s3 = 7.9 × 10–9

−9 . 9 × 10 –3 3s = 7 s = 1.3 × 10 4 in water at 25 °C is The molar solubility of PbI2(s) 1.3 × 10–3 mol L–1

10.4 Quantification of solubility: the solubility product

• The common ion effect

– Any ionic salt is less soluble in the presence of a common ion, an ion that is in the salt PbCl2(s) Pb2+ (aq) + 2Cl–(aq) Ksp = [Pb2+ ][Cl–]2 Add Pb(NO3)2(aq) to saturated solution of PbCl2 instantaneously increases [Pb2+ ] and therefore Qsp (ionic product). Qsp > Ksp PbCl2 is precipitated

10.4 Quantification of solubility: the solubility product

What is the molar solubility of PbI2 in a 0.10 M NaI solution? PbCl2(s) Pb2+ (aq) + 2Cl–(aq) Ksp = [Pb2+ ][Cl–]2 = 7.9 × 10–9

PbI2(s)

Pb2+ (aq) +

2I–(aq)

Initial concentration (M)

0

0.10

Change in concentration (M)

+s

+2s

Equilibrium concentration (M) + 2s)2 = 7.9 ×s10–9 K = s(0.10 sp

0.10 + 2s

Ksp = s(0.10)2 = 7.9 × 10–9 Molar solubility of PbI2 in 0.10 M NaI solution is 7.9 × 10–7 M (s)

10.4 Quantification of solubility: the solubility product

• Prediction of precipitation

– Qsp > Ksp precipitate will form – Qsp < Ksp no precipitate will form AgCl(s) Ag+(aq) + Cl–(aq) Ksp = [Ag+][Cl–] = 1.8 × 10–10 [Ag+] = 5.0 × 10–7 mol L–1 [Cl–] = 5.0 × 10–5 mol L–1 Qsp = 2.5 × 10–11 Qsp < Ksp no precipitate will form

10.5 Colligative properties of solutions • Nonvolatile solutes – Solutes that can’t evaporate from solution – Vapour pressure of a solution of nonvolatile solutes is lower than the pure solvent

• Colligative properties – Depend on number of solute particles in solution rather than chemical identities

10.5 Colligative properties of solutions • Molarity – Amount of substance in a particular volume of solution amount of solute (mol) molarity (c ) = volume of solution (L) – Solutions (usually) increase in volume with increasing temperature – The molarity of a solution changes as the temperature changes

10.5 Colligative properties of solutions • Molality – Preferred method of expressing solution composition when colligative properties involved – Amount of solute per kilogram of solvent amount of solute (mol) molality (b ) = mass of solvent (kg) – Temperature independent

10.5 Colligative properties of solutions • Mole fraction – The number of moles of a particular component divided by the total number of moles of material in the solution – The mole fraction of A, XA, in a solution containing substances A, B and C nA XA = nA + nB + nC – Temperature independent

10.5 Colligative properties of solutions • Raoult’s law – Boiling point of a solution containing a nonvolatile solute is higher than that of the pure solvent – Boiling point of a solvent is the temperature at which the vapour pressure of the solvent is equal to the atmospheric pressure

10.5 Colligative properties of solutions • Raoult’s law psolution = Xsolvent p*solvent – psolution – vapour pressure of the solution – Xsolvent – mole fraction or solvent in the solution – p*solvent – vapour pressure of pure solvent – For a simple two component system – Provided the solution is sufficiently dilute

10.5 Colligative properties of solutions • Raoult’s law psolution = Xsolvent p*solvent Xsolvent = 1 – Xsolute psolution = (1 – Xsolute)p*solvent psolution = p*solvent – Xsolutep*solvent Δp = Xsolutep*solvent Δp = p*solvent – psolution

10.5 Colligative properties of solutions • Raoult’s law – A solution that obeys Raoult’s law is called an ideal solution – These solutions are generally dilute and have only small interactions between their constituent molecules

10.5 Colligative properties of solutions • Solutions containing more than one volatile component – For component A pA = XAp*A

– For component B pB = XBp*B

– Total pressure ptotal = XAp*A + XBp*B

10.5 Colligative properties of solutions • Boiling point elevation and freezing point depression – boiling point elevation ΔTb ΔTb = Kbb

– freezing point depression ΔTf ΔTf = Kfb

10.5 Colligative properties of solutions • Boiling point elevation and freezing point depression – Molal boiling point elevation constant • Kb • Units – K mol–1 kg

– Molal freezing point depression constant • Kf • Units – K mol–1 kg

10.5 Colligative properties of solutions • Boiling point elevation and freezing point depression – Molality of the solution •b • Units – mol kg–1

– Kbb and Kf b are properties of the solvent only and independent of the identity of the solute

10.5 Colligative properties of solutions • Osmosis and osmotic pressure – A membrane keeps mixtures and solutions organised and separated. – Semipermeable membranes allow selective substances to pass through – Dialysis occurs when a dialysing membrane allows both water and small solute particles through – Osmosis is a net shift of only solvent through an osmotic membrane

10.5 Colligative properties of solutions • Osmosis and osmotic pressure

10.5 Colligative properties of solutions • Osmosis and osmotic pressure – Osmotic pressure • Π

– In dilute aqueous solution Π = cRT ΠV = nRT

n c = V

– This is the van’t Hoff equation for osmotic pressure

10.5 Colligative properties of solutions • Osmosis and osmotic pressure

Isotonic

Hypertonic

Hypotonic

Osmometer

10.5 Colligative properties of solutions • Measurement of solute dissociation – Molal freezing point depression constant for water is 1.86 K mol–1 – 1.00 mol kg–1 NaCl freezes at about – 3.37 °C – NaCl(s)  Na+(aq) + Cl–(aq) – Solution has a a total molality of dissolved solute particles of 2 mol kg–1 – Theoretically, a 1.00 mol kg–1 NaCl solution should freeze at –3.72 °C

10.5 Colligative properties of solutions

i =

( ∆Tf ) measured

( ∆Tf ) calculatedas nonelectrolyte

10.5 Colligative properties of solutions • Measurement of solute dissociation

10.5 Colligative properties of solutions

mol of ionised ∆Tf 1.9 0K % ionisation = × 100 % = − 1 Kf mol of acid available 1.8 6K mol kg 0.02 −1 % ionisation = × 10 0% b = 1.0 2 mol kg 1.00 %ionisation= 2% b=

10.5 Colligative properties of solutions

O 2 C6 H5 C O H benzoic acid

O H O C6 H5 C C C6H5 O H O benzoic acid dimer