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Revision notes on energetics - enthalpy changes associated with dissolving & lattice enthalpy - for Advanced A/AS Level Theoretical-Physical Chemistry:
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Part 2 ΔH Enthalpy Changes contd. Lattice Enthalpy, Enthalpies of Solution, Enthalpies of Ion Hydration 2.1ac Important enthalpy definitions and what happens when an ionic compound dissolves in water and the energetics of why? What energy changes are associated with an ionic solid dissolving in water? A particle model is used to illustrate the dissolving of an ionic solid in water and the subsequent hydration of the free ions. The enthalpy of solution, lattice enthalpy and the hydration enthalpies of gaseous ions are then introduced and are all connected using a thermochemical cycle to theoretically explore 'some' of the factors that affect how soluble a salt/oxide etc. is soluble, i.e. does it dissolve appreciably or is it insoluble. |
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2.1 What happens when an ionic compound dissolves in water and the energetics of why? Section 2.1 discusses a dissolving enthalpy cycle another application of Hess's Law in which the process of dissolving an ionic compound such as a halide salt or a metal oxide, is broken down into theoretical stages to help understand the factors involved in deciding whether of not a substance will dissolve readily or not at all. For the moment, only enthalpy changes will be considered but eventually entropy changes must be discussed too!
2.1a The changes that occur in the dissolving process Despite the strong ionic bonding forces in most salts or simple binary compounds like oxide or chloride crystals i.e. the strong electrostatic attraction between positive ions (cations) and negative ions (anions) many ionic compounds readily dissolve in water. Therefore, not surprisingly, a great deal of energy is required to separate the ions, but dissolving can still take place. So how can we explain this?
In some cases cations become hydrated via dative covalent bonds to form a complex ions e.g. Li(H2O)4]+, Cu(H2O)4]2+, Mg(H2O)6]2+, Al(H2O)6]3+ etc. because in the case of water, the most electronegative part of the highly polar water molecule (>Oδ) will be attracted to the positive ion and, since the oxygen atom has two lone pairs of electrons, it is also the source of the dative covalent bond by donation of one of these pairs of electrons into a vacant metal ion orbital. In the case of anions, the positive ends of the water molecules (Hδ+) will orientate themselves towards the negative anion and the water molecules become weakly associated with anion, but no covalent bonds are formed. This solvation of the ions means the ions are effectively bigger particles which makes the distance between the positive and negative ion centres greater, and by the laws of electrostatics, the attractive forces is weakened and the hydration process is always exothermic. PLEASE note that dissolving a solute in this situation cannot be simply regarded as a physical change. The ionic lattice is broken down, but NOT by melting, and chemical bonds are formed between the cation and water molecules to form hydrated ions or aquaions such as those listed above. 2.1b Diagram illustrating the dissolvingsolvationhydration process for sodium chloride crystals forming salt solution
In order to try to understand processes you can use a Hess's Law cycle to break the process down into theoretical steps, each of which can measured experimentally or theoretically calculated. The relative magnitude of each energy change can help understand why a substance will dissolve or be insoluble. However, this cycle only uses enthalpy values and excludes entropy changes which I'll deal with later. 2.1c The connection between lattice enthalpy, enthalpies of ion hydration and enthalpy of solution (i) The energy change for a substance dissolving in a solvent is called the enthalpy of solution.
Comments on the enthalpy of solution values
You can derive an alternative route to form the solution employing Hess's Law ... (ii) The ionic crystal lattice is vapourised into its gaseous positive ion (cation) and gaseous negative ion (anion) constituents.
(iii) The gaseous ions interact with water to give the hydrated ions in aqueous solution.
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