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Chemistry Forums for Students => High School Chemistry Forum => Topic started by: mithrilhack on July 03, 2005, 11:00:49 PM

Title: pH of Ethanol
Post by: mithrilhack on July 03, 2005, 11:00:49 PM
Does ethanol have a pH? Wikipedia states it has a pKa of 15.9...what does that mean?
Title: Re:pH of Ethanol
Post by: Mitch on July 04, 2005, 01:17:03 AM
Well, pH can change by either adding more H+ in solution or taking some out.
Title: Re:pH of Ethanol
Post by: DrCMS on July 04, 2005, 03:52:51 AM
Ethanol does not have a pH.  

As pH by its very definition is the log of the hydrogen ion concentration in aqueous solution.  If you don't have water you dont have a pH.  You can stick a pH probe in pure ethanol and get a value on the meter but that's not its pH.
Title: Re:pH of Ethanol
Post by: xiankai on July 04, 2005, 04:53:33 AM
checking the website where u derived the pKa from, a notation comes with it to "(H+ from OH group)"

it seems that ethanol has H+... in a wierd way?
Title: Re:pH of Ethanol
Post by: mithrilhack on July 04, 2005, 01:23:35 PM
Well actually I want to know if ethanol mixed with water would change the pH of the water.
Title: Re:pH of Ethanol
Post by: lemonoman on July 04, 2005, 01:36:26 PM
Not sure if this will help or not but...

When ANYTHING is put into water, there will be an 'equilibrium'  involved.  Water is called the 'universal solvent' for this reason (even though some things seem NOT to dissolve, they do, to a very small degree)

Anyways, when ethanol (C2H5OH) is put into water, a few different equilibrium could occur:

C2H5OH + H2O <==> C2H5OH2+ + OH-

which 'produces' a hydroxide (the easiest way to visualize basicity)

Or, possibly:

C2H5OH + H2O <==> C2H5O- + H3O+

where H3O+ is H2O with the classic H+, that you mentioned earlier, attached to it.

Technically speaking, both of these equilibria occur to some degree.  One will occur more than the other - and so one of OH- or H+ will be more plentiful in the solution (as opposed to a perfectly neutral solution, where the exact same amount of each is present).  If there is more OH-, then the pH will be basic (and the pH value will tell you exactly how much OH- there is) - on the other hand, if the pH is acidic then there is more H+ and you can figure the math of that out too, if you like.
Title: Re:pH of Ethanol
Post by: Mitch on July 04, 2005, 02:01:06 PM


C2H5OH + H2O <==> C2H5O- + H3O+


That's where knowing the pKa comes in. Ka is equal to the above quoted equilibrium. So you tell us the extent of disocciation of ethanol in aqueous solution. From there we'll move on determining the pH with different mol amounts of ethanol.
Title: Re:pH of Ethanol
Post by: AWK on July 05, 2005, 02:44:50 AM
pH of diluted ethanol solutions in water is practically the same as that of water. It can be calculated from two equilibriums - Kw (ion product for water) and Ka.
Title: Re:pH of Ethanol
Post by: Borek on July 05, 2005, 03:47:44 AM
http://www.chembuddy.com/?left=pH-calculation&right=pH-weak-acid-base (http://www.chembuddy.com/?left=pH-calculation&right=pH-weak-acid-base)
Title: Re:pH of Ethanol
Post by: TakeItEasy on July 20, 2005, 07:55:33 AM
Ethanol does not have a pH.  

