maxpayne
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What about; two single molecule reaction?
As we all know this: H2SO4 + 2NaOH -> Na2SO4 + 2H2O
... my question is: What if single molecule of H2SO4 meets with single molecule NaOH, will they react? If not, why, when we all
say that acids reacts with bases.
Also does probability allows such situation?
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Vargouille
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They will react to form sodium bisulfate.
H2SO4 -> H+ + HSO4-
NaOH -> OH- + Na+
H+ + OH- <=> H2O
Na+ + HSO4- <=> NaHSO4
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blogfast25
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maxpayne:
In watery medium there really are no ‘molecules’ like that.
NaOH is a stong base and dissociates completely in water:
NaOH(s) === > Na+(aq) + OH-(aq)
H2SO4 is a strong acid and also more or less completely dissociates in water:
H2SO4(l) + H2O(l) === > H3O+(aq) + HSO4 -(aq)
When both solution are mixed together the only chemical reaction taking place is:
H3O+(aq) + OH-(aq) === > 2H2O (l)
The sodium and sulphate ions play no part whatsoever. They’re so-called ‘spectator ions’.
[Edited on 5-8-2012 by blogfast25]
[Edited on 6-8-2012 by blogfast25]
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maxpayne
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blogfast25, thank you very much, if I only had a teacher like you back then 
Let's keep up on this: NaOH(s) === > Na+(aq) + OH-(aq)
This dissociation is what I suspected it was happening, but please, tell me how this ions look like in solution, are they really free or bonded
somehow and from where the energy came, that is needed for dissociation??
This energy question is my big ???.
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blogfast25
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They’re solvated to some degree: surrounded by water molecules that are loosely ‘bound’ to them because H2O is a permanent dipole.
When dissolving an ionic solvent, the energy needed is basically governed by Gibbs Law: ΔG = ΔH – TΔS. ΔH is the Enthalpy of
Solution and is basically the energy to break up the ionic lattice, it is always positive (ENDOthermic). ΔS is the Entropy change, it is always
positive (solutions are more disordered than ionic lattice) and thus - TΔS is always negative. Whether an ionic compound is soluble in water
basically depends on whether ΔG is negative or not (negative = soluble).
But there’s a complication, at least in many cases. When solvated ions are formed, usually solvation energy is released and as a result the
dissolution of many ionic compounds is EXOthermic. Interesting examples are both NaOH and H2SO4, the dissolution process of which in both cases
releases considerable amounts of Enthalpy: the solutions heat up as you make them.
In reality we cannot distinguish between the energy needed to break up the lattice and energy released due to solvation. These overall energies are
determined experimentally.
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maxpayne
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Very, very interesting.
It is a large subject to study this Gibbs free energy, and I have too many questions now to ask here about energy origins, etc. So let's get back to
that spectator ions.
You actually said before that sodium and sulfate ions "watch" the reaction in solution of other "relatively free" ions. So, it seems to me that sodium
sulfate begins to exist only when there is no water and it is wrong to say that when I mix before mentioned solutions, I get Na2SO4. Only when I
remove the water (evaporating) I produced Na2SO4?
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blogfast25
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Yes, that's quite correct. Although we tend to see a sodium sulphate solution as a, erm... sodium sulphate solution, more modern interpretations tend
to see the solute (sodium sulphate)/solvent (water) system as a system in its own right with its own properties.
Certainly a sodium sulphate solution doesn't share many properties with the 'pure substance', Na2SO4(s),of which in itself different (solid) hydrates
exist (each with their own properties, such as MP and BP). These hydrates can co-exist in equilibrium with sodium sulphate solutions of various
concentrations (see sodium sulphate - water phase diagrams).
[Edited on 6-8-2012 by blogfast25]
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maxpayne
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blogfast25, thank you for your time and effort, I see you have quite knowledge in this "invisible chemistry", and you certainly did answered as a
professional. I did learned something useful now, and it seems that things are not always as they look at the first glance.
[Edited on 6-8-2012 by maxpayne]
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blogfast25
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You're welcome. There are plenty good texts available online about this subject. Search and ye shall find.
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bbartlog
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Although the case of H2SO4 + NaOH is pretty straightforward, I think there are some other ones that are less so. For example
Al + Cl2 --> ??
