Solutions

Solutions

1. solutions and solubility

   • a solution is a homogeneous mixture of a solute and a solvent
     • a solvent is a substance that will dissolve another substance
     • A solute is a substance that can be dissolved in a solvent
   • examples of solutions
     • solid in liquid
       • Salt water
     • Liquid in liquid
       • cordial in water
       • if 2 liquids mix and form a solution (a homogeneous mixture), these liquids are said to be miscible
     • Gas in liquid
       • soft drink

2. Concentration

   • ratio of solute:solvent
   • Concentration is defined as a quantity of solute dissolved in a fixed quantity of solution.
     • The concentration is NOT a measure of how much solution you have. It is a measure of how close the particles of solute are within the solution
   • when concentration is increased, the interparticle space between solute particles is reduced
   • Moles/L
     • C=n/v
       • n = number of moles
       • V = volume of solution
       • C = concentration (mol/L)
     • To change to grams/L:
       • m = n x M
   • Parts Per Million (ppm)
     • used to describe very low concentrations
     • ppm = mg (solute)/kg (solution)
     • x grams/L by 1000
   • ALL UNITS FOR CONCENTRATION ARE QUANTITY OF SOLUTE/QUANTITY OF SOLUTION

3. Dilutions

   • when water is added/removed to/from a solution, the number of moles of solute does not change
     • Concentration changes
   • solution before = solution after
     • C1V1 = C2V2
       • *only used when water is added or removed from a solution

4. Solubility rules

   • Insoluble salts
     • some compounds have forces that cannot be broken by water. these salts only dissociate a little in water.
       • e.g. AgCl is insoluble, when it is added to water and stirred, a very small number of ions dissociate and most of the solid is still as a solid
   • soluble salts
     • solid ionic compounds exist in the form of an ionic lattice
       • When an ionic compound dissolves, the ions in the lattice dissociate and become ions in the water
   • oxides are strongly insoluble

5. Precipitate forming reactions

   • precipitate forms when two soluble salts are mixed to produce an insoluble salt (precipitate)
     • e.g. when barium nitrate (aq) reacts with sodium sulfate, the barium ions collide with the sulfate ions in the solution, they will join together to form an insoluble precipitate
       • the salt produced in this reaction is an insoluble salt, which is indicated by (s)

6. saturated solutions

   • a saturated solution is one where no more solute will dissolve under the given conditions of temperature and pressure.
   • solubility can be described by making a saturated solution of a substance and then measure how much of the solute has dissolved
   • the solubility of a solid in a liquid depends on the temperature. In general most solids are more soluble as the temperature is increased
   • Unsaturated solution
     • an unsaturated solution can continually dissolve a solute under a given temperature/pressure
   • supersaturated solution
     • a supersaturated solution has dissolved more solute than expected under a given temperature and pressure

7. Solvents and solubility

   • When a solvent and solute are mixed in an attempt to form a solution:
     • force of attraction exists between solute particles
       • force of attraction exists between solvent molecules
         • must be a force of attraction between the solute and solvent molecules that is strong enough to:
           • ii) Separate the solvent molecules
           • i) Separate the solute molecules.
           • must be similar intermolecular forces and/or bonding between the solvent and solute particles ("like dissolves like")
             • can be soluble due to equivalence of intermolecular forces or insoluble due to the mismatch of intermolecular forces
           • SOLUBLE
             • solute particles are attracted to the water particles strongly enough to cause the solute particles to separate from each other and spread uniformly throughout the water
           • INSOLUBLE
             • the attraction between solute and solvent is not strong enough to separate solute-solute and solvent-solvent forces of attraction
           • POLAR solvents will dissolve POLAR solutes e.g. NH3, CH3CH2OH and C12H22O11 will readily dissolve in water
             • NON-POLAR solvents will dissolve NON-POLAR solutes. e.g. oil and wax dissolve in kerosene but not water
   • Temperature
     • When the temperature increases, the speed of both solvent and solute particles increases (direct)
       • Force of attraction between solute and solvent:
         • For ionic substances in water the solubility increases with temperature.
         • For molecular gases in water, the solubility decreases with temperature.
       • solubility of all gases in water decreases
   • Solutions, freezing and boiling points
     • ion-dipole forces are stronger than hydrogen bonding
       • ion-dipole interactions are the electrostatic force of attraction between polar molecules and ions. e.g. NaCl and H2O
     • the boiling point of water is dependent on the forces of attraction between the water molecules
       • e.g. if NaCl solid is added the boiling point of water increases. This is because the ion-dipole forces are greater, therefore requiring more energy to break resulting in a higher boiling point
     • the freezing point of salt water will be lower than pure water. this is because the large sodium ions make it more difficult for water particles to come close enough to form ice therefore, temp has to be decreased.
     • ions present directly effect the ability of water to change state

8. Unique properties of water

   • Melting/boiling point
     • Water has a high melting and boiling points compared with other covalent molecular substances.
       • The presence of hydrogen bonding between the water molecules means a larger amount of energy than usual is required to break the intermolecular forces between the water molecules. This energy is only available at higher temperatures.
   • Density of the solid phase compared with the liquid phase
     • In most substances the solid phase is denser than the liquid phase. Water’s liquid phase is actually denser that its solid phase. Consequently, ice floats on water.
       • when approaching the freezing temperature, molecules arrange themselves in a way which each molecule forms 4 hydrogen bonds to 4 neighbouring molecules
         • This creates a very open arrangement, meaning the water molecules in ice are more widely spaced. This is due to its v-shape structure which creates air pockets
   • Surface Tension
     • The surface tension of water is higher than of most covalent molecular substances
       • The attraction of surface molecules to ones below them have no opposite forces, therefore, they have an overall downwards force of attraction. The particles in the body of the liquid have no net force

9. Ionic equations

   • Step 1: write a balanced molecular equation
   • Step 2: split aqueous solutions into ions (strong acids/bases and soluble salts)
   • Step 3: Cancel out spectator ions
   • Write the ionic equation. This is the real reaction as it shows which chemical is reacting with which other chemical
   • observations
     • gases are colourless and odourless except: NO2, NH3, halogens
       • Halide - ion of halogen
         • all colourless in solution

10. Solubility curve