WHAT HAPPENS WHEN A SALT OR SUGAR DISSOLVES IN WATER

A PEDAGOGICAL HISTORY

 

Kevin C de Berg

Avondale college, Cooranbong NSW 2265, Australia

kdeberg@avondale.edu.au

 

Have you ever had a relaxing salt bath?  If you haven’t it is worth trying particularly if you have aching muscles after a vigorous game of basketball or rugby.  When you throw crystals of salt into the tub water they seem to disappear in the dissolving process.  No more solid salt appears to be present.  Is the salt still present in the water even though you cannot see it?  You can answer this question just by tasting the water.  There is no doubt that the water tastes salty.  As well, if you took a sample of the salty tub water in a tray and let it evaporate over the next couple of days in the sunlight salt crystals could be recovered from the solution.  You may want to try this.  So what do you think happened when the salt dissolved in the water?  It would appear that the salt is still present in the water when it dissolves but in such small size particles that you cannot see them.  We could model this dissolving process like this.

 

 

A similar model would apply to the dissolving of sugar in water.  The sugar particles are broken down by the water into smaller particles that you cannot see with the naked eye.  Let’s explore the dissolving process a little further.  Apart from changing the taste of water, what other properties of water might be changed when you dissolve salt or sugar in water?  Well, over a hundred years ago, scientists discovered that water will freeze at lower temperatures than  zero degrees celsius when you dissolve salt or sugar in water.  If you live in a climate where you get plenty of snow you will know that salt is often spread on the roads to reduce the amount of ice forming so that travelling is safer.  Cars easily slide on ice so it is in the interests of road safety to lower the freezing temperature of water by adding salt to it.  This means there is less chance of ice forming and therefore less risk involved in car travel.  But by how much is the freezing temperature of water lowered when amounts of salt or sugar are added to it?  To answer this question a scientific investigation is required.  But what is involved in a scientific investigation?  Let’s find out by observing how scientists in the past approached this problem.

 

The first recorded scientific investigation of this problem was carried out by a French scientist F.W. Raoult in the 19th century.  Raoult decided that he had to reduce the number of variables as much as possible so that he could decide what factors affected the lowering of the freezing point.

 

In this experiment what do you think can vary from one investigation to the next? Think about this for a moment.

 

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Well, the amount of salt or sugar can vary, the amount of water can vary, and the type of salt or sugar can vary.  Let’s concentrate on sugars first.  A scientist might select one sugar to study first and see how the freezing temperature varies with amount of sugar while keeping the amount of water constant.  Suppose we measured how much the freezing temperature of water was lowered when 1 gram, 2 grams , and 3 grams of glucose was dissolved in 100 grams of water.  You would get results something like that shown in Table 1.

 

Table 1  Freezing point depression for different amounts of glucose in 100 grams of water.

 

 

Amount(grams)

 

Depression (degrees celsius)

 

1

 

0.103

 

2

 

0.204

 

3

 

0.308

 

What happens to the freezing point depression as the amount of sugar changes? As the amount of sugar doubles or triples, how does the depression change?

 

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From the data above you should realise that freezing point depression is proportional to the amount of sugar dissolved.  This tells us a lot about the dissolving process.  Suppose we didn’t control the variables and added 1 gram of glucose to 100 grams of water, 2 grams of glucose to 150 grams of water, and 3 grams of glucose to 80 grams of water and measured the depression in the freezing point.  Results similar to those in Table 2 would result.

 

Table 2.  Freezing point depression for different amounts of glucose in different amounts of water.

 

 

Amount of glucose

 

Amount of water

 

Freezing point depression

 

1 gram

 

100 grams

 

0.103

 

2 grams

 

150 grams

 

0.138

 

3 grams

 

80 grams

 

0.385

 

 

You can see that when we do not minimise the number of variables that can change in a given experiment it is very difficult to extract a relationship between the variables.  The results in Table 2 show that the depression in the freezing temperature does not double when the amount of sugar dissolved doubles because we have allowed the amount of water to vary as well.  This is an important feature of a scientific investigation.

