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Theoretical–Physical Advanced Level Chemistry

Equilibria – Chemical Equilibrium Revision Notes PART 8.7

 8.7 (i) The Solubility of covalent compounds

and (ii) Miscible and immiscible liquids

What factors control how soluble a covalent compound is in water or organic solvents. Why do some liquids mix (miscible) and others not (immiscible)? In both situations the role of intermolecular forces and intermolecular bonding is discussed.

(c) doc b KS4 Science GCSE/IGCSE Notes on reversible reactions and chemical equilibrium

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibrium and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series

8.7 Solubility of solid covalent compounds and miscible/immiscible of liquids

An approach to the covalent case studies described below

(Solubility of inorganic salt like compounds are considered in Thermodynamics Part 2)

(solubility product is dealt with in Equilibria section 4.2 and partition calculations in section 4.1)

  • The solubility of one material in another is a consequence of the physical and chemical interactions involved.

  • A solvation process may involve intermolecular force interactions or a chemical change or both.

  • When solids dissolve in liquids, a resulting saturated solution (maximum solubility) is another equilibrium situation.

  • Similarly mixing two immiscible liquids is also an equilibrium situation even if only a trace of one dissolves in the other.

  • It is important that you have read the section 8.2.1 on the nature of intermolecular forces because they all get mentioned again here, but not in as much detail, I will assume a working knowledge of them in section 8.7 on solubility and miscibility.

  • In considering whether A (liquid or solid solute) will dissolve in a liquid (e.g. the solvent) three possible interactions should be considered i.e. the intermolecular forces involved and solvation processes.

  • The three 'interactions' will simply be A...A, A...B and B...B associations.

  • If either A...A or B...B associations are strong then one of them will dominate and A and B will be less likely to mix if two liquids (miscible) i.e. if two liquids, they are likely to be immiscible OR solid A won't dissolve in liquid B.

  • However if an A...B interaction is strong or comparable to A...A or B...B associations then two liquids are likely to be miscible and mix with each other OR solid A dissolves in liquid B.

  • This conceptual way of thinking is a bit more sophisticated than 'like dissolves like', but the latter has much to commend it as a rough rule of thumb.

  • If two liquids dissolve appreciably in each other but still two layers are formed above and below a meniscus, the two liquids may be referred to as partially miscible.

  • If you can think like this, then most cases of what dissolves/mixes with what are usually quite understandable!

The cases discussed below are: Case study 8.7.1 Mixing hydrocarbons and water * Case study 8.7.2 The solubility of  alcohols in water * Case study 8.7.3 The solubility of carboxylic acids in water * Case Study 8.7.4 Why are pairs of some organic solvents miscible and others immiscible? * Case Study 8.7.5 Are dissolved molecules always in the same molecular state whatever the solvent? * 8.7.6 The solubility of some gases in water


 

Case study 8.7.1 Mixing hydrocarbons and water

  • The solubility of hydrocarbons in water is extremely low and for practical purposes it is virtually zero.

  • It doesn't matter whether it is alkanes (lower–small or higher–large), alkenes, arenes (aromatic hydrocarbons like benzene, methyl benzene etc.) or even crude oil itself, non of them can associate with water.

  • Hydrocarbon – hydrocarbon intermolecular forces are only the weak instantaneous dipole – induced dipole interactions.

  • The polar water – oil interaction will be of a permanent dipole – induced dipole nature, but even if this is greater than the hydrocarbon – hydrocarbon attraction, what will dominate this situation is the strongest of the intermolecular forces, namely the hydrogen bonds between water molecules.

  • ....δ:O–Hδ+ llll δ:O–Hδ+....  rules ok!

  • If a hydrocarbon molecule is 'shaken' into water, the hydrogen bonds will be disrupted and will tend to reform and quite literally squeeze out the hydrocarbon molecules.

  • I would expect methane will be the most soluble hydrocarbon in water because it is the smallest possible hydrocarbon and its 'insertion' will cause the minimum disruption to water's hydrogen bonds.

  • solubility of methane, ethene, ethane etc.?


 

Case study 8.7.2 The solubility of  alcohols in water

linear aliphatic alcohol structural formula Mr solubility g/100g water solubility mol/dm3
methanol CH3OH 32 miscible miscible
ethanol CH3CH2OH 46 miscible miscible
propan–1–ol (1–propanol) CH3(CH2)2OH 60 miscible miscible
butan–1–ol (1–butanol) CH3(CH2)3OH 74 8.0 1.08
pentan–1–ol (1–pentanol) CH3(CH2)4OH 88 2.8 0.32
hexan–1–ol (1–hexanol) CH3(CH2)5OH 102 0.6 0.059
heptan–1–ol (1–heptanol) CH3(CH2)6OH 116 0.093 0.008
octan–1–ol (1–octanol) CH3(CH2)7OH 130 0.054? 0.004?
nonan–1–ol (1–nonanol) CH3(CH2)8OH 144 ~insoluble ~insoluble
decan–1–ol (1–decanol) CH3(CH2)9OH 158 ~insoluble ~insoluble
  • The first three lower members of the alcohol homologous series are completely miscible with water.

