In this experiment, we aim to investigate the effect of sodium carbonate on hard water.

Jakub Figurski, 4A

GCSE CHEMISTRY COURSEWORK: INVESTIGATING HARD WATER

1. PLANNING

Aim

In this experiment, we aim to investigate the effect of sodium carbonate on hard water.

Method

Firstly, 25.00cm3 of a sample of water was pipetted into a clean conical flask. The pipette was used because this measures and delivers a very accurate amount of liquid (to the nearest hundredth of a cm3, or 0.01cm3), but care had to taken to make sure it was not held by the bulge, as this could have made the amount inaccurate (as body temperature heats the apparatus up and makes it expand). Also, the pipette is used to avoid spillage, as the thin tube can put the water straight into the appropriate container. Furthermore, a pipette filler was used for safety purposes, because blowing or sucking could get unwanted and potentially dangerous substances into the mouth. 0.500g of sodium carbonate (a white powder) was added to the water, having been carefully weighed on an electronic balance — this was very exact and had a cover to maximise accuracy (it protected the balance from being affected by wind or other particles). The balance measured the mass of the powder to the nearest thousandth of a gram (i.e. 0.001g). Then a burette on a stand was obtained, a considerable amount of soap solution funnelled into it through the top (so the level reached near the top) and 1cm3 of the solution added to the water. The burette (glass tube with measurements marked on the side and a stopcock at the bottom) was used because it very accurately lets out 1cm3 every time. The conical flask was then corked (so no substances were able to escape through the top) and shaken for ten seconds (this timed by a stop-watch). It was observed whether a permanent lather (one that lasted for more than ten seconds — again a stop-watch was used) was created. If so, the result of 1cm3 being added was recorded. If not, then the process of adding a cm3 of soap solution and shaking the flask was repeated, as many times as necessary to produce a permanent lather, and the exact amount of solution needed was noted down. The flask was washed out carefully, once with tap water and duly with distilled water (so that the apparatus was thoroughly cleaned, avoiding any contamination). The experiment was promptly repeated all over again, but this time with other amounts of sodium carbonate (1.000g, 1.500g, 2.000g, 2.500g), and the results recorded also in a table (using a tally chart). Afterwards, the different results of the tests were compared to make out any trends.

Diagrams

Prediction

I predict that the more sodium carbonate is added to the water, the less soap is needed to create a lather. Eventually, though, the powder will lose its effectiveness, due to the fact that all the calcium ions will have already been removed. So these are the results I would expect to obtain, in graph form:

(N.B. The amount of sodium carbonate is on the x-axis because it was the quantity which we controlled, but the amount of soap solution needed was the quantity which we were investigating, and was therefore beyond our control.)

Theory

Now, the explanation of the theory behind the experiment. Soaps are compounds of sodium or potassium. A common soap, and the one used here, is sodium stearate (Na+St- where St- = C17H35COO-). Sodium stearate is soluble in water, but calcium stearate is not. Hard water contains dissolved calcium ions — Ca2+(aq) — which cause the water to form ‘scum’, a solid white precipitate. Hard water also makes soap very difficult to lather, again because of the calcium ions.

So when soap is mixed with hard water, this is what happens:

sodium stearate + calcium hydrogencarbonate 

(soap) (in hard water)

 calcium stearate + sodium hydrogencarbonate

(scum)

2NaSt(aq) + Ca(HCO3)2(aq)  CaSt2(s) + 2NaHCO3(aq)

The ionic equation for this reaction is:

Ca2+(aq) + 2St-(aq)  CaSt2(s)

(N.B. Only ions which make the scum are shown, not the ‘spectator’ ions, such as the ones from 2NaHCO3(aq). Two stearate ions are needed to remove one calcium ion.)

Instead of calcium hydrogencarbonate, calcium sulphate (CaSO4) could be written in the equation:

2NaSt + CaSO4  CaSt2 + Na2SO4

(from soap)(from hard water)(scum)(left in solution)

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Therefore, in the experiment, this is what happened once the sodium carbonate powder (Na2CO3(s)) was added:

CaSO4 + Na2CO3  CaCO3 + Na2SO4

We can see that this is a double-decomposition reaction, with the sodium ‘winning’ the sulphate away from the calcium because it is higher in the reactivity series of elements. So the sodium carbonate effectively lessens the effect of calcium ions in the water, thus softening it and making it easier for a lather to be formed. Therefore, all this must mean that the more sodium carbonate is used, the more calcium ions are removed, the softer the ...

