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Topic: Salt Water  (Read 7355 times)

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Offline microphone

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Salt Water
« on: January 29, 2007, 11:01:34 PM »
Hi,

If I had a water solution with 3 mol of sodium ions and 6 mol of chloride ions, I guess that if I evaporated the water the 3 mol of sodium would form salt with 3 mol of chloride. What happens to the other 3 mol of chloride ions?

Thanks,
mic.

Offline Bakegaku

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Re: Salt Water
« Reply #1 on: January 30, 2007, 12:02:01 AM »
You couldn't have a solution with chloride ions without any positive ion like Na+ or H+ to balance the charge.
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Offline english

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Re: Salt Water
« Reply #2 on: January 30, 2007, 12:17:12 AM »
You couldn't have a solution with chloride ions without any positive ion like Na+ or H+ to balance the charge.

Exactly.

Having 3 mol of Na+ and 6 mol of Cl- leads to Na3+Cl6-, which makes no chemical sense.

1 mol of NaCl in aqueous solution yields 1 mol of sodium and chloride ions. 
                                                           
« Last Edit: January 30, 2007, 12:58:44 AM by k.V. »

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Re: Salt Water
« Reply #3 on: January 30, 2007, 04:12:34 AM »
Unless there is other - not mentioned - cation in the solution.
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Offline maakii

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Re: Salt Water
« Reply #4 on: January 30, 2007, 11:00:03 AM »
Is it  theoretically possible to pump out Na+ ions such as the sodium pumps that work in your nerves? That would lead to a negatively charged solution with excess Cl- ions..

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Re: Salt Water
« Reply #5 on: January 30, 2007, 01:21:32 PM »
Is it  theoretically possible to pump out Na+ ions such as the sodium pumps that work in your nerves? That would lead to a negatively charged solution with excess Cl- ions..

That would be the chemical equivalent of a capacitor. Like in the flash on you camera,  charges is built up, and released to produce the energy that we see in an instant :-)

Fascinating thought. I wonder if it is possible...
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Offline microphone

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Re: Salt Water
« Reply #6 on: January 30, 2007, 10:29:44 PM »
I guess my main question is can an ion exist by itself outside of a solvent?

Also, are all solutions electronically stable? ie equal positve charge to negative charge.

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Re: Salt Water
« Reply #7 on: January 31, 2007, 08:09:42 PM »
I guess my main question is can an ion exist by itself outside of a solvent?

You create cations among others, while recording MS-spectra.
There is ions over a candleflame. Reason why I know, is that the air above a candle is leading electricity (Your 9 V battery doesn't have the potential, try 5000 V).

Charge is a matter of electrons, right? So why should it be so difficult to image that you may create a potential inside a solution, as long a there is a greater force creating it, lets say - An explosion?
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Offline microphone

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Re: Salt Water
« Reply #8 on: February 01, 2007, 10:18:18 PM »
OK, what if I go to the ocean and pull a glass of water out of the ocean. I then slowly heat the glass so the water evaporates.

Will all the cations find anions to forms salts with or will some ions be left out? If some are left over what happens to them?

Offline english

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Re: Salt Water
« Reply #9 on: February 03, 2007, 03:16:33 PM »
OK, what if I go to the ocean and pull a glass of water out of the ocean. I then slowly heat the glass so the water evaporates.

Will all the cations find anions to forms salts with or will some ions be left out? If some are left over what happens to them?

When you get that sample, all of the salt ions are paired, solvated and dissociated of course.  When you evaporate the water, mass is conserved, so the ions "find" each other again and reform crystals.

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Re: Salt Water
« Reply #10 on: February 11, 2007, 05:37:59 PM »
Thanks k.v. what topic in a chemistry book should I look up to read about ion pairing? could you please explain ion pairing. what does dissociated mean in a chemistry context? Does it mean separated, if so isn't that the same as being solvated?

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Re: Salt Water
« Reply #11 on: February 12, 2007, 01:41:12 AM »
Thanks k.v. what topic in a chemistry book should I look up to read about ion pairing? could you please explain ion pairing. what does dissociated mean in a chemistry context? Does it mean separated, if so isn't that the same as being solvated?

I would recommend ionic equilibria.  Pairing is just referring to electrostatic forces that hold ions together.

Solvated and dissociated have different meanings.  We say that NaCl is solvated (surrounded by solvent molecules) and thus becomes dissociated (Na+ and Cl- ions separate).

Most species in solution become solvated and do not dissociate to an aprreciable extent, i.e. precipitates, weak acids.

In the context of acid chemistry, dissociate means to separate a proton (H+).  We are ionizing the acid.

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Re: Salt Water
« Reply #12 on: February 12, 2007, 03:14:38 AM »
thanks k.v.

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Re: Salt Water
« Reply #13 on: February 13, 2007, 10:31:23 PM »
To get more specific on the reason why Na and Cl ions decide to pair up and there are no free ions in an evaporated solution of a particular sample, say ocean water, you must bring the effects of water into account.

Technically, oceanic water contains other ions, i.e. Mg2+, Ca2+, but we'll ignore these for simplicity because some pretty strange things can happen with those ions and what you know of as the salt crystal.


In any particular sample of salt water, Na and Cl ions are insulated from each other by water molecules, such that their charges are stabilized—their charges are spread out over a larger area.

However, even with the added solvated effect of water, the Na and Cl ions still have a really weak ion-ion interaction.  So when the solution is evaporated of all the water molecules, you're left with those Na and Cl ions completely paired, with no free ions.  This weak interaction between the ions in solution (in water) lock them into place in such a way as to direct them towards each other again so they can reform the stronger, former ionic bond as water evaporates.

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