  • This is because all the A...A, B...B and B...B interactions are all similar, i.e. they can all hydrogen bond with each other.

  • However, although the hydroxy O–H group is hydrophilic ('solvent liking') and readily hydrogen bonds with water, as well as itself, there is a counter trend of the hydrophobic ('solvent hating') alkyl group which disrupts the hydrogen bonds in water.

  • Therefore, as the alkyl hydrocarbon chain increases in length, its greater hydrogen bond disruption results in a rapidly decreasing solubility so that by the time you reach octanol all successive alcohols are effectively insoluble.

  • Note that if more than one hydroxy group is present then the solubility will be increased e.g.

    • Liquid propane–1,2,3–triol (glycerol) has the structure HOCH2CH(OH)CH2OH (Mr = 92), so with its triple hydrogen bonding sites the glycerol – water interactions are strong and both liquids are completely soluble in each other i.e. fully miscible, despite being as bulky as pentan–1–ol on a molecular mass basis.

    • Sugars such as glucose (Mr = 180), despite the higher molecular mass, are highly soluble in water due to the multiple hydrogen bonding sites on the molecule (5 O–H groups).

    • Also note the increase in viscosity ('stickiness') of glucose (solid) > liquid propane–1,2,3–triol > liquid propane–1,2–diol > propan–1–ol,

    • which is accounted for in terms of increased hydrogen bonding as well as the increased instantaneous dipole – induced dipole forces from the increasing number of electrons.


 

Case study 8.7.3 The solubility of carboxylic acids in water

linear carboxylic acid structural formula Mr solubility g/100g water solubility mol/dm3
methanoic acid HCOOH 46 miscible miscible
ethanoic acid CH3COOH 60 miscible miscible
propanoic acid CH3CH2COOH 74 miscible miscible
butanoic acid CH3(CH2)2COOH 88 miscible miscible
pentanoic acid CH3(CH2)3COOH 102 5.0 0.49
hexanoic acid CH3(CH2)4COOH 116 1.1 0.095
heptanoic acid CH3(CH2)5COOH 130 0.24 0.017
octanoic acid CH3(CH2)6COOH 144 0.068 0.0047
nonanoic acid CH3(CH2)7COOH 158 0.03 0.0018
decanoic acid CH3(CH2)8COOH 172 ~insoluble ~insoluble
  • The solubility trend is similar to that of the linear alcohols in case study 8.7.2 as are the reasons for the trend.

  • The first four lower members of this carboxylic acid homologous series are completely miscible with water and the solubilities, on a molar basis, tend to be a bit higher than for alcohols, presumably because they are more polar and have two sites for H bonding (C=O and O–H).


 

Case 8.7.4 Why are some pairs of  organic solvents miscible and others immiscible?

(in these examples the cases are based on mixing equal volumes of each liquid)

  • Propanone and benzene are miscible

    • (polar) propanone – propanone intermolecular forces = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole)

    • (non–polar) benzene – benzene intermolecular forces =  (instantaneous dipole – induced dipole)

    • Although propanone is polar, most of its intermolecular force is still derived from instantaneous dipole – induced dipole attractive forces and the extra permanent dipole – permanent dipole force does not prevent miscibility.

  • Propanone and butan–1–ol are miscible

    • (polar) propanone – propanone intermolecular forces = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole)

    • (more polar) butan–1–ol – butan–1–ol intermolecular forces = (instantaneous dipole – induced dipole) + (hydrogen bonding)

    • Both molecules are polar and butan–1–ol can hydrogen bond with propanone as well as with itself, so miscibility is no problem.

  • Glycerol (propan–1,2,3–triol) and propanone are immiscible

    • (polar) propanone – propanone intermolecular forces = (instantaneous dipole – induced dipole) + (permanent dipole – permanent dipole)

    • (much more polar) glycerol – glycerol intermolecular forces = (instantaneous dipole – induced dipole) + (multiple hydrogen bonding)

    • Glycerol has the structure HOCH2CH(OH)CH2OH, so with its triple hydrogen bonding sites the glycerol – glycerol attraction will outweigh any glycerol – propanone interactions.

  • Most organic liquids readily dissolve in butan–1–ol (1–butanol)

    • Being highly polar with its hydroxy group, butan–1–ol will readily dissolve many other polar organic compounds.

