Transcript Thermodynamics and Further Inorganic Chemistry

Thermodynamics and Further
Inorganic Chemistry
Contents
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Thermodynamics
Periodicity
Redox Equilibria
Transition Metals
Reactions of Inorganic Compounds in
Aqueous Solution
Thermodynamics
• Enthalpy Change
• Free Energy and Entropy Change
Enthalpy Change (ΔH)
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Enthalpy of formation
– is the enthalpy change that occurs when one mole of compound in its standard
state if formed from its element in their standard states under standard
conditions.
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Ionisation enthalpy
– of an element is the enthalpy change that accompanies the removal of 1 electron
in each atom of one mole of gaseous atoms
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Enthalpy of atomisation
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Bond dissociation enthalpy
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of an element is the enthalpy change that occurs when 1 mole of gaseous atoms
are formed from the element in its standard state.
the enthalpy required to break and separate one mole of bond so that the
gaseous atoms exert no force on each other.
Electron affinity
– of an element is the enthalpy change that accompanies the addition of one
electron to each atom in one mole of gaseous atoms.
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Lattice enthalpy
– of an ionic compound is the enthalpy change which accompanies the formation
of one mole of ionic compound from its constituent gaseous ions.
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You will need to be able to use the quantities in energy cycle calculations,
by applying Hess’s Law. Remember it doesn’t matter which route a
chemical change takes the net energy change stays the same.
Free-Energy Change (ΔG) and
Entropy Change (ΔS)
• ΔH, whilst important, is not sufficient to explain spontaneous
endothermic
• The concept of increasing disorder (entropy change, ΔS) accounts
for this phenomenon.
• E.g. the production of CO2 from hydrogencarbonates with acid is
endothermic but occur without heating, because the products are
significantly more disordered than the products.
• The balance between entropy and enthalpy determines the
feasibility of a reaction and is called the change in free energy.
• ΔG = ΔH - TΔS (derivation not required).
• ΔG must be negative for a reaction to proceed spontaneously
Periodicity
• All trends listed here are for progression from left to right
of the periodic table.
• Reactions with water Na and Mg
– Reactivity decreases
• The elements progress from metals to non-metals.
• Period 3 oxides
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The amount of oxygen per mole of the element increases.
The oxides go from ionic to covalent in nature
The oxides go from basic to acidic
Boiling points start high because of the strong ionic forces but
become higher in the middle oxides have mixed ionic covalent
nature but still form giant lattices so have high boiling points. The
latter few elements form simple molecules and so have relatively
low boiling points.
– With water the metal oxides form alkalis, aluminium and silicon
oxide are insoluble, the non metal oxides form acidic solutions
Periodicity
• Period 3 chlorides
– The amount of chlorine atoms per mole of the
elements increases across the period.
– Sodium and magnesium chloride are ionic where as
aluminium chloride is covalent as are the chlorides of
the other elements.
– The boiling points decrease as the compounds
progress from giant ionic in structure to simple
molecular.
– Ionic chlorides form neutral solution form neutral
solutions in water. Covalent chlorides are hydrolysed
by water and form acidic solutions.
Redox Equilibria
• Oxidation and reduction are electron transfer reactions,
– Oxidation is the loss of electrons and reduction is the gain of
electrons
• In redox reactions the oxidation state of the elements
involved changes.
• When looking at redox reactions it is sometimes useful to
look at half equations involved.
• Na  Na+ + e- the sodium has lost an electron and has
therefore been oxidised
• 2H2O + 2e-  H2 + 2OH- the hydrogen has lost an
electron and has therefore been reduced.
• These can be combined as follows:
• 2Na + 2H2O  NaOH + H2
• Half equations are also useful when examining the
reactions that occur during electrolysis
Transition Metals
• Transition metal characteristics arise from
an incomplete d sub-level.
• These characteristics include complex
formation, high density, high melting
points, formation of coloured ions, variable
oxidation state and catalytic activity.
Variable Oxidation States
• Transition elements are able to form more
than one ion, each with a different
oxidation state, by losing the 4s electrons
and different numbers of 3d electrons.
• When forming ions, the 4s electrons are
lost first, before the 3d electrons.
Transition Metal Complexes
• A ligand is a molecule or ion that bonds to a
metal ion
– By forming a coordinate bond (one in which both
bonding electrons come from the same element)
• A complex ion is a central metal ion surrounded
by ligands.
• The coordination is the total number of
coordinate bonds from ligands to the central
transition metal ion.
Shapes of Molecules
• Six coordinate complexes form octahedral
shapes.
• Four coordinate complexes form
tetrahedral shapes
The effects of partially filled dorbitals
• In the partially filled d-orbitals of a transition
metal complex there is an energy gap between
the filled and unfilled orbitals. Visible light is of
the right energy to promote an electron from one
energy level to the next.
• The partially filled orbitals in transition metals
allow them to change oxidation states readily
and thus facilitate the progress of some
reactions
Reactions of Inorganic Compounds
in Aqueous Solution
• Lewis Acids and Lewis Bases
• Metal Aqua ions
• Substitution reactions
Lewis Acids and Lewis Bases
• Lewis acids are lone pair acceptors such
as the central metal ion in transition metal
ion.
• Lewis bases are lone pair donors such as
ligands.
• These definitions of acids and bases
extend acid -base chemistry to reactions
which do not involve proton exchange.
Metal Aqua Ions
• Metal aqua ions are formed in aqueous
solution:
• [M(H2O)6]2+ and [M(H2O)6]3+
• These aqua ions can be present in the
solid state for example copper sulphate is
only blue as its aqua ion, in its anhydrous
state it is white.
Substitution Reactions
• The ligands NH3 and H2O are similar in size and are uncharged,
therefore ligand exchange between these two ligands occurs without
change of co-ordination number.
• The Cl− ligand is larger than these uncharged ligands and that
ligand exchange can involve a change of co-ordination number.
• Substitution of unidentate ligands (ligands which forms one bond
with the central metal) with a bidentate or a multidentate ligand
leads to a more stable complex.
• This is because there will be more ions in solution therefore higher
entropy.
• [M(U)6]2+ + 3B [M(U)3]2+ + 6B
4 ions in solution 7 ions in solution
less entropy
more entropy
Summary
• Thermodynamics
– Standard enthalpy changes can be used to calculate energy changes in
unknown reactions
• Periodicity
– The trends in period three reactions can be explained in terms of the
structure and bonding of their compounds
• Redox Equilibria
– When looking at redox reactions it is sometimes useful to write half
equations.
• Transition Metals
– The chemistry and properties of transition metals can be explained by
their partially filled d-orbitals.
• Reactions of Inorganic Compounds in Aqueous Solution
– The idea of aid base chemistry can be extended by the concept of
Lewis acids and bases.