 

In another experiment Raoult decided to keep the amount of sugar and the amount of water constant and see how the freezing point changed from one sugar to the next.  Now we could measure amount as a mass or a number of particles.  As it turns out Raoult dissolved 0.1 moles (6.022 x 1022 particles or 18 grams for glucose and 34.2 grams for cane sugar) of a sugar in 100 grams of water and measured the freezing point of the water.  This was done for glucose and cane sugar.  Raoult then calculated what he called the Molecular Lowering Factor which was a measure of how much the freezing point of water would be depressed if one mole of the sugar had been dissolved in 100 grams of water.  This was obtained by multiplying the freezing point depression for 0.1 moles of sugar by ten to give the depression for one mole of sugar.  He could do this because as we saw in Table 1, the depression was proportional to the amount of sugar dissolved.  Raoult’s results for the two sugars are shown in Table 3.

 

Table 3.

 

 

Sugar

 

Molecular Lowering Factor

 

Glucose

 

18.6

 

Cane sugar

 

18.5

 

What would the molecular lowering  factor have been if Raoult had used 500 grams of water as the standard instead of 100 grams?

 

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Now Raoult could have chosen to use 1 gram of sugar instead of 1 mole of sugar in 100 grams of water as the standard.  This would have led to a gram lowering factor instead of a molecular lowering factor.  That is, the gram lowering factor would have been how much the freezing point would have been depressed for the dissolving of one gram of sugar in 100 grams of water.  The gram lowering factors for the two sugars are shown in Table 4.

 

Table 4.

 

 

Sugar

 

Gram Lowering Factor

 

Glucose

 

0.103

 

Cane sugar

 

0.054

 

What is the difference in the depression relationship when measuring the amount of sugar in grams compared to measuring the amount of sugar in moles?

 

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I hope you can see that when you use moles for the sugars the molecular lowering factors seem to be the same or nearly the same for both sugars.  This is not the case for the gram lowering factors.  That is, the amount by which the freezing point is lowered for different sugars depends on the number of sugar particles dissolved in the water, not the mass of the sugar particles dissolved in the water.  Do you see how our knowledge of the dissolving process has been advanced by taking a scientific approach to the problem?

 

Can you identify from what we have said so far what might constitute a scientific approach as opposed to some other approach?

 

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So Raoult observed that the depression of the freezing point of water for different sugars depends on the number of particles of sugar dissolved in the water rather than the mass of the particles.  Presumably the same principle should work for salts.  So he investigated a range of salts in the same way.  He dissolved 0.1 moles of the salt in 100 grams of water, measured the freezing point, and then determined how much the freezing point would be lowered if one mole of salt had been dissolved in 100 grams of water.  Raoult’s results for a series of salts is shown in Table 5.

 

Table 5.

 

 

Sugar

 

Molecular Lowering Factor

 

Potassium hydroxide

 

35.3

 

Sodium hydroxide

 

36.2

 

Potassium chloride

 

33.6

 

Sodium chloride

 

35.1

 

Potassium nitrate

 

30.8

 

 

What do you think of these results?  Raoult could not explain these results because he was expecting to get a molecular lowering factor somewhat close to the results he got for the sugars, 18.6, because he had dissolved the same number of moles of salt as sugar in the same amount of water.  This often happens in scientific investigations where unexpected results surface like this.  But this means that we need to further our knowledge of the dissolving process yet again to interpret the data.

 

Can you see why the salts might be behaving differently to the sugars?  Write down some ideas you might have at this stage that might help interpret the results.  Study the magnitudes of the numbers because this could be a clue.

 

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Now don’t feel defeated if you found it difficult to interpret the results because it took another five years from when these results were published to when an interpretation was suggested in the scientific literature.  In 1887 Svante Arrhenius, a Swedish chemist, suggested that, since the molecular lowering factor for the salts was approximately double the sugar value, the salt particles must have been splitting in two.  But what exactly were the salts producing when they split in two?  Two of the suggestions made were as follows.  Let’s consider the case of common salt, sodium chloride.

 

 

Case 1

 

NaCl ®  Na  +  Cl

 

Case 2

 

NaCl ®  Na+  +  Cl-

 

 

Now a number of chemists were very outspoken in their opposition to this idea of a salt splitting up in two or dissociating into two parts.

 

Study each case carefully and see if you can think of any criticisms that could be levelled at each case.

 

Case 1: Your objections

 

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Case 2: Your objections

 

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What did some chemists say about Case 1?

If sodium chloride was splitting up into sodium and chlorine, why didn’t one observe a very violent reaction between sodium and water (if you haven’t seen this reaction ask your teacher to demonstrate it carefully to you) and why couldn’t one smell the chlorine present (like a swimming pool smell)?

 

Do you think this is a legitimate objection on the basis of your general chemical knowledge?  Does Case 2 remove these objections?  Explain.

 

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What did some chemists say about Case 2?