    • However, the alkyl hydrocarbon tail means the hydrogen bonding between the alcohol molecules doesn't totally dominate the situation and many relatively non–polar organic liquids will also dissolve in butan–1–ol.


 

Case 8.7.5 Are dissolved molecules always in the same molecular state whatever the solvent?

  • For most organic compounds the molecular state is the same irrespective of the solvent e.g.

    • Sugars e.g. C6H12O6 dissolving in water or alcohol.

    • Wax dissolving in other solvents such as hydrocarbon liquids or chloroalkane solvents.

    • Ethanol dissolving in water or other organic solvents such as propanone.

  • Carboxylic acids

    • In water, typical dissolved carboxylic acids exist mainly as the solvated molecule, a few % ionised and virtually non of the dimer which can be formed by hydrogen bonding.

    • However they dimerise in most organic solvents and are quite soluble in many of them.

    • This because most organic solvents such as hydrocarbons, chloroalkanes are not as polar as the acid itself (or water!).

    • Therefore the strongest intermolecular force that can operate is the hydrogen bonding between two carboxylic acid molecules.

    • In other words A...A solute interactions are stronger than A...R solute – solvent OR B...B solvent – solvent intermolecular forces.

  • Iodine

    • Non–polar iodine has a very low solubility in water because the strongest interactions possible in the mixture are H2O...H2O via hydrogen bonding.

    • So when the iodine molecules try to dissolve, its interaction with water is so weak that it fails to disrupt the hydrogen bonds between water molecules, hence the solubility is low.

    • Iodine is much more soluble in less polar solvents like hydrocarbons and chloroalkanes where the A...A, A...B and B...B intermolecular attractive forces are more comparable.

    • In both pure water and organic solvents iodine exists as the iodine molecule I2(aq) or I2(organic solvent).

    • Although iodine has a low solubility in pure water, it is much more soluble in potassium iodide solution which is a much more convenient for delivering a useful solution of iodine.

    • What actually happens when you add solid iodine to potassium iodide solution is that a chemical reaction takes place that more than compensates for the disruption of hydrogen bonding in water. The iodine forms the 'tri–iodide' ion.

      • (i)  I2(s) (c) doc b I2(aq)  and then  (ii) I2(aq) + I–(aq) (c) doc b [I3](aq) ....

      • .... and the 2nd equilibrium shifts the first equilibrium to the right, i.e. an increase in solubility,

      • but remember, this is due to a chemical reaction, NOT a change in the intermolecular forces.

  • Hydrogen chloride gas


 

8.7.6 The solubility of some gases in water

  • The maximum solubility is quoted as the volume of gas dissolved in 1 cm3 of water at 0oC at 1 atm pressure of the gas (101325 Pa).

    • This corresponds to the equilibrium X(g) (c) doc b X(aq) where X is the gas dissolving in water.

    • This is only quoted specifically where the gas reacts with water moving this equilibrium further to the right i.e. increasing solubility.

  • From Avogadro's Law, equal volumes of gases contain the same number, the data represents solubility on a molar basis.

  • Solubility of gases decreases on increase in temperature. and increases with increase in pressure of the gas.

  • In most cases the solubility is low because the dissolving molecule is disrupting the hydrogen bonds between water molecules.

  • * However, in a few cases, there is a chemical reaction with water OR strong hydrogen bonding occurs between solute and the solvent water, and in such cases the solubility is considerably increased.

  • When comparing any two of the molecules in the list, you can produce reasonable arguments to explain either the difference or similarity in solubility and I've quoted a few such arguments.

Gases in alphabetical order Solubility at 0oC Structure Mr electrons Comments
* Ammonia 1300 NH3 polar molecules 17 10 Extremely soluble, can strongly hydrogen bond with water and a few % of the dissolved ammonia reacts with water – ionisation.

 llll δ–N–Hδ+ llll δ–O–Hδ+ llll δ–N–Hδ+ llll δ–O–Hδ+ llll  etc.

NH3(g) (c) doc b NH3(aq) + H2O(l) (c) doc b NH4+(aq) + OH(aq)

Argon 0.056 Ar atoms 40 18 Weak instantaneous dipole – induced dipole interactions, very low solubility
Carbon dioxide 1.71 CO2 linear molecules 44 22 The effect of the permanent dipoles (δ+C=Oδ–) is cancelled out because of the linearity of the molecule, but its greater solubility compared to molecules like oxygen, nitrogen etc. may be due to hydrogen bonding? and the tiny % of the ionisation reaction:

CO2(g) (c) doc b CO2(aq) + H2O(l) (c) doc b H3O+(aq) + HCO3(aq)

