- #1
sven222
- 23
- 1
So I have an chem assignment which I've been going through for the past week or so, but the last question has provided somewhat of a stumbling block. Here it is,
So from the question, I have gathered the following,
High pH solutions absorb at higher wavelenghts. Whereas low pH solutions absorb at lower wavelengths. Given the equation C6H5OH --> C6H5O- + H+, one would assume that when the equilibrium of the reaction lies to the right, then the wavelength of the absorbance would decrease, because the concentration of H+ has increased (ie lower pH).
Now, if this is the case, then why? Does the increased H+ concentration stabilise the ground state as opposed to the excited electronic state? This would explain the decrease in wavelength (ie hypsochromic (blue) shift), which in turn is an increase in energy. In other words, more energy is required for an electronic excitiation to occur (stabilised ground state by the H+).
That sounds alright, but then I realized that phenol absorbs at around 270 nm. According to my theory, it should absorb at 400 nm (equil. would lie to the left, thus less H+, thus higher pH, thus high wavelength). And then it turns out the phenolate ion absorbs at about 280 nm, not the 400 nm that one would expect from the initial data. This begs the question, what causes the 400 nm absorption?
I think I'm almost there, but I've tangled myself up in all of this information. I'm not looking for a direct answer, but rather just a push in the right direction. If anyone could do that, it would be much appreciated :)
At pH 13, the absorbance of a particular phenolsolution is 1.5 at 400 nm and 0.0 at 270 nm. At pH 4, the values for the solution of the same cocentrations are 0.0 and 1.0 at these two wavelengths respectively. At pH 9, the values are 0.9 and 0.4 respectively.
(a) Explain these spectral changes
(b) Calculate the pKa of the phenol.
etc etc (few more related Q's)
So from the question, I have gathered the following,
High pH solutions absorb at higher wavelenghts. Whereas low pH solutions absorb at lower wavelengths. Given the equation C6H5OH --> C6H5O- + H+, one would assume that when the equilibrium of the reaction lies to the right, then the wavelength of the absorbance would decrease, because the concentration of H+ has increased (ie lower pH).
Now, if this is the case, then why? Does the increased H+ concentration stabilise the ground state as opposed to the excited electronic state? This would explain the decrease in wavelength (ie hypsochromic (blue) shift), which in turn is an increase in energy. In other words, more energy is required for an electronic excitiation to occur (stabilised ground state by the H+).
That sounds alright, but then I realized that phenol absorbs at around 270 nm. According to my theory, it should absorb at 400 nm (equil. would lie to the left, thus less H+, thus higher pH, thus high wavelength). And then it turns out the phenolate ion absorbs at about 280 nm, not the 400 nm that one would expect from the initial data. This begs the question, what causes the 400 nm absorption?
I think I'm almost there, but I've tangled myself up in all of this information. I'm not looking for a direct answer, but rather just a push in the right direction. If anyone could do that, it would be much appreciated :)