Acidity of phosphoric acid

[thushan]

Hmm...how did you heat and cool it?

[thatricksta]

Hmm...how did you heat and cool it?

Heated in a beaker with boiling water and the net backwards reaction occurred vigorously.

We used an ice bath to cool it and when that didn't work I put it in the freezer to see if I could drive the backwards reaction, came out frozen and the same colour hahaha

[thushan]

hmm...it may be possible that colour change does not occur before solution freezes - I don't know what happens with frozen salt solutions...

[thushan]

hmm...it may be possible that colour change does not occur before solution freezes - I don't know what happens with frozen salt solutions...

You could make sure by grabbing some 1 M sulfuric acid and mixing with a few drops of methyl violet indicator, just to make sure it does turn yellow at all...

[Mao]

Ignore above post: i just found out colour change is between pH ~ 0.15 - 3.2

I don't think temperature change was enough to decrease pH to 0.15. Phosphoric acid is somewhat weak; a 0.01 M solution would have pH = 2.25.

Probs concentration of ur solution was quite low, so pH would be much closer to 3 than it is to 0.15.

We had a solution of 1M and its base pH was approximately 1, so that also rules out that... haha, thank you though

That's not quite right.. I'm fairly sure 1M H3PO4 doesn't get near a pH of 1.
I will do some number crunching later tonight to confirm this.


@Thushan, your interpretation of my post is spot on. But I didn't look up the pKa of the indicator.

[thushan]

Solved the quadratic, pH of 1 M phosphoric acid is 1.09 - if Ka = 7.25 E -3.

[thatricksta]

Ignore above post: i just found out colour change is between pH ~ 0.15 - 3.2

I don't think temperature change was enough to decrease pH to 0.15. Phosphoric acid is somewhat weak; a 0.01 M solution would have pH = 2.25.

Probs concentration of ur solution was quite low, so pH would be much closer to 3 than it is to 0.15.

We had a solution of 1M and its base pH was approximately 1, so that also rules out that... haha, thank you though

That's not quite right.. I'm fairly sure 1M H3PO4 doesn't get near a pH of 1.
I will do some number crunching later tonight to confirm this.


@Thushan, your interpretation of my post is spot on. But I didn't look up the pKa of the indicator.

Wiki says it does... hahahaha
Our indicator also suggested it was a pH of 1... green is the mix between blue/violet and yellow, therefore the indicator is in its 'middle'.

[thushan]

yeah, you're right tricksta it'd be 1.09

[Mao]

Ignore above post: i just found out colour change is between pH ~ 0.15 - 3.2

I don't think temperature change was enough to decrease pH to 0.15. Phosphoric acid is somewhat weak; a 0.01 M solution would have pH = 2.25.

Probs concentration of ur solution was quite low, so pH would be much closer to 3 than it is to 0.15.

We had a solution of 1M and its base pH was approximately 1, so that also rules out that... haha, thank you though

That's not quite right.. I'm fairly sure 1M H3PO4 doesn't get near a pH of 1.
I will do some number crunching later tonight to confirm this.


@Thushan, your interpretation of my post is spot on. But I didn't look up the pKa of the indicator.

Wiki says it does... hahahaha
Our indicator also suggested it was a pH of 1... green is the mix between blue/violet and yellow, therefore the indicator is in its 'middle'.

I must still be drunk or something. I read 1M as 0.1M. /continuation of herp-a-derp


Seeing as the %ionization argument is a hit-and-miss, I have crunched some thermodynamics data. Here is a very crude estimate of what happens:

1. Equilibrium concentration of the various species at 1M H3PO4 are:
From this, we note that the second and third ionisation are small, and can be neglected.

2. From the enthalpy of formation data, has (very rough estimate)

3. Applying the Arrhenius equation, the Boltzmann factor increases by a factor of 2.5 as temperature is decreased from 298K to 273K. I.e. Ka1 increases by a factor 2.5. (7.25E-3 M to 1.81E-2 M)

4. Solving for equilibrium concentrations again, we get (increased by a factor of 1.5),

I don't have the Ka of methyl violet handy, but this crude estimate is something to think about: Is a pH change of ~0.2 noticeable by eye?

This of course depends a lot on the properties of the indicator, at the turning point (indicator ratio HI:I is 1:1), this pH change corresponds to a 20% increase of HI and 20% decrease of I-. The further we deviate from this turning point, the change becomes smaller (i.e. at HI:I of 2:1, this change corresponds to 12% increase of HI). This is probably my best bet on why you couldn't visibly see any change in color.

(This result is somewhat dependent on my estimate of the change in enthalpy. If my estimate is out by a factor of 2 (unlikely, as -50 kJ/mol is a quite strong binding force), the pH change approximately doubles.)

[thushan]

Hmm fair enough - btw Mao is it not the Van 't Hoff equation that we're using here? I thought Arrhenius eqn was primarily for kinetics...

[thatricksta]

Thanks
I think that is probably too complex for a 10 mark summary report but I understand how that works I think which is excellent

[Mao]

Hmm fair enough - btw Mao is it not the Van 't Hoff equation that we're using here? I thought Arrhenius eqn was primarily for kinetics...

Haha, I have never heard of that equation. It is the same thing, and easily derived from the Arrhenius equation. (In fact it's not even Arrhenius, it's the Boltzmann factor that we're using, which pops up everywhere)