Electrostatic force of attraction between an ion and a polar molecule
Formed from the electrostatic force of attraction between an ion and a dipole
Note: “Intermolecular force” is a misnomer, as this is in ionic substances, not covalent substances
Stronger than permanent dipole-dipole and dispersion forces
Ion – dipole forces form between ions of an ionic compound and molecules of water. We say the ions are hydrated.
If the ion – dipole forces formed between ions and water are strong enough, they can pull ions off the surface of an ionic lattice, and the ionic substance is soluble.
Not all ionic substances dissolve. In insoluble substances, the ionic bonding is too strong, and the ion – dipole bonds formed are not strong enough to break up the ionic lattice.
Solubility
Soluble or Insoluble?
Solutions form when solvent and solute mix
Process involves breaking bonds/attraction in these individual substances, which is endothermic, thus it requires energy
In the solution, attractions form between the solvent and solute, and this bond formation is exothermic, releasing energy
Whether the solution forms overall depends on whether the overall process is energetically favourable
i.e. if the amount of energy released from forming solute-solvent bonds is greater than the amount of energy breaking solute-solute and solvent-solvent bonds
Dissolving can be exothermic or endothermic, but if too much energy is required, it won’t happen
Like dissolves like (do not use this phrase)
If the solvent and solute are able to form good intermolecular bonds with each other, this will favour dissolving
Polar solvents and polar solutes can form dipole-dipole interaction (and sometimes H-bonds), which favours dissolving
Non-polar substances cannot form dipole-dipole interactions or h-bonds, so won’t mix with polar substances
Non-polar solutes dissolve well in non-polar solvents. All substances form only dispersion forces, meaning no strong bonds are broken
Soluble
Insoluble
When a solute is soluble in a specific solvent
When a solute is insoluble in a specific solvent
The intermolecular forces between the solvent and solute are stronger than the intermolecular forces present between the solvent’s atoms, and the solute’s atoms
The intermolecular forces between the solvent & solute are weaker than the intermolecular forces present between the solvent & solvent and between the solutes
This means the difference between the energy released in the formation of the intermolecular forces between the solvent & solute and the energy required to break the solute – solute and solvent – solvent intermolecular forces decreases
This means the difference between the energy released in the formation of the intermolecular forces between the solvent & solute and the energy required to break the solute – solute and solvent – solvent intermolecular forces is large
Thus the solvent – solvent and solute – solute IMF are able to be disrupted and the solute – solvent IMF are able to be established as this process is energetically favorable
Thus the solvent – solvent and solute – solute IMF are not able to be disrupted as this process is energetically unfavorable and so the solute does not dissolve.
Water - The Universal Solvent
As a highly polar molecule, water is great at dissolving polar molecule substances and many ionic substances
Water can form a range of intermolecular forces to stabilise a solution
Dipole-dipole forces with polar molecules
Hydrogen bonds with hydrogen bonding polar molecules
Ion-dipole forces with ions from ionic compounds
Aqueous Solutions of Covalent Compounds
Many polar molecules are soluble in water due to the strong dipole-dipole forces or hydrogen bonds formed
Organic compounds are usually non-polar, but if they have –OH (alcohol) or –NH (amine) groups, they can hydrogen bond, so this will increase their solubility in water
Acids are molecular substances that ionise in water to form ions. This means they can form ion – dipole forces with water, and so are usually very soluble
Solubility of ionic substances
All ionic substances are soluble to some extent, so we categorise substances by the concentration of a saturated solution at 25˚ C.
A substance is described as soluble if a saturated solution is 0.1 mol L-1 or higher.
A substance is described as slightly soluble if a saturated solution is 0.01 mol L−1 to 0.1 mol L−1.
A substance is described as insoluble if a saturated solution is below 0.01 mol L-1.
Precipitation
Precipitation Reactions
An insoluble solid that forms during an aqueous reaction is called a precipitate.
A reaction which forms a precipitate is called a precipitation reaction
The limewater test for carbon dioxide is a precipitation reaction
Limewater is actually a dilute solution of calcium hydroxide
The calcium hydroxide reacts with carbon dioxide to form calcium carbonate, which is insoluble in water
Uses of Precipitation Reactions
Most precipitation reactions are very fast reactions that occur between ions.
This makes them very useful for identifying specific ions based on the type of precipitate formed.
Precipitation reactions have a number of other uses:
Production of coloured pigments for paints and dyes
Removal of toxic chemicals from water
Separation of reaction products.
Spectator Ions
In ionic precipitation reactions there are often ions that are not involved in the reaction. These are known as spectator ions.
The spectator ions are easily identified using the ionic equation, where they are the ions that don’t change their state of matter, or don’t participate in the reaction
You should always write net ionic equations, omitting spectator ions, in this course.
Identifying Positive Ions
Many metal ions can be identified by flame tests
However, precipitation reactions can also be used to identify positive ions, as many metals form hydroxide precipitates with characteristic colours
Identifying negative ions: Halides
Formed from group 17 elements, the halogens
Detected using silver nitrate solutions
Colourless ions in a solution
I−,Cl−,Br−
The different silver halide precipitates can be distinguished by their differing colours
Chloride: white AgCl precipitate
Bromide: cream AgBr precipitate
Iodide: yellow AgI precipitate
Isolating the Precipitate
Precipitate from a precipitation reaction can be separated from the reaction mixture by filtration
A Buchner funnel and flask can be used to accelerate the process
This apparatus uses a vacuum pump to draw the mixture through the filter
The filtrate is finally washed and dried
More Solubility
Solubility
Concentration of a saturated solution
Saturated: Solution where no more solute will dissolve at the given temperature
Unsaturated: Less solute than needed to make a saturated solution. Most solutions we use will be unsaturated
Supersaturated: Unstable solution, usually formed as a saturated solution curves down. If it is disturbed, it will crystallise, forming crystals of solute, and leave a saturated solution
Solubility: Mass of a substance that will dissolve in 100 g of water at 25˚ C
Measured in g/100g
The solubility of a substance depends on the temperature of the solvent
Solubility Curves
Solubility curves show the solubility of substances as temperature changes
Generally, the solubility of ionic solids increases with temperature
There is more energy to overcome bonding in the solid
Saturation and Crystallisation
Solutions become supersaturated when they contain more solute than they would normally be able to dissolve
This can happen when a saturated solution is cooled, or another change in conditions occurs that causes solubility to decrease
The solute will stay in solution until a ‘seed’ crystal is added.
This causes it to crystallise out of the solution very quickly.
When it does this, it gives out heat energy.
Solubility of Gases
Many gases are soluble in water.
For example, fish can breathe because of the oxygen dissolved in water.
Solubility of gases in water is often low, as most gases are not very polar. More polar gases are more soluble, in general.
As temperature increases, solubility of gases decreases
This is because the gases have more kinetic energy, which increases motion and breaks the IMF between the gases and the solvent, causing them to escape the solution
