Disassociation of ionic compound in a solution

[NOM4DCAL]

I am trying to understand what kind of experiments have been preformed which have led us to the conclusion that ionic compounds like NaCl in something like water disassociate. From my understanding of the term disassociate this means to no longer be together or to be separate. So I wonder if this was the case how are saltwater purifiers working how is it that salt water purifiers do not sometimes pull more Na+ then Cl-. Maybe we are being taught using incorrect verbiage. Like not that they are completely dissociating but instead like we already teach ionic bonds are not particularly strong bonds and thus are easily persuaded to form other compounds which have stronger bonds. I am asking because I cannot seem to get a straight answer from any of my professors on the subject. A thought experiment I often ask them when they are adamant that they truly are dissociated is if we took 100mL’s of a solution of aqueous NaCl ( about 10% salt by volume) then take a syringe of 10mL’s of the solution and tested how much Na+ and Cl- was in the 10ml syringe there would still be 10% NaCl in the syringe so how can we say that the compound has dissociated? Is there some magic happening that causes them to be present in the same amount every time or is there still a bond between them?

[santosh]

Ionic compounds like NaCl dissociate in water, meaning they separate into Na⁺ and Cl⁻ ions. However, this doesn't mean they're completely separated with no attraction. The ions are still attracted to each other but are surrounded by water molecules, which weakens the bond enough to let them move independently. In solutions, the ions are evenly distributed, so if you take a sample, you'll still find the same ratio of Na⁺ to Cl⁻. Saltwater purifiers don't pull more Na⁺ than Cl⁻ because both ions are separated and processed equally in electrolysis.

[joan598martin]

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@NOM4DCAL wrote:

I am trying to understand what kind of experiments have been preformed which have led us to the conclusion that ionic compounds like NaCl in something like water disassociate. From my understanding of the term disassociate this means to no longer be together or to be separate. So I wonder if this was the case how are saltwater purifiers working how is it that salt water purifiers do not sometimes pull more Na+ then Cl-. Maybe we are being taught using incorrect verbiage. Like not that they are completely dissociating but instead like we already teach ionic bonds are not particularly strong bonds and thus are easily persuaded to form other compounds which have stronger bonds. I am asking because I cannot seem to get a straight answer from any of my professors on the subject. A thought experiment I often ask them when they are adamant that they truly are dissociated is if we took 100mL’s of a solution of aqueous NaCl ( about 10% salt by volume) then take a syringe of 10mL’s of the solution and tested how much Na+ and Cl- was in the 10ml syringe there would still be 10% NaCl in the syringe so how can we say that the compound has dissociated? Is there some magic happening that causes them to be present in the same amount every time or is there still a bond between them?

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Hello @NOM4DCAL,
You're raising valid points about the nature of ionic compounds like NaCl in water! Here’s a concise explanation:

When NaCl dissolves in water, it dissociates into Na⁺ and Cl⁻ ions because water molecules surround and separate the ions due to hydrogen bonding. This process is well-documented through experiments like conductivity tests, where a solution’s ability to conduct electricity proves the presence of free-moving ions (since pure water doesn't conduct electricity well).

Your concern about saltwater purifiers hinges on their design to maintain charge balance—they wouldn't selectively pull more Na⁺ than Cl⁻.

Rather than “magic,” it’s the dynamics of ionic solutions ensuring that ions remain evenly distributed in any given sample, balancing out by their very nature.

Opened up further thoughts or looking to explore something else? Let me know!


Best Regards,
Joan Martin