As pH by its very definition is the log of the hydrogen ion concentration in aqueous solution.  If you don't have water you dont have a pH.  You can stick a pH probe in pure ethanol and get a value on the meter but that's not its pH.
:oif that's not the pH of ethanol, then what is it?
Title: Re:pH of Ethanol
Post by: DrCMS on July 20, 2005, 11:18:32 AM
Thats the point the meter will give a reading but by definition in a non-aqueous system the number can NOT be pH.  To be a true pH reading you must have water.
Title: Re:pH of Ethanol
Post by: Dude on July 20, 2005, 04:23:28 PM
Come to think of it, I don't know the answer to Quark's question.  pH is operationally defined as a cell potential generated by a certain amount of H+ in an aqueous buffer solution upon a glass electrode.  If one places the electrodes in ethanol or hexane, a value comes out.  One can logically argue based upon activity that the value isn't referable to the aqueous buffer solutions that were used to calibrate the instrument.  However, what creates the potential to cause a value to register?  Now I'm confused as to why I really can't divide by 0 other than 20 years of educators telling me that it is undefined.
Title: Re:pH of Ethanol
Post by: Borek on July 21, 2005, 07:25:02 PM
why I really can't divide by 0 other than 20 years of educators telling me that it is undefined.

For division:

if x/y=a then ay=x

but it is not satisfied for y=0.

From what I know pH scale can be generalized for use in other solvents than water, thus it seems for me ethanol has a pH - and the neutral one of 9.45, as it autodissociates just like water does (but the ethanol ionic product is - at 10-18.9 - substantially smaller then Kw).
Title: Re: pH of Ethanol
Post by: X on July 05, 2007, 07:06:38 PM
If anyone still cares, when measuring pH of ethanol, pHe is measured. It is an apparent pH with special equipment and circumstances, since to have a pH, an aqueous solution is necessary.
Title: Re: pH of Ethanol
Post by: Medic851 on November 22, 2007, 06:16:32 PM
Of course ethanol has a pH

Everything organic has a pH and you could probably argue that inorganic compounds have a pH depending on the definition of an acid (bronsted lowry, lewis etc..)

If Wikipedia is correct in stating that the pKa is 15.9 than that value is related to pH by the negative logarithm of pKa.

So.... pH = (-log)pKa---------- if you use the inverse log function on your calculator which is 10^x power times the pKa 15.9 you should derive a number around 7.9 which represents the pH. So Ethanol is slightly basic which is consistent with the idea that that OH (hydroxide ion) the characteristic feature of alcohols is present in solution.
Title: Re: pH of Ethanol
Post by: Medic851 on November 22, 2007, 06:29:07 PM
Well that sounds good but its not entirely correct.


pKa = -log Ka

overall equation pH = pKa + log(base/acid)

The relationship of the Henderson-Hasselbach equation helps us to understand the pH of solutions based on relative concentrations of the acid and base. Using this equation one can create buffers, calculate pH or figure out concentrations of acids or bases in solution. Given three values-- the fourth can be solved for algebraically.
Title: Re: pH of Ethanol
Post by: Borek on November 22, 2007, 06:49:01 PM
So.... pH = (-log)pKa---------- if you use the inverse log function on your calculator which is 10^x power times the pKa 15.9 you should derive a number around 7.9 which represents the pH.

You are comparing apples and oranges. pH is a measure of H+ activity in water solution.

Quote
So Ethanol is slightly basic which is consistent with the idea that that OH (hydroxide ion) the characteristic feature of alcohols is present in solution.

Completely wrong. Just because ethanol molecule ends with -OH  doesn't mean it dissociates giving OH-. If anything, it is slightly acidic, able to produce ethanolates - salts in which proton from the -OH is replaced by metal.

Also note, that pH of pure water changes with temperature:

07.47
257.00
506.63
756.35
1006.14

Does it mean that cold water is basic and hot acidic? In all cases [H+]=[OH-]
Title: Re: pH of Ethanol
Post by: Medic851 on November 23, 2007, 05:40:46 PM
So then what is the pH of lets say aq. Ethanol CH3CH2OH at STP? I think that was the orignial question.

