Enthalpy Changes, Reaction Rates + Chemical Equilibrium

• Activation energy (Ea)


  The minimum amount of energy required to start a chemical reaction
• Standard enthalpy change of combustion


  The enthalpy change when one mole of a substance is burnt in excess oxygen, under standard conditions in their standard states
• Standard enthalpy change of formation

  The enthalpy change when one mole of a compound is formed from its constituent elements, under standard conditions in their standard states
• Standard enthalpy change of neutralisation

  The enthalpy change when one mole of water is formed in a neutralisation reaction under standard conditions, in their standard states
• Spirit burner experiment (calorimetry)

  Liquid fuels (such as methanol) can be easily burnt using small spirit burners:

  
  

  1) Using a measuring cylinder, measure out 150cm^3 of water. Pour water into a beaker. Record the initial temperature of the water to the nearest 0.5 degrees Celsius.

  2) Add methanol to the spirit burner. Weigh the spirit burner containing methanol

  3) Place the spirit burner under the beaker. Light the burner and the methanol whilst stirring the water with the thermometer

  4) After 3 minutes, extinguish the flame. Immediately record the maximum temperature reached by the water

  5) Re-weigh the spirit burner containing methanol. Assume the wick has not been burnt.

  Possible problems --> Incomplete combustion (less energy given out), heat losses, (non-standard conditions)

  
• Average bond enthalpy


  The energy required to break one mole of a specified type of bond in a gaseous molecule
• Exothermic - Bond enthalpy


  Energy released when forming bonds is greater than the energy required to break bonds

  
• Endothermic - Bond enthalpy

  Energy required to break bonds is greater than the energy released when forming bonds

  
• Why might the actual bond enthalpy differ from the average value?


  Bonds may be in different environments/Incomplete combustion/Heat loss/Non-standard conditions
• Why do Br2 and I2 not exist in the gaseous state under standard conditions?

  They have London forces between the molecules (which energy is required to break them)
• What 2 conditions are needed for an effective collision?

  1) Particles must collide with energy greater than/equal to the activation energy

  2) Particles must collide with correct orientation
• Rate of chemical reaction

  How fast reactants are being used up/how fast products are being formed
• Rate of chemical reaction - Explanation

  1) The rate of reaction is fastest at the start of the reaction, as the concentration of reactants is at its highest

  2) The rate of reaction slows down as the reaction proceeds, as the reactants are being used up and so the concentration decreases

  3) When one of the reactants have been completely used up, the concentration stops changing and rate of reaction is 0
• Effect of increasing concentration of the rate on reaction


  The number of particles will increase while occupying the same volume. This brings the particles closer together and they collide more frequently. Given that the particles have sufficient energy and correct orientation, effective collisions increases (increasing the rate of reaction)
• Effect of increasing pressure on the rate of reaction

  Increasing the pressure increases the concentration of gas molecules as the number of the gas molecules occupy a smaller volume. The gas molecules are closer together and so collide more frequently. Given that the gas particles have sufficient energy and correct orientation, effective collisions increase (increasing the rate of reaction)
• Monitoring the production of a gas using gas collection (e.g. decomposition of H2O2)

  1) Hydrogen peroxide is added to a conical flask and the bung is replaced

  2) Initial volume of gas is measured in the measuring cylinder

  3) Manganese dioxide (MnO2) is catalyst and is quickly added to the conical flask. The bung is then replaced and start a stopwatch

  4) Volume of gas produced in the measuring cylinder is recorded at regular intervals

  5) Reaction is complete when no more gas is produced

  2H2O2(aq) (MnO2 - catalyst) ---> 2H2O + O2(g)

  
• What is a catalyst?

  Substance that increases the rate of reaction without being used up in the overall reaction

  
• Role of catalyst


  - Can react with a reactant to form an intermediate

  - Provide a surface on which a reaction can take place

  - Increases the rate of a chemical reaction by providing an alternative pathway of lower activation energy

  
• Homogeneous catalyst


  A catalyst that has the same physical state as reactants

  
  

  For example, making esters with sulfuric acid (they are in the same physical state - liquids) or ozone depletion - the reactant (O3) and catalyst (Cl radical) are both gases
• Heterogeneous catalyst


  A catalyst that has a different physical state to reactants.

  
  

  They are usually solid and absorb reactant molecules onto their surface
• Economic importance and benefits of catalysts

  - Reduces temperature needed for industrial chemical reactions

  - Requires less energy, electricity or fossil fuels (increases profitability)
• Environmental importance and benefits of catalyst


  Using less fossil fuels cuts carbon dioxide emissions
• Boltzmann distribution


  The shaded area of the Boltzmann distribution show that only a small proportion of molecules have more energy than the activation energy

  
  

  - No molecules have 0 energy - the curve starts at the origin

  - The area under the curve is equal to the total number of molecules

  - There is no maximum energy for a molecule - the curve never reaches the x-axis

  
• Boltzmann distribution - temperature (features)

  As the temperature increases, kinetic energy of the molecules increases. The peak of the graph is lower on the y-axis and more to the right/further along the x-axis. The number of molecules stays the same (area under curve also remains the same)
• Boltzmann distribution - temperature

  As temperature increases, a greater of proportion of molecules will have energy greater than/equal to the activation energy. Collisions will be more frequent as the molecules are moving fast. There will be a greater proportion of effective collisions increasing the rate of reaction