If sodium chloride was splitting up into positive sodium ions and negative chloride ions, why didn’t the enormous attractive forces that exist between opposite charges bring these ions back together again to form NaCl?  How could these charges be produced by the simple action of dissolving in water?  Enormous electrical forces would be required to generate these charges.

 

Do you think these are legitimate objections? Explain.

 

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How did Arrhenius react to these objections to his dissociation hypothesis?  He agreed that the objections to Case 1 were justified.  Yes, one would expect a reaction between sodium and water and would expect a chlorine smell akin to that of a swimming pool.  But Arrhenius indicated that these objections couldn’t apply to Case 2 because a sodium ion (Na+) is not the same as elemental sodium (Na) and a chloride ion (Cl-) is not the same as a chlorine atom (Cl) or elemental chlorine (Cl2).

 

Do you agree with his statement? Explain.

 

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Arrhenius was very much in favour of Case 2 and submitted conductivity results to support the notion of the existence of ions in solution.  In a conductivity experiment electricity is passed into a solution and if the solution does not conduct electricity the current falls to zero but if the solution conducts then a nonzero value of current will flow in the circuit.  The fact that salt solutions conducted electricity was supportive in Arrhenius’ view of the existence of positive and negative ions in solution.  Now, some of Arrhenius’ opponents agreed that a solution could not conduct electricity unless charges were present in solution but they suggested that the ions were created by the outside electrical source rather than simply forming from the dissociation of the salt in water.  A conductivity experiment might use equipment like that below.

 

 

 

Do you think, therefore, that the conductivity experiments categorically prove that salts spontaneously split up into their ions in water?  OR Do you think it is better to say that the results of conductivity experiments were consistent with a model suggesting the presence of ions in solution from spontaneous dissociation of salts in water?  Justify your answer.

 

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Professor Henry Armstrong was strongly critical of the ionic dissociation model proposed by Arrhenius.  He called it the “nonsensical hypothesis of ionic dissociation.  It asserts that sodium chloride is so loosely strung together that it falls to pieces when dissolved in water.  No valid motive is suggested for such self-sacrifice and the hypothesis is in entire opposition to the teachings of chemical experience and inapplicable to the explanation of the greater number of facts”.  Some chemists, including Armstrong, thought that the dissociation model was based simply on numerical relations obtained from physical measurements and could not be relied upon to be a genuine chemical model based on chemical experience.  Armstrong argued that Arrhenius and his ionist friends such as Ostwald from Germany had ignored the effect of water in the dissolving process.  Because of the existence of hydrated salts such as Na2CO3.10H2O, Armstrong argued that dissolving was an association process with water rather than a dissociation process in water.  He argued that properties such as the lowering of the freezing point depended on the number of free hydrone molecules (H2O) present in the solution.  When a salt or sugar was dissolved in water the number of free hydrone (H2O) molecules was reduced due to an association with salt or sugar particles.  The almost double molecular lowering factor of the freezing point for salts like sodium chloride, seen in Table 5, was due to two hydrone molecules being affected by the addition of NaCl.  The addition of a sugar particle affected only one hydrone molecule and so gave half the molecular lowering factor of salts like sodium chloride.

 

Do you think Armstrong is overreacting to the dissociation model OR do you think he has some legitimate concerns?

 

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Consider salts like calcium chloride, CaCl2, and barium chloride, BaCl2, which consist of three particles per formula unit.  When Raoult in 1882 determined how much the freezing point of water would be lowered if 1 mole of these salts had been dissolved in 100 grams of water, he obtained the following results shown in Table 6.

 

Table 6.

 

 

Salt

 

Molecular Lowering Factor

 

Calcium chloride

 

49.9

 

Barium Chloride

 

48.6

 

 

How would Arrhenius have explained these results?

 

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How would Armstrong have explained these results?

 

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A solution of calcium chloride also conducts electricity.

How would Arrhenius have explained this result?

 

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How would Armstrong have explained this result?

 

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So a situation arose in chemistry where there was two models for understanding what happens when a salt dissolves in water and each model was vehemently defended by its proponents and vehemently attacked by its opponents.  There was the model of spontaneous dissociation into ions and the model of spontaneous association with water.  Now we don’t have powerful enough microscopes to actually see particles in solution to decide which of the two models is correct.  We have to judge indirectly by measuring properties like freezing point depression and conductivity.  Sometimes we still cannot categorically decide which model is correct even though a range of properties is studied.  Scientists have to hope that there will arise a property which can help us decide which model is more applicable in describing the dissolving process.  Now, it turned out that around 1915 a British scientist by the name of William Bragg was studying the effect of passing X-rays into salt crystals and from the X-ray pictures obtained concluded that no discrete molecules of NaCl existed as such but rather sodium  chloride existed as a three-dimensional lattice of sodium ions and chloride ions.  A picture is shown below.

