I have been given $\mathrm pK_\mathrm a$ values of an amino group, a carboxyl group and a side chain of cysteine. How can I find the ionic charge on it at different $\mathrm{pH}$ values?


2 Answers 2


For this, you can use the Henderson-Hasselbalch equation. Using the degree of dissociation, $\alpha$, this can be written as $$\mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log\frac{\alpha}{1- \alpha}$$ Rewriting to solve for $\alpha$: $$\alpha = \frac{1}{10^{\mathrm{p}K_\mathrm{a} - \mathrm{pH}} + 1}$$ As stated above, $\alpha$ is the degree of dissociation, meaning the degree at which $\ce{H+}$ is dissociated from the ionizable group. At $\mathrm{pH} = \mathrm{p}K_\mathrm{a}$, $\alpha = 0.5$, meaning that $50\%$ of the ionizable groups in question are deprotonated. If $\mathrm{pH} = \mathrm{p}K_\mathrm{a} + 1$, about $90\%$ of all groups are deprotonated and if $\mathrm{pH} = \mathrm{p}K_\mathrm{a} + 2$, about $99\%$ of all groups are deprotonated. (If you want to, take a look at the shape of the $\alpha$ distribution for $\ce{-SH}$ in cysteine at Wolfram Alpha.)

From a textbook I found the following $\mathrm{p}K_\mathrm{a}$ values for cysteine: \begin{align} \mathrm{p}K_\mathrm{a}(\ce{-COOH}) &= 1.9\\ \mathrm{p}K_\mathrm{a}(\ce{-NH_3+}) &= 8.35\\ \mathrm{p}K_\mathrm{a}(\ce{-SH}) &= 10.5 \end{align}

From these values, $\alpha$ can be calculated for each ionizable group at the desired pH and this will give you the net charge of the amino acid. Upon deprotonation, the following changes in charge occur for the ionizable groups: \begin{array}{lclcr} \ce{-COOH} &:& 0 &\rightarrow &-\\ \ce{-NH_3+}&:& + &\rightarrow &0\\ \ce{-SH} &:& 0 &\rightarrow &-\\ \end{array}

As an example, let's calculate the charge of cysteine at pH 10. Using the Henderson-Hasselbalch in a spreadsheet yields the following $\alpha$ values: \begin{array}{lcr} \ce{-COOH} &:& 0.9999\\ \ce{-NH_3+} &:& 0.9781\\ \ce{-SH} &:& 0.2403\\ \end{array}

This gives the following total charge for cysteine at pH 10: $$(-1)\cdot0.9999+(+1)\cdot(1-0.9781)+(-1)\cdot0.2403=-1.218$$

  • $\begingroup$ How can I calculate pI from this ? $\endgroup$
    – biogirl
    Commented Nov 27, 2013 at 7:00
  • $\begingroup$ No problem! pI is the pH at which the molecular charge is zero. This can be calculated using the tools explained above. The easiest approach would probably be to find pI by calculating total charge at different pH values in a spreadsheet. $\endgroup$
    – oveoyas
    Commented Nov 27, 2013 at 17:24
  • $\begingroup$ So do I have to do this by trial and error ? No other option ? $\endgroup$
    – biogirl
    Commented Nov 27, 2013 at 18:00
  • $\begingroup$ You could use the equations above to derive an analytical expression. What you need is an expression for each of the three alphas (as given above) and the "equation" used to calculate the total charge at the bottom of the answer with total charge set to zero. From this it should be possible to solve for pH, i.e. for pI. Personally, I think using a spreadsheet is easier. $\endgroup$
    – oveoyas
    Commented Nov 27, 2013 at 18:28
  • $\begingroup$ ok so i would have to use variables and solve algebra ...oohooh...I undrstnd why u find spreadsheet easier :P $\endgroup$
    – biogirl
    Commented Nov 27, 2013 at 18:30

I just wanted to show you how you could set up the equation if you don't know that :-). If you find an expression for the ionic charge of one side group as a function of pH, you're well on the way. You just do that for all the side groups, and add them together, and say that overall you want that to equal 0.

Then you can solve for pH in a program that does such a thing. You can also graph this, and find the isoelectric point (pI) by visual inspection. I've done both below, mostly for fun, since oveoyas has provided a sufficient answer :-)

Plot of charge as function of pH


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