Chemical Equilibrium

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  • Chemical equilibrium and acid - base reactions.
    • Dynamic equilibrium
      • Macroscopic properties remain constant e.g temperature, concentration and pressure.
      • Closed system - nothing must be added, nothing must be lost.
      • Equilibrium can be reached from either direction
      • Dynamic at a molecular level - constant chnages occur, but these are in balance.
      • Rate of forwards reaction = rate of backwards reaction.
    • Le Chatelier's principle
      • If a system in dynamic equilibrium is subjected to a change, the position of equilibrium will shift to minimise the effect of the change.
      • This basically states that if you change something (temperature of the system/pressure of the system/concentration of one of the chemicals), then the system (reactants + products) will work to oppose that change.
      • Pressure changes will only affect reactions that involve gases.
      • We will use the Harber Process to explain Le Chatelier's principle.
        • A certain temperature 450 Degree Celsius - and pressure (~200atm) are used in the production of ammonia.
    • Changing the temperature of a system in dynamic equilibrium
      • The forwards reaction here is exothermic. This means that the forwards reaction leads to an increase in the temperature.
      • The backwards reaction must therefore be endothermic. The backwards reaction leads to a decrease in temperature.
      • Increasing Temperature
        • 1. System works to decrease temperature.
        • 2. Backwards reaction is promoted.
        • 3. Equilibrium shifts to the left hand side.
        • 4. Concentration of reactants increases.
      • Decreasing Temperature
        • 1. System works to increase temperature.
        • 2. Forwards reaction is promoted.
        • 3. Equilibrium moves to the right hand side.
        • 4. Concentration of reactants decreases.
      • Increasing Pressure
        • 1. System works to decrease pressure.
        • 2. Forwards reaction is promoted.
        • 3. Equilibrium shifts to the right hand side.
        • 4. Concentration of reactants decreases.
      • Decreasing Pressure
        • 1. System works to increase pressure.
        • 2. Backwards reaction is promoted.
        • 3. Equilibrium shifts to the left hand side.
        • 4. Concentration of reactants increases.
    • Acids & Bases
      • A base is something that accepts protons in aqueous solution.
        • HCl + NaOH   -> NaCl + H20
          • Neutralisation reaction.
      • Carbon Dioxide as an acidic gas
        • 1. CO2 + H20   H2CO3
        • 2. CO2 + H2O H+ + HCO3-
        • 3. HCO3- H+ + CO(3)2-
        • Using Le Chatelier's principle
          • For the first equilibrium - increasing levels of CO2 dissolving into the ocean lead to an increase in H+ (an increase in acidity) as equilibrium moves to the RIGHT hand side. This corresponds to a decrease in pH.
            • For the 2nd equilibrium - increasing acidity (H+) levels lead to a decrease in the level of carbonate present as equilibrium moves to the LEFT hand side.
      • An acid is something that donates protons, H+(aq) in aqueous solution.
        • Hydrochloric acid - HCL(aq) --> H+ (aq) + cl-(aq)
        • Nitric Acid HNO3(aq) --> H+(aq) + NO3-(aq)
        • Sulphuric acid- H2SO4(aq) -->2H+(aq) + SO4 2-
        • A base is something that accepts protons, H+ in aqueous solution.
          • HCL(aq) + NaOH(aq) ---> NaCl(aq) +H2O(l)
          • 2HNO3(aq) + Ca(OH)2 ---> Ca(NO3)2(aq) + 2H2O(l)
      • Acid-base titrations
        • 1. One reactant is measured with a graduated or volumetetric pipette and placed into a conical flask
          • 2. The other reactant is placed into a burette and is slowly run into the conical flask until the reaction is complete - often shown with an indicator. The flask should be swirled after each addition.
            • 3. The volume run out of the burette is known as the 'titre'
              • 4. Values should be recorded to an appropriate level of precision. Burettes usually measure to 0.10cm3, so it is not possible to be more precise than 0.05cm3. Values will always have 2.d.p - the final d.p being 0 or 5.
                • 5. Repeat until titre is constant.
                  • 6. Mean Titre is the average volume needed to completely react with the solution in the conical flask. (Calculated using consistent data).
      • Making a Standard Solution
      • Preparing a standard solution
        • The mass of the solid will be measured using a balance. To the beaker containing the solid, a small amount of distilled water will be added.
          • The contents will be stirred thoroughly to ensure that all soluble material has dissolved. This solution will be transferred to a volumetric flask. Distilled water will be used to swill out the residue from the beaker, into the volumetric flask.
            • Distilled water is added to the volumetric flask so that the solution reaches the mark (often 250cm3) The flask will be inverted several times to ensure thorough mixing of the contents.
              • Having 250cm3 of the solution will allow for making several repeats of the titration (as it is used to use 25cm3 samples each time)
  • 1. Measure mass of solid using a balance.
    • 2. Add a small amount of distilled water to the conical flask.
      • 3. Stir contents thoroughly so it dissolves.
        • 4. Transfer the resulting solution to a volumetric flask.
          • 5. Swill out residue in conical flask with distilled water into volumetric flask.
            • 6. Add distilled water to volumetric flask so solution reaches 'the mark' (normally 250cm^3)
              • 7. Invert the flask several times to ensure solution is thoroughly mixed.
    • Making a Standard Solution

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