 

 

 

 

In other words, when sodium chloride is dissolved in water, the water doesn’t create the ions.  It simply separates the ions already present in the crystal from each other.  This meant that in a conductivity experiment the ions already existed in solution most likely before any outside voltage was applied.  This discovery lent strong support to the ionic dissociation model.  Armstrong, however, was horrified that chemists were resorting to the use of physical measurements in chemistry and had this to say about Bragg’s work with sodium chloride and X-rays.  “Chemistry is neither chess nor geometry whatever X-ray physics may be.  Such unjustified aspersion of the molecular character of our most necessary condiment must not be allowed any longer to pass unchallenged.  A little study of the Apostle Paul may be recommended to Professor Bragg as a necessary preliminary even to X-ray work; that science is the pursuit of truth.  It were time that chemists took charge of chemistry once more and protected neophytes against the worship of false gods”.

 

What do you think was motivating Armstrong in this outburst of criticism?

 

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I want you to now cast your minds back to 1884 and consider how strong in your view the evidence of the lowering of freezing point data supported an ionic dissociation model.  To help you assess this I have included the data already discussed with data for other compounds in the one table (Table 7) below.

 

Table 7.  Molecular Lowering Factors of solutions of the following compounds in water.

 

 

Methyl alcohol

 

17.3

 

Hydrochloric acid

 

39.1

 

Ethyl alcohol

 

17.3

 

Nitric acid

 

35.8

 

Glycerol

 

17.1

 

Sulphuric acid

 

38.2

 

Cane sugar

 

18.5

 

Potassium hydroxide

 

35.3

 

Formic acid

 

19.3

 

Sodium hydroxide

 

36.2

 

Phenol

 

15.5

 

Potassium chloride

 

33.6

 

Acetic acid

 

19.0

 

Sodium chloride

 

35.1

 

Butyric acid

 

18.7

 

Calcium chloride

 

49.9

 

Ether

 

16.6

 

Barium chloride

 

48.6

 

Ammonia

 

19.9

 

Potassium nitrate

 

30.8

 

Aniline 

 

15.3

 

Magnesium sulphate

 

19.2

 

Oxalic acid

 

22.9

 

Copper sulphate

 

18.0

 

 

Which data do you think fits the model and which data, if any, doesn’t fit the model?

 

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Now assess how strong the data in the table supports an ionic dissociation model.

 

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Even though Armstrong’s water association model fell out of favour in the early 20th century due partly to the discovery of ionic lattice structures for salts through X-ray analysis, his insistence that water must be considered to have an important role in the dissolving process proved to be correct.  Water was shown to have a high dielectric constant which meant that oppositely charged ions could exist freely in solution without combining due to powerful forces of attraction.  Water, if you like, was able to absorb these powerful forces of attraction up to a point.  If highly concentrated salt solutions were made or solutions containing highly charged ions were made however, the chance of ion-pairs forming could be enhanced.  So, in a concentrated sodium chloride solution for example, the chance of ion-pairs of the type, Na+.Cl-, forming becomes higher.

 

Do these ideas help explain any of the anomalies in Table 7?  Explain.

 

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So, ever since you started high school you have been expected to know that when common salt dissolves in water, the following process occurs:

 

                                               NaCl ® Na+  +  Cl-

 

And for more complicated salts like sodium carbonate:

 

                                      Na2CO3.10H2 ®  2 Na+  +  CO32-  + 10 H2O

 

As you can see from this assignment this idea proved very controversial.  This little project has been designed to show you how this knowledge came about and what evidence was used to support it.  Of course, I have not been able to discuss all the evidence that was used to support the ionic dissociation model but enough has been included to give you an idea.  The project hopefully has helped you understand how scientists think and how they decide between conflicting models.

 

Is the process of deciding between conflicting models simple or complex? Justify your answer.

 

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To what extent does a scientist’s personality affect the acceptance or rejection of a scientific model?  Draw upon examples in this article in your answer as well as other factors.

 

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What might be the roles of evidence and belief in the construction of scientific knowledge?  Draw upon examples in this article in your answer as well as other factors.

 

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