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Therefore, in the experiment, this is what happened once the sodium carbonate powder (Na2CO3(s)) was added:

CaSO4 + Na2CO3  CaCO3 + Na2SO4

The graph (on pg. 3) takes this shape because the more sodium carbonate is added, the less soap solution it takes to give a permanent lather (as the water is made softer and softer). The graph levels off after a certain time, though. This is because eventually there are no more calcium ions to remove by the sodium carbonate. That completes my prediction.

But for the experiment to be fair and successful, there are a number of key factors:

having the correct amount of sodium carbonate (so it must be weighed carefully, with no wind interference) — it is accurate to 0.001g (3 decimal places), so is extremely precise;
having the correct amount of soap solution (so it is made sure that the burette is used carefully and accurately);
having the correct amount of water (so the pipette must be used with great care, and not held by bulge, which renders the amount inaccurate) — correct to 2 decimal places (to 0.01g), again very accurate;
the source of water (so one has to ensure that the same source of water is being used each time the experiment is repeated);
the time taken to shake the conical flask (so a stop-watch is used);
the cleanliness of the apparatus (so the equipment is rinsed out with normal water and then distilled water, to maximise hygiene and prevent contamination);
the temperature of the laboratory (so it is ensured that it remains the same throughout the duration of the tests).

All these variables can be controlled except the last one — but the temperature of the laboratory should be the same anyway, so temperature should not affect the results at all.

Furthermore, there are particular ranges of values which should be considered in the experiment. In regard to the amounts of sodium carbonate tested, there should be a minimum of 0.5g and a maximum of 5g, as this certainly provides us with a suitable range to distinguish any trends in the results. Also, to really have a sufficient total of results for analysis, no less than five different values of sodium carbonate should be tested. Moreover, the whole experiment should be repeated at least once and averages calculated in order to eliminate any anomalous results which are obtained. So, in essence, one can be sure that the experiment is producing reliable (by repeating it to remove fluke results) and accurate (by using the equipment efficiently and carefully, as well as weighing and measuring out substances exactly, to several decimal places — the sodium carbonate is weighed to three decimal places, whilst the water is measured out to two d.p.) results.

Preliminary Experiments

To verify my prediction, some extra preliminary work was done.

Firstly, many different samples of water (e.g. from the laboratory, from Norwich, Evian) were tested by a similar process to the one in this experiment. Same amounts of the types of water were put into conical flasks (using pipettes) and then the soap solution was poured into them via the burette. It was recorded how many cm3 were needed to produce a lather, as in this coursework experiment. These were the results produced:

This shows that the calcium ions have the effect of reducing the amount of lather. This is because the distilled water, which is pure H2O, with nothing dissolved in it — including calcium ions, was proved to be very soft (producing a lather very quickly with the soap solution), whilst the Norwich Tap Water, which has many calcium ions dissolved in it, was very hard (a lather was made only after much soap solution was added). All this information allowed me to confirm my prediction that the more sodium carbonate is used and the more calcium ions can be removed by it, the faster a lather will be seen and so the softer the water will be.

Secondly, and more importantly, another preliminary test was conducted. 25.00cm3 of tapwater was measured out, and 0.500g of sodium carbonate added to it. Then the process was exactly the same as in the coursework experiment, but only samples of 0.500g and 2.000g of sodium carbonate (Na2CO3) were tested, as well as a control with just plain water being tested.

I would therefore expect the results to show a similar trend in the coursework experiment, too, with the hardness being reduced with every amount of sodium carbonate added, and so less soap solution being required each time.

2. OBTAINING EVIDENCE

Safety

Here are the safety precautions which we took doing the experiment:

The soap solution is flammable, so be careful to not ignite it by keeping it away from fire.
Wear safety spectacles at all times to prevent the soap solution getting into the eyes.
Take care with wet glassware (which is slippery so it could slip from the hands and smash).
Take care with the fragile glass pieces of equipment (to prevent glass shattering).
Use the pipette filler to fill the pipette with water (as blowing and sucking could get unwanted and potentially dangerous substances into the body).
Hold the pipette near the end where the filler will go.
When shaking the conical flask, hold it securely and keep a thumb over the bung to stop substances escaping from the top of the flask.
N.S.P. (normal safety precautions) apply.

Results

[N.B. For every individual test, 25.00cm3 of water was used.]

3. ANALYSING AND CONSIDERING EVIDENCE

Analysis

In the experiment, it was seen that the more sodium carbonate was added to the water used, the less soap solution was necessary to produce a lather. This can be best illustrated by a graph of the results (see pg. 12).

