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Topic: Is NaCl a base?  (Read 79338 times)

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Offline gonzo

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Is NaCl a base?
« on: October 17, 2007, 07:22:41 AM »
I've become a bit confused and was hoping someone could clear this up for me.. I didn't think NaCl was a base, in fact I thought Na+ and Cl- exactly neutralized each other. Even though Na is an alkali metal, is NaCl an alkali salt?

This has me confused too: http://chemed.chem.purdue.edu/demos/demosheets/11.5.html - it says that NaCl is a weaker conjugate base..

Could someone clear this up for me please?  ???

Offline AWK

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Re: Is NaCl a base?
« Reply #1 on: October 17, 2007, 07:48:16 AM »
Statement in your link is a misconception. It mixes Bronsted theory in which Cl- anion is a base with Arrhenius theory in which a molecule may show basic pH becouse it is a base or after hydrolysis, but even that this is not the case for NaCl.

I think it will be a stormy discussion after my statement. I personally think chemed.chem.purdue is after all very valuable link.
« Last Edit: October 17, 2007, 08:12:08 AM by AWK »
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Offline gonzo

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Re: Is NaCl a base?
« Reply #2 on: October 17, 2007, 08:41:33 AM »
I'm still confused.. if there's likely a discussion regarding your statement it's because it's not clear whether NaCl is a base or not? I thought this would be simple enough, but.. isn't a base defined as a substance that will accept H+? How will NaCl accept a proton?

Offline Borek

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Re: Is NaCl a base?
« Reply #3 on: October 17, 2007, 09:53:55 AM »
Cl- can act as a very weak base and accept proton:

Cl- + H+ = HCl

Na+ can act as a weak acid:

Na+ + H2O = NaOH + H+

and while it is still very weak, it is at least 1000 times stronger acid than Cl- is a base.

Purdue page can be confusing, IMHO mostly because it treats all substances listed as either acids or bases, not mentioning their dissociation to individual ions with their own characetristics.
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Offline gonzo

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Re: Is NaCl a base?
« Reply #4 on: October 17, 2007, 10:34:36 AM »
Thank you..

So the answer is that overall NaCl can act as a (weak) base, but it's not a base on its own, is it? I'm not really sure if that even makes sense, I understand your examples, so I'm not sure how to phrase it..

If we dissolve NaCl in H2O we have a saline solution that would measure as pH neutral, wouldn't it? Or does it in fact measure as a very weak base?

Offline Borek

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Re: Is NaCl a base?
« Reply #5 on: October 17, 2007, 10:52:00 AM »
Depending on how deeply you want to dig it will be either neutral (answer than 99.99% of teachers will accept) or very slightly acidic (fact that remaning 0.01% of teachers is aware of ;) ).

In NaCl solution you have very weak Bronsted-Lowry base (Cl-) and very weak Bronsted-Lowry acid (Na+). They are so weak, you may treat them as simply neutral. If you want to be a nitpicker - Na+ is strong enough Bronsted-Lowry acid to give measurable change in pH (0.1M NaCl solution should have pH of 6.98). Cl- - while technically a Bronsted-Lowry base - is too weak a base to give observable effects.

But for all practical purpose you may safely ignore both Na+ being an acid and Cl- being a base.
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Offline AWK

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Re: Is NaCl a base?
« Reply #6 on: October 17, 2007, 11:37:36 AM »
As Borek pointed out, pH of NaCl solutions will be slightly below 7, but this is caused by "ionic strenth effect" (influence of ionic strength on activity of H+ ion), and not by different basicity of ions in NaCl.
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Offline gonzo

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Re: Is NaCl a base?
« Reply #7 on: October 17, 2007, 12:00:48 PM »
I really appreciate the answers!

I mistyped before with base/acid - so I'm still confused I guess ;)

If NaCl can act as a weak conjugate base (according to purdue.edu), then what is up with that? Shouldn't it have been a weak conjugate acid?

Offline AWK

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Re: Is NaCl a base?
« Reply #8 on: October 17, 2007, 12:40:40 PM »
As I and Borek pointed out, term conjugate is not applicable to NaCl.
According to Bronsted-Lowry theory acid donates proton, base accepts proton. Na+ from NaCl is not a Bronsted acid or base. Cl- is of course a very very weak Bronsted base.
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Offline gonzo

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Re: Is NaCl a base?
« Reply #9 on: October 17, 2007, 01:18:05 PM »
Thank you!

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