Anyone who’s studied chemistry will be overly familiar with titrations. It’s an analytical technique that can be used to find the concentration of a solution (the amount of a solute dissolved in it). I put this graphic together primarily to remind my A level students of some of the key aspects of titrations, but as well as being a handy guide for them, it’s a useful introduction to the technique for non-chemists too!
Students are usually introduced to titrations in the context of reactions between acids and alkalis. As many of you might well recall, this is known as a neutralisation reaction. Titration allows us to work out the concentration of, for example, an acid of unknown concentration, by using a fixed volume of it and measuring how much of an alkaline solution of known concentration is needed to react with all of it.
To measure the fixed volume of the acid solution, chemists often use pipettes. These are long glass tubes which measure a fixed volume, and are also more precise than using measuring cylinders. Pipettes are filled using pipette fillers, which come in a variety of shapes and sizes; some older chemists might remember the practice of mouth-pipetting, which is now frowned upon for fairly obvious reasons! Pipettes aren’t just dunked straight in to the bottle of the solution – this could introduce contamination, so we first pour a suitable amount of solution into a beaker, then use the pipette to measure a precise amount of it. Prior to this, the pipette should have been rinsed with distilled water, followed by the solution it is to be filled with, again to avoid contamination.
Once the solution is in the pipette, it’s then transferred to a conical flask. Conical flasks are better than beakers for this procedure because they can be easily swirled without risk of the contents spilling. Pipettes are actually calibrated to retain a very small amount of solution in the tip when emptied, so although it’s tempting to force this out by blowing down the pipette or by squeezing the pipette filler, this is actually detrimental to the results gained from the titration.
Once you’ve got your acid of unknown concentration in the conical flask, it’s time to set up the burette with your alkali of known concentration. Burettes are tall, thin, graduated glass tubes, with a tap at the bottom that can be opened and closed to allow the solution inside to flow out. The scale on its side allows the amount of solution that’s been allowed to flow out to be read off.
As with the pipette, the burette should be rinsed with distilled water followed by the solution it is to be filled with to avoid contamination issues. It’s easiest to fill with a small funnel on top, though you also need to take care when doing this otherwise it’s easy to send your solution fountaining out of the top of your burette! The burette should be filled up to above the zero line, and then lowered down to it by opening the tap. The meniscus, the bottom of the water level in the burette, should be level with the zero mark.
Filling the burette this way is also useful because it means the space under the tap is also filled with liquid. This is important, as the burette is calibrated to include this volume. If you don’t do this, or if an air bubble is present in the space under the tap, the volume of solution you record as being added will be slightly higher than the amount you’ve actually added, leading to incorrect titre values that will affect your calculated results.
With both conical flask and burette filled, you’re ready to start the titration. First, an indicator is commonly added to the conical flask. For acid-alkali titrations, this is a chemical that undergoes a colour change at certain acidities. Two commonly used examples are phenolphthalein and methyl orange. The indicator changes colour at the end point, when all of our acid has reacted with our alkali. Without it, both solutions are colourless, so it would be impossible to tell!
Alkali solution is run from the burette into the acid solution in the conical flask, swirling the flask as it is added. When the end point is reached, the burette tap is closed, and the volume of alkali added is recorded. A white tile can be placed underneath the conical flask to aid with the ease of spotting the end point colour change. The volume of alkali added is referred to as the titre value; multiple titres are usually taken until concordant results are obtained. These are results that are no more than 0.10 centimetres cubed away from each other.
Once concordant results are obtained, we now know the volume of our known concentration alkali needed to react with a known volume of our acid of unknown concentration. We can now use this to work out the acid’s concentration. To do this, we use a very simple equation: n = cv. In this equation, n is equal to the number of moles, c is the concentration in moles per decimetre cubed, and v is the volume in decimetres cubed.
First, we’ll work out the number of moles of alkali we added from the burette. We know the volume we added, and the concentration of the solution, so we simply multiply these together to find the number of moles. Note that there is one catch: in this equation, volume needs to be in decimetres cubed, not centimetres cubed, so we’ll need to divide our volume in centimetres cubed by 1000 to get it into decimetres cubed. Then it’s good to go in the equation.
Once we know how many moles of alkali we’ve used, we need to know the equation for the reaction so we can work out how many moles of acid it should react with. Let’s take the reaction between sodium hydroxide and hydrochloric acid as an example:
What does this equation tell us? Well, there aren’t any numbers in front of any of the chemical formula, which means that it reads: “1 mole of sodium hydroxide reacts with 1 mole of hydrochloric acid to give one mole of sodium chloride plus one mole of water.” This tells us that the sodium hydroxide and the hydrochloric acid react in a one-to-one ratio. However many moles of sodium hydroxide we had, it reacts with the same number of moles of our hydrochloric acid.
So, we now know the number of moles of hydrochloric acid, and the volume we used. Now, we simply need to use another form of the n = cv equation to calculate the acid’s concentration. By rearranging it we arrive at c = n/v; plugging in the numbers, first remembering to convert the volume to decimetres cubed from centimetres cubed, gives us our unknown acid’s concentration. Job done!
Of course, chemists will be at pains to point out that acid-base titrations are by no means the only use for the titration technique – they can be used for a number of other reactions too, but that’s beyond the remit of this post. There are also back titrations, which are a whole different beast!
Got your own titration tips and tricks you want to share? Drop them into the comments below!
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