CHEMISTRY 5

CHEMISTRY 5

1. STRONG AND WEAK ACIDS

   • STRONG
     • Fully ionise in solution
       • Lots of H+ ions
     • Hydrochloric
       • HCl --> H+ + Cl-
     • Nitric
       • HNO3 --> H+ + NO3-
     • Sulphuric
       • H2SO4 --> 2H+ + SO4 2-
   • WEAK
     • Partially ionise in solution
       • A few H+ ions
     • Ethanoic
       • CH3COOH <=> CH3COO- + H+
   • Why does HCl react faster & more vigorously with Mg than CH3COOH?
     • HCl has lots of H+ ions: lots of collisions between HCl and Mg, therefore reaction is more vigorous and the Mg reacts faster
   • THE DIFFERENCE IN CONDUCTIVITY
     • In order for any chemical to conduct, it must have freely moving charged particles
       • Strong acids have lots of H+ ions in solution, and so conduct better than weak acids, which only have a few H+ ions
   • Explain why the pH of a weak acid is much higher than the pH of a strong acid of the same concentration.
     • The higher the concentration oh hydrogen (H+) ions, the lower the pH
   • Descalers
     • Scale is calcium carbonate that forms on the element inside kettles
       • This prevents the element from heating the water as effectively.
     • Descalers are weak acids that react with the scale
       • A strong acid should not be used because it would react with the scale AND the metallic element
       • Sulphamic acid is a suitable descaler

2. IONIC EQUATIONS AND PRECIPITAION

   • PRECIPITATION
     • A precipitate is an insoluble solid formed when 2 solutions are mixed together.
     • Testing metal ions with sodium hydroxide
       • Cu2+ = blue precipitate
         • copper sulphate
           • CuSO4
         • +
         • sodium hydroxide
           • 2NaOH
         • -->
         • copper hydroxide
           • Cu(OH)2
         • +
         • sodium sulphate
           • Na2SO4
         • The copper ions and the hydroxide ions have a very strong attraction for one another, this is why they form the precipitate.
       • Fe2+ = green precipitate
       • Fe3+ = rust precipitate
     • Why is it so fast?
       • Ions are spread completely through each solution, so ions from the two solutions collide and react - immediately the liquids mix
         • As the ions are also moving rapidly, the frequency of the collisions is high
           • High collision frequency = quick reaction
     • PREPARING A CLEAN DRY SAMPLE
       • Step 1: MIX
         • Make the preciipitate
       • Step 2: FILTER
         • Filter the precipitate from the solution through filter paper - removes soluble impurities
       • Step 3: WASH
         • Distilled water is poured over the precipitate to remove traces of solution
       • Step 4: DRY
         • Leave the precipitate in a warm place on an evaporating dish to evaporate the water
     • TESTING HALIDES WITH LEAD NITRATE
       • Chloride = white precipitate
       • Bromide = cream precipitate
       • Iodide = yellow precipitate
       • Barium chloride can be used to test for sulphate ions
   • IONIC EQUATIONS
     • Only the ions that react with each other to form the precipitate are included: spectator ions left out
       • Spectator ions are the ions which don't take part in the reaction and are still moving around in solution at the end

3. GAS VOLUMES

   • Gases produced in a reaction can be collected in a gas syringe, measuring cylinder or burette, or measured using a balance (decrease in mass = gas released)
     • The gas syringe is the best method because all gases are contained in the apparatus and none can escape
   • RATE CURVES
     • Steepest at the start because no particles have been used yet, more collisions & more chance of successful collisions = higher rate
       • Towards the middle: less steep: more particles have been used up = slower rate
         • Levels off at the end: one of the chemicals has been completely used up = no more collisions
     • LIMITING REACTANT: chemical that is used up
     • REACTANT IN EXCESS: still there at the end of the reaction
     • If any FACTOR that affects the RATE of the reaction is altered, then the STEEPNESS is altered
       • Temperature, surface area, concentration,
     • Where the graph LEVELS OFF depends on the number of PARTICLES used

4. EQUILIBRIA

   • The concentration of products and reactants stays constant & the forward and backwards rates are equal
   • If any ccndition is altered, the equilibrium moves to oppose the change
     • Think of it like a SEESAW
   • CONCENTRATION
     • A + B <=> C + D
       • If you increase the concentration of A, the equilibrium moves to the right to use up the A that has been added
       • If you decrease the concentration of D, the equilibrium moves to the right to replace what has been lost
     • If a chemical is added, that side of the equilibrium is heavier, so it moves to the other side to USE UP the added chemical
     • If a chemical is removed, that side of the equilibrium is lighter, so the equilibrium moves to that side to replace the chemical lost
   • PRESSURE
     • No. of particles: greater no. of particles = greater pressure
     • IF PRESSURE IS INCREASED: the equilibrium moves in the DIRECTION of the FEWER particles to REDUCE PRESSURE and OPPOSE the CHANGE
     • IF PRESSURE IS DECREASED: the equilibrium moves in the DIRECTION of MORE particles to INCREASE PRESSURE and OPPOSE the CHANGE
   • TEMPERATURE
     • EXOTHERMIC = NEGATIVE DELTA H:
       • If temp INCREASED, equilibrium moves to the LEFT, to TAKE IN HEAT and oppose the change
       • If temp DECREASED, equilibrium moves to the RIGHT, to GIVE OUT heat and oppose the change
     • ENDOTHERMIC = POSITIVE DELTA H
       • If temp INCREASED, the equilibrium moves to the RIGHT, to GIVE OUT heat and oppose the change
       • If temp DECREASED, the equilibrium moves to the LEFT, to TAKE IN heat and oppose the change
     • If the temp is increased, it will always go in the endothermic direction
   • THE CONTACT PROCESS
     • Making sulphuric acid from sulphur, air and water: useful in fertilisers, paints & dyes.
     • 1) S + O2 --> SO2
       • 2) 2SO2 + O2 <=> 2SO3
         • 3) SO3 + H2O --> H2SO4
         • Most important step: determines amount of product
     • increasing the temperature moves the position of equilibrium to the left and increases rate of reaction so a compromise temperature of 450 degrees C is used
     • addition of Vandadium (V) oxide catalyst increases rate but does not change position of equilibrium
     • even at low pressure, the position of equilibrium is already on right so expensive high pressure is not needed.