Posted by Kyo on October 30, 2007 at 17:34:20:
In Reply to: [Chemistry] - Practice Qns for 'A' level Examinations posted by Kyo on October 30, 2007 at 17:30:29:
Ans to Q1
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Ans to Q2
C15 H30 O3
(Note : the tricky bit about this qn refers to the fact that you cannot determine directly from the products how many moles/mass of O2 was present in the organic compound, because even after adding up the oxygen in both products CO2 and H2O, you don't know how much of the oxygen came from the compound, how much came from the air. You have to find the moles/mass of oxygen in the compound indirectly, by taking (total mass of compound 0.43g) - (mass of C in products) - (mass of hydrogen in products).
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Ans to Q3
During storm, atmospheric pressure at ground level decreases sharply.
Gas pressure within coal mine > gas pressure outside coal mine.
Gases inside coal mine rush out via entrance to outside.
Coal mines contain varying levels of poisonous gases, eg. methane, asphyxiant gases (specifically carbon monoxide).
The observation that the man died at the entrance, supports hypothesis of poisonous gases in mine.
(In H2/H1 context, other relevant info for this qn refers to ligand displacement of CO over O2)
Carbon monoxide ligand has approx 200 times greater affinity for haemoglobin compared to oxygen.
Fe-CO dative/coordinate bond is stronger than dative/coordinate Fe-O2 bond.
CO displaces the oxygen in oxyhaemoglobin to form carboxyhaemoglobin.
HbO2 + CO -> HbCO + O2
Consequently, the brain and body tissues/organs suffer from oxygen deprivation, resulting in brain/tissue/organ damage, loss of consciousness, and eventually death.
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Ans to Q4 :
By Hess Law, Solution enthalpy = (endothermic) Lattice Dissociation enthalpy + Solvation (hydration) enthalpy.
(Note : lattice enthalpy is usually defined exothermically; when defined endothermically, ie. the energy required to break apart one mole of ionic solid to form gaseous ions; this term is properly called lattice dissociation enthalpy).
Solvation enthalpy (a subset of which is hydration enthalpy, when the solvent is water), refers to the exothermic (ie. heat releasing) process of forming ion-dipole bonds/forces-of-attraction between the ions and the polar water molecules (hence the term ion-dipole; ion-dipole bonds are weaker than ionic bonds, but stronger than hydrogen bonds and dipole-dipole van der Waals forces. This is common sense because ion-dipole bonds are half ionic (ie. full formal charge) and half dipole (ie. partial/delta charge).
Hence, by Hess Law, solution process can be regarded as first investing energy in form of heat (enthalpy is heat transfer under constant pressure; energy is heat transfer under constant volume), to break apart solid ionic compound into gaseous ions (ie. lattice dissociation process), then getting your energy returns when ion-dipole bonds form (ie. solvation process).
Hence, whether solution occurs depends on whether investment outweights returns, or whether returns outweighs investment; ie. whether it is thermodynamicaly feasible. But so far we're only considering about enthalpy, which will suffice for explaining solubilities of hydroxides and sulphates (but not carbonates, which requires another variable in Gibbs Free Energy equation, more on this later).
As we go down Grp II, because charge density decreases (since ionic radius increases), BOTH lattice dissociation enthalpy and solvation (hydration) enthalpy decreases, for BOTH hydroxides and sulphates/carbonates.
However, the rate at which the lattice dissociation enthalpy falls, relative to the rate at which solvation (hydration) enthalpy falls, is different for hydroxides vs sulphates/carbonates. (Note that nitrates are always completely soluble).
For small hydroxide anions, as we go down Grp II, the rate of fall of lattice dissociation enthalpy is faster than the rate of fall of solvation (hydration) enthalpy.
For large sulphate/carbonate anions, as we go down Grp II, the rate of fall of lattice dissociation enthalpy is slower than the rate of fall of solvation (hydration) enthalpy.
The reason for this difference in fall relates to the fact that lattice dissociation enthalpy is affected by the inter-ionic distance. Because OH- ions are so small, as we go down Grp II cations, the inter-ionic distance is more significantly changed. Because SO4 2- and CO3 2- ions are so large, as we go down Grp II cations, the inter-ionic distance is less significantly changed.
In effect, the result is the SOLUTION (ie. lattice dissociation + solvation) enthalpies of Grp II hydroxides enthalpies become MORE exothermic as we go down the group; while the SOLUTION enthalpies of Grp II sulphates/carbonates become LESS exothermic as we go down the group.
Thus, solubilities of Grp II hydroxides increase down the group, while solubilities of Grp II sulphates/carbonates decrease down the group.
(Bonus section : why Grp II carbonates solubilities first decrease from Be to Sr, then increase at Ba?)
Gibbs Free Energy (ie. feasibility and spontaneity of reactions, eg. dissolving a compound in water) does not only consider enthalpy, but also another important variable - entropy.
Each time you dissolve a compound in water, entropy simultaneously increases and decreases (of course, there will be a net increase OR decrease in entropy; this depends on the ionic compound and the solvent).
The entropy increase is caused by a greater degree of disorder when you mess up the orderly lattice structure of an ionic compound during the solution process. The decrease in entropy is caused by a lesser degree of disorder of the solvent (eg. water) molecules when they arrange themselves around the cations and anions, during the formation of ion-dipole bonds.
At the top of any Grp (eg. I or II), the high charge density cation has a greater ordering effect (ie. entropy decreasing effect) on the solvent/water molecules. Thus, near the bottom of the group, eg. Barium, the net entropy effect is an INCREASE in entropy.
All the while, whether considering hydroxides or sulphates or carbonates, and whether considering top or bottom of any Group, BOTH enthalpy and entropy variables need to be considered in Gibbs Free Energy equation.
However, for hydroxides and sulphates, the enthalpy argument will suffice, ie. enthalpy effects outweights entropy (in case of sulphates) and/or is non-opposing with entropy (in case of hydroxides).
But for solubilities of Grp II carbonates, enthalpy effects outweights entropy effects from Be to Sr, then entropy effects outweights enthalpy effects from Ba onwards.