[Medic851]

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.

[Borek]

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.

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:

0	7.47
25	7.00
50	6.63
75	6.35
100	6.14

Does it mean that cold water is basic and hot acidic? In all cases [H⁺]=[OH^-]

[Medic851]

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.

[Medic851]

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"

[AWK]

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

[fireemblem555]

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.

[Loyal]

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.

[Borek]

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?

[maya.mnkr]

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 

[Borek]

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.

[Richakachad]

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 pK_a of Ethanol. Remember the derivation of pK_a is:

HB + H₂O  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 K_a:

K_a = [H⁺][B^-] / [HB]

pK_a = -log₁₀(K_a)

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.

[Borek]

And they did.

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

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.

[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.