The graph makes clear some trends in the results obtained. At first, when the mass of sodium carbonate powder is gradually increased (by 0.500g each time), the amount of soap solution needed to make a lather steadily decreases. Then, after 3.000g of powder is used, the graph begins to level off into a plateau, indicating that however much more sodium carbonate one adds, the quantity of the necessary soap to produce a lather will remain the same. This is explained below.

Water becomes ‘hard’ when it is dissolved with calcium (Ca2+) ions (or magnesium, Mg2+, ions). To soften this water, one has to remove these ions. One method of removing the calcium ions is to add sodium carbonate. When the calcium ions (from the water) and the carbonate ions (from the sodium carbonate) react together, calcium carbonate is formed:

Ca2+(aq) + CO32-(aq)  CaCO3(s)

And this is the equation with all the ions involved:

CaSO4 (aq) + Na2CO3 (aq)  CaCO3 (s) + Na2SO4 (aq)

(calcium sulphate from hard water)(sodium carbonate)(calcium carbonate)(sodium sulphate)

So here the sodium ‘wins’ the sulphate away from the calcium because it is higher in the reactivity series of elements. The result of the reaction is that the soap cannot react well with the calcium ions (because they are ‘locked’ in the insoluble calcium carbonate), and thus this makes it more difficult for ‘scum’ to be formed, and a lather appears faster. So the sodium carbonate effectively lessens the effect of calcium ions in the water, thus softening it and making it easier for a lather to be formed. The calcium carbonate is precipitated and this removes the hardness from the water. Therefore, all this must mean that the more sodium carbonate is used, the more calcium ions are removed, the softer the water is and the more likely it is for a lather to be created on the surface of the water.

When the soap solution is promptly added to this from the burette, the whole experiment takes shape. Soaps are compounds of sodium or potassium. Here, sodium stearate was used. Sodium stearate is soluble in water, but calcium stearate is not. Hard water contains dissolved calcium ions which cause the water to form ‘scum’, a solid white precipitate. Hard water also makes soap difficult to lather, again because of the calcium ions. So when soap is mixed with hard water, this is what happens:

2NaSt + CaSO4  CaSt2 + Na2SO4

(from soap)(from hard water)(scum)(left in solution)

The ionic equation for this reaction is:

Ca2+(aq) + 2St-(aq)  CaSt2(s)

(N.B. Only ions which make the scum are shown, not the ‘spectator’ ions,

such as the ones from Na2SO4. Two stearate ions are needed

to remove one calcium ion.)

Therefore, when the conical flask, in which all the substances used in this experiment are, is shaken, a combination of factors has to be considered. By this time, the sodium carbonate powder has served to soften the water and remove the permanent hardness, by the calcium ions (the cause of hardness) being combined to form calcium carbonate. The calcium ions are the ones which cause precipitate formation when reacting with soap, so by reacting it with the sodium carbonate one makes it into a substance that does not affect the lathering of water. As a result, a lather is more likely to appear quicker (i.e. with less soap) as more sodium carbonate is progressively added — and this is the trend shown by the results.

However, after 3.000g of sodium carbonate was added, the graph started to level off into a plateau which would continue at the 1.50 cm3 of soap solution mark regardless of how much powder was used for the test. This is because the sodium carbonate loses its effectiveness due to the fact that all the calcium ions will have already been removed by it, and so it is just not possible to remove any more, however much powder is added. Therefore, the amount of soap solution ought to stay the same, too, as it indeed does, resulting in the graph levelling off into a plateau.

All this proves that my prediction was quite correct, because I did rightly predict that the more sodium carbonate was added to the water, the less soap would be needed to create a lather, and that the graph of results would even off. On the other hand, my prediction graph envisaged that line of best-fit would see an initial sharp fall, which did not actually happen. So it has been proved that as one gradually increases the quantity of sodium carbonate used, the amount of soap solution needed to produce a lather also gradually decreases, until the plateau is reached.

4. EVALUATING

Evaluation

Overall, the experiment was a success in that it showed what it was supposed to — that sodium carbonate certainly has a definite effect on the amount of lather made (the trend clearly and well seen on the graph, pg. 12), instigating the faster production of it when soap solution is added. Furthermore, the experiment was efficient and easy to use, because only one process was carried out at a time, and the fact that it was done in pairs meant that the tests were conducted quicker, which resulted in factors (such as temperature) staying constant throughout most of the experiment. The tests followed a simple, repetitive process, making them easy to carry out.

