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Lattice Enthalpies

  • Enthalpy Change= The heat energy transferred in a reaction at constant pressure. 
  • Lattice enthalpy is defined as either Lattice Formation Enthalpy, or Lattice Dissociation Enthalpy. 
  • Lattice Formation Enthalpy= The enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions. 
  • Lattice Dissociation Enthalpy= The enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its gaseous ions under standard conditions. 
  • The purely ionic model assumes that all ions are spherical, and have their charge evenly distributed around them. 
  • The experimental lattice enthalpy, from the Born-Haber cycle is evidence that ionic compounds usually have some covalent character. 
  • Ions in a lattice aren't usually exactly spherical.
  • Positive ions polarise neighbouring negative ions to different extents. 
  • The more polarisation there is, the more covalent the bonding will be. 
  • If the experimental lattice enthalpies are a lot bigger than the theoretical lattice enthalpies, it tells you that the bonding is stronger than the calculations from the ionic model predict. The differnce shows that the ionic bonds are quite strongly polarised, so they have quite alot of covalent character.
  • If the experimental and theoretical values are close together, the structure of the lattice for this compound is quite close to being purely ionic (almost no polarisation).
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Enthalpies of Solution

  • When a solid ionic lattice dissolves in water:

(1) The bonds between the ions break (endothermic) (Lattice enthalpy of dissociation).                   (2) Bonds between the ions and water are made (exothermic) (Enthalpy change of hydration).       (3) Enthalpy change of solution is the overall effect on the enthalpy of these two things. 

  • Enthalpy chnage for a reaction= sum of enthalpies of bonds broken- sum of enthalpies of bonds formed. 
  • Mean bond enthalpies are only approximations. This is because a given type of bond will vary in strength from compound to compound. This value can even vary within a compond. 
  • Mean bond enthalpies are avaerages of these bond enthalpies. 
  • Only the bond enthalpy of diatomic molecules will always be the same. 
  • You get more exact results for working out mean bond enthalpy, from experimental data from the specific compounds. 
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Entropy Change

  • Entropy= a measure of the number of ways that particles can be arranged and the number of ways that the energy can be shared out between the particles. 
  • The symbol for entropy is S.
  • Substances really like disorder, so the particles move to increase the entropy. 
  • Things that effect entropy are:
  • (1) Physical state- In solid compounds, particles remain in fixed positions, and only vibrate around this point, therefore disorder is very low and so is entropy. Whereas gaseous compounds have very random arrangements, so they have more disorder and so the highest entropy. 
  • (2) Dissolving- Dissolved particles can move freely, and are no longer held in a fixed position, therefore there is a lot of disorder and the highest entropy. 
  • (3) Quantity of particles- The greater the number of particles you have, the more ways they and their energy can be arranged, meaning the entropy increases as the number of moles increase.
  • Spontaneous reactions are ones that you don't have to supply energy to. These can be endothermic or exothermic. 
  • This occurs because the entropy increases such a lot that the reaction occurs by itself, without you supplying any energy. 
  • Examples of this are reactions where there is a change in state to increase the entropy (evaporation of water), or where there is a greater number of moles of products than reactions (sodium hydrogencarbonate with hydrochloric acid). 
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Entropy Change Calculations

  • In a reaction there is the entropy change between the reactants and products (called the entropy change of the system) and the entropy change of the surroundings (because energy is transferred from or to the system). 
  • DeltaS= DeltaS (system) + DeltaS (surroundings)
  • DeltaS (system)= S (products) - S (reactants)
  • DeltaS (surroundings)= - (DeltaH/ T)
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Free- Energy Change

  • Free energy change (Delta G) is a measure used to predict how feasible a reaction is. 
  • If the free energy change is negative or equal to zero, then the reaction might happen by itself. 
  • If free energy change shows that a reaction is feasible, it may have such a high activation energy and be so slow that you wouldn't notice it happening at all. 
  • DeltaG= DeltaH (in Jmol-1) - (T (in K) x DeltaS (system) (in JK-1mol-1))
  • If a reaction is exothermic and has a positve entropy change, then the free energy change is always negative- these reactions are feasible at any temperature. 
  • If a reaction is endothermic and has a negative entropy change, then the free energy change is always positive- these reactions are not feasible at any temperature. 
  • If DeltaH is positve and DeltaS (system) is postive the reaction won't be feasible at some temperatures, but will be at higher temperatures, as then they result in an increase in entropy
  • If DeltaH is negative and DeltaS (system) is negative then the reaction will be feasible at some temperatures but not at higher temperatures, as it results in a decrease in entropy
  • T= DeltaH/ DeltaS (system)
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