CHEMISTRY SPECIFICATION NOTES
OCR Chemistry Salters Nuffield B AS/A-Level
Key:
Definitions are in turquoise
Core practicals are in orange
Equations and tests/reactions in green
Topic 8: Oceans (O)
(a) the factors determining the relative solubility of a solute in aqueous and non-aqueous
solvents. Intermolecular bonds, ion–dipole bonds and ionic bonds should be considered.
In ionic solids, ions are held together by their opposite electrical charges. They are held
together by a positive ion - cation - and a negative ion, an anion. In other words, ionic
bonds. The ions build up into a giant ionic lattice where ions are arranged in fixed positions.
Overall the attractions are stronger than the repulsions which is why the lattice holds
together. An example of this is NaCl.
Ionic substances in a solution may dissolve. When they dissolve, the ions become
surrounded by water molecules. The positive and negative ions separate and behave
independently. Not all ionic substances dissolve though. For the ions to occur, two things
must happen: they must be separated from the lattice in an endothermic process, and then
hydrated in an exothermic process.
(b) the terms hydrated ions, enthalpy change of solution (∆solH), lattice enthalpy (∆LEH) and
enthalpy change of hydration of ions (∆hydH). P.401 - 403 and (c) the dependence of the
lattice enthalpy of an ionic compound and the enthalpy change of hydration of ions on the
charge density of the ions.
Lattice enthalpy (∆LEH): the enthalpy change when one mole of solid is formed by the
coming together of separate ions. All lattice enthalpies have negative quantities as they are
exothermic. To break the lattice, energy must be put in. This is endothermic. Lattice
enthalpy becomes more negative when ionic charges increase and ionic radii decrease.
The lattice enthalpy becomes more negative for ions with greater charge density.
Hydrated ion: An ion that has water molecules bound to them through ion-dipole bonds.
Water is a dipole, as oxygen is more electronegative than hydrogens and attracts the
electrons in the covalent bond more strongly. Water molecules may bind weakly or strongly
to ions. The hydrogen bonds that attract water molecules to each other must also be
broken to do this.
, Enthalpy change of hydration of ions (∆hydH): the strength of the attractions between ions
and water molecules. When water is acting as a solvent, aq is used. Enthalpies of hydration
also depend on concentration of solution produced, charge and small radii. Hydration of
ions is also exothermic. When dealing with solvents other than water, enthalpy of solvation
(∆solvH) is used.
Na + (g) + aq -> Na + (aq)
The greater the charge density of the ions, the greater the electrostatic attraction and the
more exothermic the lattice enthalpy, and also the greater the attraction of water molecules
and the more exothermic the hydration enthalpy.
Enthalpy change of solution (∆solH): the enthalpy change when one mole of a solute
dissolves to form a very dilute solution. It can be measured experimentally. It can also be
calculated using: (∆solH) = ∆hydH(cation) + ∆hydH(anion) - ∆LEH. The hydration of ions
favours dissolving and helps to supply the energy needed to separate the ions from the
lattice.
(i) the solution of an ionic solid in terms of enthalpy cycles and enthalpy level diagrams
involving these terms and (ii) use of these enthalpy cycles to perform calculations p.404 -
406
The solubility of a substance depends on the relationship between ∆LEH and ∆hydH. It can
first be represented as an enthalpy cycle to show the dissolving of an ionic solid.
∆hydH is always negative. To get from NaCl to the aqueous ions, ∆LEH must be subtracted.
OCR Chemistry Salters Nuffield B AS/A-Level
Key:
Definitions are in turquoise
Core practicals are in orange
Equations and tests/reactions in green
Topic 8: Oceans (O)
(a) the factors determining the relative solubility of a solute in aqueous and non-aqueous
solvents. Intermolecular bonds, ion–dipole bonds and ionic bonds should be considered.
In ionic solids, ions are held together by their opposite electrical charges. They are held
together by a positive ion - cation - and a negative ion, an anion. In other words, ionic
bonds. The ions build up into a giant ionic lattice where ions are arranged in fixed positions.
Overall the attractions are stronger than the repulsions which is why the lattice holds
together. An example of this is NaCl.
Ionic substances in a solution may dissolve. When they dissolve, the ions become
surrounded by water molecules. The positive and negative ions separate and behave
independently. Not all ionic substances dissolve though. For the ions to occur, two things
must happen: they must be separated from the lattice in an endothermic process, and then
hydrated in an exothermic process.
(b) the terms hydrated ions, enthalpy change of solution (∆solH), lattice enthalpy (∆LEH) and
enthalpy change of hydration of ions (∆hydH). P.401 - 403 and (c) the dependence of the
lattice enthalpy of an ionic compound and the enthalpy change of hydration of ions on the
charge density of the ions.
Lattice enthalpy (∆LEH): the enthalpy change when one mole of solid is formed by the
coming together of separate ions. All lattice enthalpies have negative quantities as they are
exothermic. To break the lattice, energy must be put in. This is endothermic. Lattice
enthalpy becomes more negative when ionic charges increase and ionic radii decrease.
The lattice enthalpy becomes more negative for ions with greater charge density.
Hydrated ion: An ion that has water molecules bound to them through ion-dipole bonds.
Water is a dipole, as oxygen is more electronegative than hydrogens and attracts the
electrons in the covalent bond more strongly. Water molecules may bind weakly or strongly
to ions. The hydrogen bonds that attract water molecules to each other must also be
broken to do this.
, Enthalpy change of hydration of ions (∆hydH): the strength of the attractions between ions
and water molecules. When water is acting as a solvent, aq is used. Enthalpies of hydration
also depend on concentration of solution produced, charge and small radii. Hydration of
ions is also exothermic. When dealing with solvents other than water, enthalpy of solvation
(∆solvH) is used.
Na + (g) + aq -> Na + (aq)
The greater the charge density of the ions, the greater the electrostatic attraction and the
more exothermic the lattice enthalpy, and also the greater the attraction of water molecules
and the more exothermic the hydration enthalpy.
Enthalpy change of solution (∆solH): the enthalpy change when one mole of a solute
dissolves to form a very dilute solution. It can be measured experimentally. It can also be
calculated using: (∆solH) = ∆hydH(cation) + ∆hydH(anion) - ∆LEH. The hydration of ions
favours dissolving and helps to supply the energy needed to separate the ions from the
lattice.
(i) the solution of an ionic solid in terms of enthalpy cycles and enthalpy level diagrams
involving these terms and (ii) use of these enthalpy cycles to perform calculations p.404 -
406
The solubility of a substance depends on the relationship between ∆LEH and ∆hydH. It can
first be represented as an enthalpy cycle to show the dissolving of an ionic solid.
∆hydH is always negative. To get from NaCl to the aqueous ions, ∆LEH must be subtracted.