  
• Boltzmann distribution - catalyst


  A catalyst provides an alternative pathway with a lower activation energy. This means compared to the Ea, a greater proportion of molecules will exceed the lowered activation energy (Ec). More molecules will react to form products, increasing the rate of reaction

  
• Dynamic Equilibrium


  1) The rate of the forward reaction is equal to the rate of the reverse reaction

  2) The concentrations of reactants and products do not change

  
  

  The picture on the right shows that there is an increase in the concentration of products, until equilibrium is reached and the concentration remains constant (the blue line shows a decrease in concentration of reactants until equilibrium is established - the concentration doesn't go to 0, there is still reactant)

  
• Le Chatelier's Principle


  When a system is subject to a change in temperature, pressure or concentration. The position of equilibrium will shift to counteract this change
• Effect of concentration on the position of equilibrium

  Increase in concentration - The position of equilibrium will shift to the right to reduce the effect of the increase in concentration of reactants

  Decrease in concentration - The position of equilibrium will shift to the left to reduce the effect of decrease in reactant (or increase in concentration of product)

  
• Effect of pressure on the position of equilibrium (only if reactant/product are gases)

  Increase in pressure - The position of equilibrium will shift to the side with the least moles of gas to decrease the pressure

  Decrease in pressure - The position of equilibrium will shift to the side with the larger moles of gas to increase the pressure

  
• Effect of temperature on the position of equilibrium

  Increase in temperature - The position of equilibrium will shift to the endothermic direction (forward reaction)

  Decrease in temperature - The position of equilibrium will shift to the exothermic direction (backward direction)

  
• Effect of catalyst on the position of equilibrium

  Catalysts increase the rate of forward and reverse reaction equally. It does not affect the position of equilibrium once reached. It only causes a reaction to reach equilibrium faster
• Investigating changes to the position of equilibrium with concentration

  The equilibrium between aqueous chromate ions (Cr2O42-) and aqueous dichromate ions (Cr2O72-) is sensitive to changes in acid concentration. Solutions of chromate ions turn yellow and solutions of dichromate ions turn orange

  2Cr2O42- + 2H+ --> Cr2O72- + 2H20

  Increasing the acid concentration --> Adding dilute sulfuric acid causes the position of equilibrium to shift to the right. This decreases the concentration of H+ ions and creates more products (Cr2O72-) which turns the solution more orange

  Decreasing the acid concentration --> By adding aqueous sodium hydroxide. This shifts the position of equilibrium to the left. This increases the concentration of H+ ions and creates more reactant (Cr2O42-) which turns the solution more yellow

  
• Investigating changes to the position of equilibrium with temperature

  Cobalt chloride, CoCl2, dissolves in water to form a pink solution.

  Add water to cobalt chloride in a boiling tube and add a small quantity of hydrochloric acid. Transfer the boiling tube to an ice bath (solution is pink). Then transfer the boiling tube to a boiling water bath - solution is blue.

  Forward reaction = Endothermic

  Backward reaction = Exothermic

  Increasing the temperature of the solution increases the amount of heat energy in the system. Equilibrium shifts to the right favouring the endothermic reaction as the solution turns more blue - more CoCl42-.

  Decreasing the temperature of the solution decreases the amount of heat energy in the system. Equilibrium shifts to the left favouring the exothermic reaction as the solution turns more pink - more [Co(H2O)6]2+

  
• Operational Conditions (Haber Process)


  Le Chatelier's Principle is used to get the best yield of ammonia:

  - Increasing pressure will shift the position of equilibrium to the side with the fewest moles of gas to reduce the pressure. The position of equilibrium shifts to the right and the yield of ammonia increases. However, very high pressures are expensive and so a compromise of 200atm is used,

  - Using a low temperature would shift the position of equilibrium to the right and produce a high yield of ammonia (the forward direction is exothermic - decrease temperature), However, it would produce the product very slowly. So a compromise of 450 degrees Celsius is used (300-500)

  
• Safety and Economic issues concerning Haber Process

  - Very high pressure requires a strong container and a large quantity of energy to generate, so would be very expensive.

  - Safety is also a concern, which could endanger workforce and the surrounding area.

  - If low temperature is used, it would be a slow process and give a very slow rate of reaction
• The equilibrium constant (Kc)

  
• Estimating the position of equilibrium from Kc

  - If Kc is very large (Kc>1), the position of equilibrium lies to the RHS - reaction mixture will mostly contain products

  - If Kc is very small (Kc<1), the position of equilibrium lies to the LHS - the reaction mixture will mostly contain reactants

  - If Kc is close to 1, the mixture will contain a similar concentration of reactant and products
• Calculating enthalpy change of neutralisation + Cooling curves


  1) Pipette 25cm^3 of 1.00moldm-3 of hydrochloric acid into a polystyrene cup (measure with measuring cylinder first)

  2) Then pipette 25cm^3 of 1.00moldm-3 of sodium hydroxide. DO NOT ADD TO THE BEAKER YET

  2) Start a stop-clock and record the initial temperature of the acid. Record the temperature at 1m, 2m and 3m (stir with thermometer).

  3) At the 4thm add sodium hydroxide to the acid. Record the temperature to the nearest 0.5 at 5m, 6m, 7m, 8m, 9m and 10m.

  4) Plot a graph of temperature against time. To correct for cooling, extrapolate the cooling curve section of the graph back to when zinc was added. Draw a vertical line from the time that the solutions were mixed to the extrapolated cooling curve