The results obtained from the experiment do fit the pattern expected, allowing the conclusion of ‘the more sodium carbonate was added to the water used, the less soap solution was necessary to produce a lather’ to be backed up by good, reliable evidence. Therefore, the graph shows that the line of best-fit goes through, or is very close to, the plotted points, thus showing that indeed as gradually more and more sodium carbonate powder was used, the amount of soap needed for a lather to form steadily decreased for the duration of the experiment. The plateau at the end of the best-fit line also validates a scientific explanation, confirming that

the sodium carbonate loses its effectiveness because all the calcium ions will have already been removed by it, and so it is just not possible to remove any more, however much powder is added (so the amount of soap for making a lather remains the same). So the results were highly successful, supporting all the scientific theories very well; in fact, no totally anomalous results were recorded.

However, the procedure, although producing suitable evidence, was not perfect. Firstly, there was a problem with temperature. Unfortunately, the temperature of the laboratory when we conducted the experiment was beyond our control, and it could have fluctuated and had an effect on the results in the process. Secondly, the temperature of the water used for the tests may not have been entirely constant, either. This was especially true of the start of the experiment on any day, when warmer water tended to come from the taps than later on — this extra heat definitely could have had an effect on the results, as it often plays a major role in affecting the speed of reactions. To combat this, a thermometer could be used to make sure that water temperature stays the same, or some water could be allowed to flow out prior to the experiment, thus making sure that the early warm water does not affect the tests. Thirdly, students were not totally clear on what exactly made up a full lather. Although it was seen whether the lather lasted for ten seconds or more, it was then sometimes still not clear if the precipitate was thick enough to be a proper lather. This problem could be overcome by measuring the lather and making sure it was a certain height (e.g. 0.5cm). Finally, the fact that the experiment was carried out in sections over the course of a couple of weeks was a big failure of the experiment, without doubt. Various conditions could have varied from session to session, so different parts of the experiment were conducted under different conditions. This means that the experiment was not very fair, as a different set of apparatus was used each time the tests were continued from the previous lesson, and this caused stages of the experiment conducted differently, and so results not quite matching the exact pattern of other results could well have been recorded. Therefore, the experiment should be done on one particular day, or at least every pair of students should use the same equipment during every session (so it could be labelled with their names).

The actual results do follow the prediction made closely, and are certainly good enough to support the conclusion that I have drawn, with the graph and its line of best-fit clearly demonstrating unmistakable patterns. Moreover, the slight anomalies seen on the graph can be easily accounted for. If a point is slightly below the best-fit line (e.g. the one circled in green on the graph), then less soap was required to produce a lather than expected for that amount of sodium carbonate. This may have been because slightly more sodium carbonate powder was added, so the water was actually a bit softer than it should have been, so inevitably less soap was needed; also, this may have been because I was not strict enough on what exactly constituted a lather, and thus recorded something as being a lather when it was not actually a full one. If a point is slightly above the best-fit line (e.g. the one circled in orange on the graph), then more soap was required to produce a lather than expected for that amount of sodium carbonate. This may have been again because I was not exactly sure what a real lather was (it was sometimes hard to distinguish between a lather and not quite a lather!), or perhaps because more sodium carbonate powder was added then should have been, so the water was actually harder than it should have been, and more soap was needed for a lather to appear. Also, the anomalies could have been recorded because a different set of equipment was used for different stages — if, for example, one electronic balance used was slightly inaccurate (because of damage, say) and another was absolutely accurate, then minor differences in exact result trends would be seen, and on the graph these appear as wrong anomalies. This is because, if a inaccurate amount of sodium carbonate was added, then the water would become either softer or harder than expected, and so the amount of soap needed to make a lather would rise or fall accordingly, too.

More work could be done to extend this investigation and thus provide more support for the conclusion.

Many different samples of water (e.g. distilled, from the laboratory, from students’ homes) could be tested by a similar process to the one in this experiment. Same amounts of the types of water could be put into conical flasks (using pipettes) and then the soap solution poured into them via the burette. It would be recorded how many cubic centimetres were needed to produce a lather, as in this coursework experiment. This would show that the calcium ions have the effect of reducing the amount of lather. This is because the distilled water, which is pure H2O, with nothing dissolved in it (including calcium ions), would be proved to be very soft (producing a lather very quickly with the soap solution), whilst another type of water, which has many calcium ions dissolved in it, would be very hard (a lather made only after much soap solution was added). All this information would allow me to confirm my conclusion here that the more sodium carbonate is used and the more calcium ions can be removed by it, the faster a lather will be seen and so the softer the water will be.