Carbon monoxide 0.035 CO diatomic molecules 28 14 Not a very polar molecule, so low solubility via  instantaneous dipole – induced dipole forces
Chlorine 4.61 Cl2 non–polar diatomic molecules 71 34 A larger molecule of 34 electrons increasing the instantaneous dipole – induced dipole forces so increasing the solubility compared to other molecules of less electrons in the table and also a tiny % of the following ionisation reaction:

Cl2(g) (c) doc b Cl2(aq) + H2O(l) (c) doc b H3O+(aq) + Cl(aq) + HClO(aq)

Ethene 0.25 C2H4 non–polar planar molecules 28 16 Weak interaction but 5 times more soluble than methane, perhaps because the pi electron clouds of the C=C double bond are more polarizable increasing the instantaneous dipole – induced dipole interaction?
Helium 0.0094 He atoms of only 2 electrons 4 2 Only the very weakest instantaneous dipole – induced dipole forces despite the minimum disruption of water's hydrogen bonds, hence extremely low solubility.
Hydrogen 0.021 H2 non–polar diatomic molecules of only 2 electrons 2 2 Only very weak instantaneous dipole – induced dipole forces despite the minimum disruption of water's hydrogen bonds, hence very low solubility.
* Hydrogen chloride 506 HCl polar diatomic molecules 36.5 18 Extremely soluble because of a ~100% reaction with water to give ~10 molar hydrochloric acid.

HCl(g) + H2O(l) ==> H3O+(aq) + Cl(aq)

Similarly hydrogen bromide (HBr) and hydrogen iodide (HI) are also very soluble for the same reason.

Hydrogen sulfide 4.68 H2S slightly polar bent molecule 34 18 Hydrogen sulphide is slightly polar but the moderate solubility is due mainly to instantaneous dipole – induced dipole forces and a tiny % chemical reaction via

H2S(g) (c) doc b H2S(aq) + H2O(l) (c) doc b H3O+(aq) + HS(aq)

Nitrogen 0.024 N2 non–polar diatomic molecules 28 14 Low solubility due to being a more bulky molecule than say H2 or He but increased electrons increases instantaneous dipole – induced dipole forces gives nitrogen a slightly greater solubility.
Methane 0.054 CH4 simplest hydrocarbon molecule 16 10 Only very weak instantaneous dipole – induced dipole forces despite the minimum disruption of water's hydrogen bonds for a small molecule, hence very low solubility.
Oxygen 0.049 O2 diatomic molecules 32 16 Low solubility due to being a more bulky molecule than say H2 or He but increased electrons increases instantaneous dipole – induced dipole forces gives oxygen a slightly greater solubility.
Sulfur dioxide 79.9 SO2 bent polar molecules 64 32 Sulphur dioxide

Henry's Law and Gas Solubility

  • Henry's Law states that the mass of slightly soluble gas that dissolves in a definite mass of liquid at a given temperature is very nearly proportional to the partial pressure of that gas (in contact with the liquid).

  • This law holds for gases that do NOT react with the liquid solvent (e.g. hydrogen chloride in water) OR have a very strong intermolecular attraction with the liquid (eg. ammonia in water).

  • One mathematical expression of Henry's Law is: px = KH c

  • where px = the partial pressure of the dissolved gas X (i.e. px above the liquid), KH is Henry's Constant (which depends on the gas solute, the liquid solvent and the ambient temperature, and c is the concentration of the dissolved gas solute.

  • For relatively non–polar molecules of low solubility you do get a 'reasonably' linear graph of concentration (x) versus partial pressure (y)

  • with the axis origin 0,0 and this corresponds to the equilibrium

    • X(g) (c) doc b X(aq) where X is the gas dissolving in water.

 

WHAT NEXT?

Part 8 sub–index: 8.1 Vapour pressure origin and examples * 8.2.1 Introduction to Intermolecular Forces * 8.2.2 Detailed comparative discussion of boiling points of 8 organic molecules * 8.3 Boiling point plots for six organic homologous series * 8.4 Other case studies of boiling points related to intermolecular forces * 8.5 Steam distillation – theory and practice * 8.6 Evidence and theory for hydrogen bonding in simple covalent hydrides * 8.7 Solubility of covalent compounds, miscible and immiscible liquids

Advanced Equilibrium Chemistry Notes Part 1. Equilibrium, Le Chatelier's Principle–rules * Part 2. Kc and Kp equilibrium expressions and calculations * Part 3. Equilibrium and industrial processes * Part 4. Partition, solubility product and ion–exchange * Part 5. pH, weak–strong acid–base theory and calculations * Part 6. Salt hydrolysis, Acid–base titrations–indicators, pH curves and buffers * Part 7. Redox equilibria, half–cell electrode potentials, electrolysis and electrochemical series

 

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