Thermodynamics

• Standard enthalpy of solution


  Enthalpy change when 1 mol of an ionic compound is dissolved in solvent/ H20 to make aq ions.
• Enthalpy of hydration


  Enthalpy change when 1 mol of gaseousions is converted into aqueous ions
• What are the units for entropy?


  JK-1mol-1
• Definition for enthalpy change


  Heat change at a constant pressure

  Gases have high entropy due to random arrangement of particles
• How is disorder measured?


  JK-1mol-1

  Symbol (S)

  From a solid to aqueous ions= entropy is positive as there is more disorder

  From gas to liquid= entropy is negative as less disorder
• What does it mean when a reaction is feasible?


  Reaction can take place. Overall if G<or= to 0 then reaction is feasible.
• Why might a reaction that is feasible not take place?


  May have a high Ea that's too high to overcome.
• Favourable conditions for Gibbs free energy reaction


  1) Enthalpy change= Negative

  2) Entropy value= Positive
• What happens if you are dividing something positive by something negative?


  You're answer will be -ve. For example: 32,000 divided by -298 will = -107 jk-1mol-1 for entropy.
• What has zero entropy?


  Only perfect crystals at absolute zero (T = 0 K) will have zero ∆S entropy.

  
  

  Normally substances don't have zero entropy.
• How does the size and charge of the ion affect the size of the lattice enthalpy of an ionic solid?


  The smaller the ion and the higher its charge, the stronger the lattice
• What can you use to work out the enthalpy change of a solution?


  1)/\Hsolution = /\ HL dissociation + /\ hydH

  2) /\H solution = - /\HL formation + /\hydH
• How will a significant entropy change occur?


  1) If there is a change of state from solid or liquid to gas 2) There is a significant increase in number of molecules between products and reactants.
• What happens to the feasiblity of a reaction when increasing entropy?


  If the reaction involves an increase in entropy (∆S is +ve) then increasing temperature will make it more likely that ∆G is negative and more likely that the reaction occurs e.g. NaCl + aq Na+ (aq) + Cl - (aq)
• What happens to the feasibility of a reaction when decreasing entropy?


  If the reaction involves an increase in entropy (∆S is +ve) then increasing temperature will make it more likely that ∆G is negative and more likely that the reaction occurs e.g. NaCl + aq Na+ (aq) + Cl - (aq)
• What will happen to the feasibility of a reaction if delta S is close to 0?


  If the reaction has a ∆S close to zero then temperature will not have a large effect on the feasibility of the reaction as - T∆S will be small and ∆G will not change much e.g. N2 (g) + O2 (g) 2NO (g)
• What are hydration enthalpies?


  Hydration enthalpies are exothermic as energy is given out when water molecules bond to the metal ions.

  
  

  The higher the charge density the greater the hydration enthalpy (e.g. smaller ions or ions with larger charges) as the ions attract the water molecules more strongly

  
  

  e.g. Fluoride ions have more negative hydration enthalpies than chloride ions.

  Magnesium ions have a more negative hydration enthalpy than barium ions.
• What does it mean when a substance is insoluble?


  If a substance is insoluble it is often because the lattice enthalpy is much larger than the hydration enthalpy and it is not energetically favourable to break up the lattice, making ΔH solution endothermic.
• Equation for atomisation enthalpy for

  1) Bromine

  2) Sodium


  1) 1/2 Br2 (l) = Br (g)

  2) Na (s)= Na (g)
• What determines the charge on an ion?


  Elements combine to make the compound which is as stable as possible - the one in which the greatest amount of energy is evolved in its making.
• Why is MgCl2 formed when Mg reacts with chlorine


  1) Mg+ and Cl- would form weak attractions and so overall lattice enthalpy would be slightly negative

  2) Mg2+ and Cl2- means more energy to ionise magnesium and a lot more energy released when lattice forms because stronger attraction between Mg2+ and Cl2-. More negative

  3)Mg3+ and Cl3- means more energy needed to ionise Mg. Higher enthalpy of formation energy. But 3rd electron closer to nucleus so more energy needed to remove. Less negative than 2-. Energetically unstable/ thermally unstable.
• How can a reaction occur even though it is endo?


  Lower Ea than expected, so easier to overcome
• What happens to the Ea when temperature is increased?


  The Ea also increases.
• What do you talk about when you get a question on comparison between a calculated value and cycle value?


  1) Is there covalent character there?

  2) Which value is bigger or smaller?

  3) What does the ionic model assume? Does it support or refute?
• What increases entropy?


  1) Entropy higher for weakly bonded compounds compared to compounds with strong covalent bonds.

  2) Entropy increases as mass of a molecule increases

  3) Increases as complexity of a molecule increases (number of atoms/number of heavier atoms)
• Using a Born-Haber cycle, a value of -905 kJ mol-1 was determined for the lattice enthalpy of silver chloride. A value for the lattice enthalpy of silver chloride using the ionic model was -833 kJ mol-1 . Explain what a scientist would be able to deduce from a comparison of these values.


  1)The model used assumes the ions are spherical and in a lattice and attractions purely electrostatic

  2)The calculated value is smaller than the cycle value or stronger attraction

  3)Indicating some covalent character or ions are polarised
• Calculate the minimum temperature at which this reaction is feasible.


  1)For a reaction to occur ΔG < 0

  2) ΔS is positive and large as a gas is evolved 1

  3) TΔS is larger than ΔH and ΔG is negative

  4) Delta S may be a gas which is more positive than solids/ liquids and has more entropy.
• How do you work out the lattice dissocation enthalpy of a substance?


  ∆H(solution) = LE + Σ(hydration enthalpies)
• Deta G equation


  Detla G= or smaller than 0
• Enthalpy of atomisation


  Enthalpy change when 1 mole of gaseous atoms is produced from an element in it's standard state.
• Enthalpy of hydration


  Enthalpy change when 1 mol of gaseous ions is turned into 1 mol of aqueous ions by becoming hydrated (dissolved in water).
• Enthalpy of solution


  Enthalpy change when 1 mol of an ionic lattice/solid is dissolved in water to produce 1 mole of aqueous ions.
• Bond dissocation enthalpy


  Enthalpy change when 1 mole of a covalent bond is broken down into it's gaseous phase/ stateof2atoms(2cl= CL2, not always 2 tho, depends on molecule)
• Lattice enthalpy of formation


  Enthalpy change when 1 mole of an ionic solid is formed from its constituent gaseous ions. (not standard states because not atoms)
• Lattice enthalpy of dissocation


  Enthalpy change when 1 mole of an ionic solid is broken down into it's constituent gaseous ions.
• Enthalpy of vaporisation


  Enthalpy change when 1 mole of liquid turned into gas.
• Enthalpy of fusion


  Enthalpy change when 1 mol of solid turned into liquid.
• What is delta G equal to for a reaction to be feasible?


  Delta G equal to or < than 0.
• What do we need the temperature to be equal to for reaction to be feasible?


  T > or equal to temperature in K
• y=mx+c equation for gibbs free energy


  y= -deltaST+delta H