Okay, I agree that perhaps alcohol is more acidic than basic but isnt it true that alcohols can act as a base in chemical reactions. They can become pronated yeilding alkyloxonium ion. This lowers the pKa making the compound more acidic. The Hydrogen ion can also leave giving CH3CH2O- which acts as a base.

you are saying that temperature is a determinate factor of a solutions pH but in all instance there is an established equilibrium. That sounds correct but how does this relate to the pH of ethanol at STP?

Furthermore, apples and oranges are both round fruit. Perhaps color and composition are differnt but the infuence of strucutre in chemistry surely can't be ignored.
Title: Re: pH of Ethanol
Post by: Medic851 on November 23, 2007, 05:55:46 PM
I am pretty sure that pKa is directly related to how much a compounds will dissociate H+ ion in solution and therefore is related to pH. (while pKb is related to how much OH- is in solution) and they are both related to Kw.

I just checked a table in my organic chemistry text. the higher the pka the LESS acidic the compound and conversely the lower the pKa the MORE acidic the compounds.

Ethanol is listed at 16
water is 15.7 so Ethanol is in fact slightly less acid (or more basic) than water.

and water is amphoteric --meaning it can act as either acid or base.

Acetic acid has a pKa of 4.7--- pretty acidic for an organic substance

Inorganic compounds such as Hydrogen Iodide HI (pka -10.4) are even more acidic.

Sulfuric acid is pKa -4.8

These values are for compounds with water as the solvent
so water then becomes pronated by acids yielding H30+ (hydronium ion) which is given a pKa of 0 by default.

so with all do respect I'm not "completely wrong"

Title: Re: pH of Ethanol
Post by: AWK on November 26, 2007, 03:06:45 AM
I am pretty sure that pKa is directly related to how much a compounds will dissociate H+ ion in solution and therefore is related to pH. (while pKb is related to how much OH- is in solution) and they are both related to Kw.

I just checked a table in my organic chemistry text. the higher the pka the LESS acidic the compound and conversely the lower the pKa the MORE acidic the compounds.

Ethanol is listed at 16
water is 15.7 so Ethanol is in fact slightly less acid (or more basic) than water.

and water is amphoteric --meaning it can act as either acid or base.

Acetic acid has a pKa of 4.7--- pretty acidic for an organic substance

Inorganic compounds such as Hydrogen Iodide HI (pka -10.4) are even more acidic.

Sulfuric acid is pKa -4.8

These values are for compounds with water as the solvent
so water then becomes pronated by acids yielding H30+ (hydronium ion) which is given a pKa of 0 by default.

so with all do respect I'm not "completely wrong"


Of course, you are not wrong theoretically, but practically change of pH after ethanol addition cannot be measured, and calculations may give difference, i guess, eg 0.0001 unit of pH
Title: Re: pH of Ethanol
Post by: fireemblem555 on January 04, 2009, 12:16:56 AM

Wikipedia says the pKa of Ethanol is 15.9

pKa=-logKa
so Ka= 10^(-15.9) or inverse log (-15.9)

Ka=1.26 x 10 ^-16

Ka = [A-][H+]             Ka= [CH3CH2O-][H+]
        ----------                    ------------------
            HA                              [CH3CH2OH]
Since ethanol is monoprotic (as far as dissociation is concerned) [A-] = [H+]




1.26 x 10^-16 = [H+][H+]
                         --------------
                              [HA]
1.26*10^-16 * [HA] = [H+] squared

For a solution having a 1M CH3CH2OH concentration

[H+] squared = 1.26 x 10^-16 mol/L
[H+] = 1.12 x 10^-8 mol/L
pH = -log [H+]
pH = -log [1.12 x 10^-8 mol/L0
pH = 7.95 which is mildly basic

Using the equation y= -log(sqrt((1.26*10^-16)x) where x is the concentration of the chemical being dissociated, the intersection point with 7 will show when the pH will cease being basic, which is at 79 mol/L a concentration that is so high as to be irrelevant.

Also as an end note, I know that ethanol can act as an acid in a sufficiently basic environment, I took Organic Chemistry in University.  I just wanted to make the definition relatively simple.