Li + O2 --> ??
The less you bet, the more you lose when you win.
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blogfast25
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Quote: Originally posted by bbartlog  | Although the case of H2SO4 + NaOH is pretty straightforward, I think there are some other ones that are less so. For example
Al + Cl2 --> ??
Li + O2 --> ??
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Sure but direct union of elements are a different class altogether!
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barley81
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Blogfast, enthalpy of solution and entropy of solution do not have to be positive! Calcium chloride dissolves exothermically, while ammonium nitrate
dissolves endothermically (enthalpy of solution for calcium chloride is negative; enthalpy of solution for ammonium nitrate is positive).
Entropy of solution need not be positive. Hydrocarbons have negative entropy of solution in water:
http://www.mpcfaculty.net/mark_bishop/solubility_entropy.htm
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maxpayne
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I think that blogfast already mentioned complication, and without further thinking, my sense tells me that mathematical equation do not explain
everything, especially in Gibbs law. For me, it is just a fancy equilibrium math, nothing else.
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blogfast25
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| Quote: Originally posted by barley81 | Blogfast, enthalpy of solution and entropy of solution do not have to be positive! Calcium chloride dissolves exothermically, while ammonium nitrate
dissolves endothermically (enthalpy of solution for calcium chloride is negative; enthalpy of solution for ammonium nitrate is positive).
Entropy of solution need not be positive. Hydrocarbons have negative entropy of solution in water:
http://www.mpcfaculty.net/mark_bishop/solubility_entropy.htm |
I never said anything of the sort: you badly read what I wrote.
ΔG has to be negative for dissolution to occur, not necessarily ΔH.
When dissolution is endothermic that’s because TΔS > ΔH. Even endothermic dissolution means ΔG < 0. And ΔS is always
positive.
Where dissolution is exothermic, that's caused by solvation energy (always negative).
Re. your link, nice page but sorry, but I don’t believe that:
<i>”Overall, the attractions in the system after hexane and other hydrocarbon molecules move into the water are approximately equivalent in
strength to the attractions in the separate substances. For this reason, little energy is absorbed or evolved when a small amount of a hydrocarbon is
dissolved in water. To explain why only very small amounts of hydrocarbons such ashexane dissolve in water, therefore, we must look at the change in
the entropy of the system. It is not obvious, but when hexane molecules move into the water layer, the particles in the new arrangement created
areactually less dispersed (lower entropy) than the separate liquids. The natural tendency toward greater dispersal favors the separate hexane and
water and keeps them from mixing.”</i>
This argument is pulled by the hair. Far more likely is that the hydrogen bonds in water cannot be easily broken by the hydrocarbon. But a weak
solution of hexane in water still has higher entropy than the separate fluids, undissolved in each other.
[Edited on 7-8-2012 by blogfast25]
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blogfast25
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No Sir. This thermodynamic principle underpins much of chemistry, including theory on what and what doesn't dissolve in each other.
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maxpayne
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| Code: | No Sir. This thermodynamic principle underpins much of chemistry, including theory on what and what doesn't dissolve in each other. |
Then explain to me origins of the energy needed for dissociation. Equation uses only quantities, it does not explain themselves.
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blogfast25
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This energy is known as lattice energy or coulombic energy. Imagine the ions of NaCl (Na+ and Cl-) to be at infinite distance from each other (at
atomic scale even infinite isn’t very large). These ions are attracted by electrostatic attraction (coulombic attraction) and thus have Potential
Energy. When allowed to move towards each other this potential energy is converted into heat (enthalpy).
But when dissolving the lattice, the exact opposite must happen and we must expend energy (enthalpy) to rip the lattice apart. Without solvation
energy the dissociation energy is therefore closely related to the lattice energy. Exact values for molar lattice energies are known and tabled, for
many binary ionic compounds.
The lattice energy usually makes up the most part of the Enthalpy of Formation of (solid!) binary ionic compounds, see Born Haber Cycle:
http://en.wikipedia.org/wiki/Born%E2%80%93Haber_cycle
[Edited on 7-8-2012 by blogfast25]
[Edited on 7-8-2012 by blogfast25]
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