Molar solution

Molar solution

A molar solution is one that contains one mole of solute per litre of solution.

The phrase may be prefixed with a number to denote other concentrations. For example, a five molar solution of aqueous hydrochloric acid (written as "5M HCl (aq)") means there are 5 moles of HCl per litre of solution. If the solvent is not mentioned (such as "5M NaOH"), it is safe to assume that the solvent is water or the one most commonly used with that solute.

For work with aqueous solutions, concentrations expressed in terms of "molarity" are most useful when performing stoichiometric calculations since easily measured volumes correspond directly to moles of chemical substances involved.

Particularly when working with dilute (aqueous) solutions at room temperature this measure of concentration is by far the most common one in use. It should however be noted that it involves the "volume" of the solution and volumes depend on temperature and pressure. Therefore a solution that is 0.5000 M at room temperature will have a lower molarity at say 80 °C without removing a single molecule from it. For non-dilute 'real' solutions volume even depends on concentration itself. Other concentration measures are available and used in circumstances where this becomes a problem.


The most common way to prepare a solution of known molarity is:

# decide which volume to prepare and make sure a clean volumetric flask is available
# calculate the number of moles of solute needed in this volume
# weigh off the right amount of solute
# put ("all") the solute in the volumetric flask (use funnel)
# add a bit of solvent to dissolve the solute
# (If any powder sticks to the funnel, rinse it into the flask)
# continue to add solvent, homogenizing and dissolving as you go
# fill the volumetric flask "up to" (and not "over"!) the calibration line
# homogenize thoroughly, applying heat if necessary
# label the flask with a clear label that does not fade or wash off
# cool solution to room temperature, and verify volume
# store, ensuring neither solvent nor solute can evaporate by providing a barrier such as a stopper

If the preparer overshoots the mark, he or she will either properly dispose of solution and start over or find a way to "determine" the actual concentration they made.


The above procedure only works if the solute (solid "or" liquid) is available in pure enough form (and you do not overshoot the mark). For some materials that is a problem. For example NaOH is available in pellets but they are hygroscopic. This means that over time they gain mass by taking up water from the atmosphere. This means that if x grams are weighed off this is not all NaOH but in part water. A solution will then end up being slightly lower in concentration than intended. This is why the actual concentration then needs to be determined by "standardization". Usually this is done by taking a material that can be weighed off precisely like potassium hydrogen phthalate and performing a number of titrations against this acid.

It is also possible to start from a strong stock solution, e.g. a 50% solution by weight and weighing off twice the weight of the desired NaOH and diluting this in a volumetric flask. These stock solutions are hygroscopic as well but not as strongly as the pellets.


*The very concept "concentration" presupposes homogeneity. Once the flask is filled to the mark homogenizing is not so easy.
* "All" of the solute must dissolve, if not the solution will eventually achieve the saturated concentration, not the calculated one.
* Some solutes are slow to dissolve.
* The solvent may already contain solutes that may or may not interfere. Even distilled water will contain CO2, O2 and N2 if left open to air. CO2 is important when dealing with acids and bases, O2 with reducing agents. Boiling the solvent first will expel gases.
* Some solutions are light sensitive or spoil at room temperature. Proper storage is important. Leaving flasks open to the air may spoil the content (Gases get in).


If lower concentrations than the one available are desired these can be made by dilution. A suitable amount of solution is pipetted into a clean flask and solvent added to the mark. It is best to use properly calibrated flasks and pipettes, rather than graduated cylinders or beakers which have graduations that are too imprecise for analytical work.

Automatic pipetters facilitate working with tiny amounts, but they should be recalibrated regularly. Note that they are typically designed for aqueous work. The plastic tips may dissolve in some other solvents.

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