# Volume expansion from phosphate buffer solution

I want to mix a phosphate buffer at $\mathrm{pH\ 7.2}$ at $20~\mathrm{^\circ C}$. I do this by dissolving $0.0036\ \mathrm{mol}$ of sodium dihydrogen phosphate monohydrate ($\ce{NaH2PO4*H2O}$) and $0.0063\ \mathrm{mol}$ of di-sodium hydrogen phosphate dihydrate ($\ce{Na2HPO4*2H2O}$) in water per liter of buffer. I add $0.15~\mathrm{M}\ \ce{NaCl}$. How much water do I add if I want $1~\mathrm{L}$ of buffer? Do I just use $1~\mathrm{L}$ of water — will the salts not expand the volume? Is it insignificant and at what point do I need take into account the volume expansion of the solution?

Notes:

$$M(\ce{NaH2PO4 . H2O}) = 137.99~\mathrm{g/mol} \Longrightarrow 137.99~\mathrm{g/mol} \times 0.0036~\mathrm{mol} = 0.496764~\mathrm{g}$$

$$M(\ce{Na2HPO4 . 2 H2O}) = 177.99~\mathrm{g/mol} \Longrightarrow 177.99~\mathrm{g/mol} \times 0.0063~\mathrm{mol} = 1.121337~\mathrm{g}$$

$$M(\ce{NaCl}) = 58.44\ \mathrm{g/mol} \Longrightarrow 58.44\ \mathrm{g/mol} \times 0.15~\mathrm{mol/L} = 8.766~\mathrm{g/L}$$

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– Jan
Sep 21, 2016 at 15:43

There is an exact way of doing this and there is a practical way of doing this.

The exact way involves a volumetric $1~\mathrm{l}$ flask. You would first fill the salts into the volumetric flash, add some water to dissolve them and then carefully add more water until the flask contains exactly $1~\mathrm{l}$ of solution. That removes the error of the salts increasing the overall volume.

However, it is also part of the exact way to add a pH-monitoring device before you have finised to make sure the pH is exactly at $7.2$ — and to carefully add phosphoric acid or sodium hydrogen phosphate if the pH is off. (A better solution would be to make sure the pH is over the desired value so phosphoric acid can simply be added; as it is a liquid it can be added and dissolved easier.) For most applications outside of thermodynamics, this is an overkill.

The practical way involves one of the following:

• using an Erlenmeyer flask, beaker or similar with a $1~\mathrm{l}$ mark; filling the salts first and then the solution up to the mark

• simply measuring $1~\mathrm{l}$ of water and adding that.

Why are the practical solutions okay? In most cases, it does not matter whether the molar concentration of, say $\ce{NaCl}$ is exactly $0.15~\mathrm{M}$ or whether it is $0.14~\mathrm{M}$. This rather unimportant experimental error is simply accepted. It also does not really matter whether the pH is exactly $7.2$ or whether it is, say $7.5$. Therefore, one often decides to go the not-so-exact but easier and faster way.

• Thanks for the help! Do I add NaCl after i mixed the salts and dH2O? Do you think a magnetic stirrer is necessary? Sep 21, 2016 at 16:26
• @Lennart No, add NaCl with the phosphates. A stirrer is only permitted if you are using the practical way and actually add a physical litre of water, otherwise you are introducing additional error ;)
– Jan
Sep 21, 2016 at 16:27
• Alright. I think I'll go with the practical way :) Regarding the pH.. is it still OK to be not-so-exact when using it in HPLC as an eluent? I'm using SEC- size exclution chromatography Sep 21, 2016 at 16:38
• @Lennart For that it is definitely okay ;) It’s how we make our HPLC-buffers, too ;)
– Jan
Sep 21, 2016 at 16:39
• Ah, ended up with 6.98 pH wanted about 7.2 pH hope it doesnt affect too much.. Sep 22, 2016 at 11:42