I apologize if I have made any errors, please correct me if I have.
Title: Re: pH of Ethanol
Post by: Loyal on January 04, 2009, 01:15:53 AM
Alcohols more readily acts as a lewis base.   Alcohols can easily form EtOH2+, but usually when this happens a reaction occurs and the product can be different depending on the conditions and acid used. 

In the case of EtO- that is a super base and usually created from sodium metal reacting with Ethanol

2Na + 2EtOH --> 2Na+ + 2EtO- + H2

This base in exceedingly powerful and it is near impossible for it to exist in water. So I would imagine the equilibrium constant for ethanol to the ethoxide ion would be rather low.
Title: Re: pH of Ethanol
Post by: Borek on January 04, 2009, 05:31:01 AM
pH = 7.95 which is mildly basic

So, you have added acid to the solution, and you have calculated that pH - after adding this acid - is higher than 7 (so the solution becomes basic), and you don't hear alarm buzzing "something went wrong" in your head?

Since ethanol is monoprotic (as far as dissociation is concerned) [A-] = [H+]

Think about it once again.

Hint: what is pH of 10-8M HCl?
Title: Re: pH of Ethanol
Post by: maya.mnkr on October 16, 2011, 11:20:05 AM
hey Borek i see this post now and i still don't understand how to calculate the pH of ethanol..

i know that for your example of HCl the answer is more complicated then just say pH=8 because you have to consider the dissociation of the water as well 
Title: Re: pH of Ethanol
Post by: Borek on October 16, 2011, 02:14:24 PM
i know that for your example of HCl the answer is more complicated then just say pH=8 because you have to consider the dissociation of the water as well 

Same with ethanol solution.

Title: Re: pH of Ethanol
Post by: Richakachad on January 07, 2012, 05:33:58 PM
i know that for your example of HCl the answer is more complicated then just say pH=8 because you have to consider the dissociation of the water as well 

Same with ethanol solution.



And they did. They calculated the pH of 1M EtOH from the pKa of Ethanol. Remember the derivation of pKa is:

HB + H2O  ::equil:: H3O+ + B-

Rate = [H3O+][B-] / [HB][H2O]

[H2O] is a constant and [H3O+] is equivalent to [H+] we can go on to say:

Rate x [H2O] = [H+][B-] / [HB]

Or if we let Rate x [H2O] be a new constant Ka:

Ka = [H+][B-] / [HB]

pKa = -log10(Ka)

The reason your HCl question doesn't work is that the pH = -log10[HCl] formula is an approximation that relies on the assumptions that HCl disassociates entirely into its constituent ions so [H+]HCl = [HCl] (which is more or less true at STP) and that at high concentrations (ie: above ~ 1mM) the contribution of H+ to the solution from the water compared to that of the HCl is negligible.

But either way:

[H+]HCl = 10-8
[H+]H2O = 10-7

[H+]Total = 1.1 x 10-7

pH = 6.96

At 1uM HCl pH = 6.996
At 100nM   pH = 6.9996

Etc, etc...

To sum up: EtOH is slightly basic - which makes sense on another level as:

CH3CH2OH2+ is going to be much easier formed than CH3CH2O- due to the lone pairs on the Oxygen attacking the occasional H+ ion it comes across.
Title: Re: pH of Ethanol
Post by: Borek on January 08, 2012, 09:38:49 AM
And they did.

No idea what you mean by that. Who "they", and what and where they did?

Quote
The reason your HCl question doesn't work is that the pH = -log10[HCl] formula is an approximation

Show where I used such a formula.

Quote
[H+]HCl = 10-8
[H+]H2O = 10-7

[H+]Total = 1.1 x 10-7

Yes, you have to account for H+ from water autodissociation. It was already mentioned several times in the thread. But you can't directly add these concentrations - presence of H+ from HCl dissociation shifts water autodissociation equilibrium to the left.