Volume 1

Individual Substances • Impurities and Purity • Mixtures of Substances • Homogeneous and Heterogeneous Mixtures • Separation of Mixtures and Purification Methods: Distillation and Recrystallization • Crystallization Experiments • Unconventional Separation Techniques • Exercises

1.4.1. Individual Substances. A substance consists of identical molecules of a particular type. For instance, oxygen gas consists of O2 molecules, water of H2O molecules, helium of He atoms (monoatomic molecules), etc. To emphasize that a substance is composed of molecules of just one type, the term an individual substance is often used (also see subsection 1.4.2 below). Some naturally occurring minerals are individual substances, both simple and complex. Examples of simple substances encountered in nature are limited. They include carbon (C) in the form of diamond, gold (Au), sulfur (S), and copper (Cu), as shown in Figure 1-7.
Figure 1-7. Samples of native simple substances. Left to right: carbon (C, in the form of diamond; source), gold (Au; source), sulfur (S; source), and copper (Cu; source).

Naturally occurring individual complex substances (compounds) are more common and include a large number of minerals, some of which are shown in Figure 1-8. Note that the formula of malachite in the caption to Figure 1-8 is written as Cu2CO3(OH)2 rather than Cu2H2CO5. The reason for the parenthesis separating and enclosing "OH" will be explained later.
Figure 1-8. Samples of naturally occurring individual complex substances (compounds). Left to right: quartz (SiO2; source), malachite (Cu2CO3(OH)2; source), and rock salt (NaCl; source).

Individual substances possess properties that are determined by molecules constituting them, regardless of their source. For example, pure white sugar (sucrose, C12H22O11) has exactly the same properties (sweetness, solubility in water, shape of crystals, melting point, etc.) regardless of its source, such as sugarcane, sugar beets or fruit. Some people believe that icing sugar is sweeter than granulated sugar. That is not true. Icing sugar may taste sweeter because it is composed of smaller particles (crystals) that melt (dissolve) in one's mouth faster. However, your coffee or tea would have the same sweetness no matter if you sweeten it with icing sugar or granulated sugar or sugar cubes, so long as you put the same mass amount of any of these commercial forms of sugar in your drink.

1.4.2. Impurities and Purity. If any substance is by definition composed of molecules of only one type, then why do we need the term "individual substance"? In reality, there is no such thing as an absolutely pure substance, regardless of whether it is native or man-made. In addition to the molecules of the main component, any real substance always contains a smaller proportion of molecules of other substances that are called impurities.

Each time you hear or read about a "pure" substance, keep in mind that no real substance can be 100% pure. If this is the case, then how pure should a substance be to be identified as pure or, in other words, individual? It depends on our standards for purity for a particular substance in a particular application. For example, silicon (Si) must be 99.999999999% pure for use in semiconductors. For solar cell applications, silicon can be 100,000 times less pure, "only" 99.9999%, and for metallurgical applications, roughly two billion times less pure, around 98%. So, a sample of silicon that is 99.99% pure would be considered super-pure in metallurgy yet unacceptably impure in the semiconductor and solar cell industries. The production of ultrapure substances can be extremely expensive.

The purity of substances used in chemical research and manufacturing is usually 90-99%. The adjective "individual" is often, but not always, added to the noun "substance" to accentuate the fact that a given substance meets the purity standards for a particular application and/or publication in a scientific journal.

1.4.3. Mixtures of Substances. Matter consisting of different substances (molecules of different type) is called a mixture. For example, if we dissolve sugar in water, the resultant syrup is a mixture of sugar and water. Plain white vinegar is a mixture of acetic acid and water. Portland cement is a mixture of mainly two substances, calcium oxide (CaO) and silicon oxide (SiO2) with much smaller quantities of aluminum oxide (Al2O3) and some other compounds.

It is not uncommon for the beginner to be confused about the difference between a simple substance, a complex substance, and a mixture. Let us repeat, again, that a simple substance consists of identical molecules made up of only one element (type of atom). A complex substance also consists of identical molecules, but those molecules are made up of two or more different elements. A mixture consists of different molecules.

Most of the matter around us is in the form of mixtures. We are very complex mixtures of myriads of chemicals that are constantly reacting. Air is a mixture of gases, primarily nitrogen (N2) and oxygen (O2) with much smaller amounts of argon (Ar), carbon dioxide (CO2), water vapors, and some other minor components. Naturally occurring waters are mixtures of water and various salts. Sea water is salty because of the sodium chloride (NaCl, table salt) and other salts dissolved in it. Milk is a mixture of water with many inorganic and organic compounds, including proteins, lipids (fats), carbohydrates, and mineral salts.

Both pure substances and their mixtures have many important uses. For example, pure active ingredients of many drugs and crop protection agents are often mixed with other substances on purpose, in order to enhance the beneficial effect. Another example is alloys, mixtures of metals, which are often superior to the individual metal components. The most widely used alloy is steel that has a higher tensile strength than its main individual component, iron.

Sometimes pure substances are mixed with other substances for purely marketing purposes. A typical example is brown sugar, which is conventionally manufactured by blending pure white sugar (99.9% pure sucrose) with molasses that gives the product the characteristic light brown color. Brown sugar is believed to be a healthier form of sugar than refined white sugar and, as a result, has a higher market value. As a matter of fact, however, there is virtually no difference between brown and white sugar in terms of nutrition and health effects.

1.4.4. Homogeneous and Heterogeneous Mixtures. There are two types of mixtures, homogeneous and heterogeneous. A mixture is homogeneous if it displays a perfectly even distribution (mixing) of its components on a molecular level. Such uniform distribution guarantees the same composition and, consequently, the same properties of the mixture throughout any given sample. For instance, if you fully dissolve sugar in water on thorough stirring or shaking, the resultant solution is a homogeneous mixture. No matter whether you taste the solution closer to the bottom of the container or closer to the top or elsewhere in the entire volume, the sweetness is the same. In contrast, heterogeneous mixtures have compositions and properties that differ in various parts because the mixing is not uniform. Figure 1-9 provides a simple illustration of homogeneous and heterogeneous mixtures.
Figure 1-9. Homogeneous (uniform) and heterogeneous (non-uniform) mixtures (adapted from source).

Mixtures of gases are always homogeneous. Solutions and many alloys are also homogeneous mixtures. A solution is a liquid containing within it two or more substances. For example, brine consists of two complex substances, table salt (NaCl) and water (H2O). Alloys, which are often called "solid solutions", are mixtures of metals obtained on melting, followed by cooling. For example, brass is an alloy of copper (Cu) and zinc (Zn), and bronze is an alloy of copper (Cu) and tin (Sn).

Mixtures prepared by simply blending two or more solids are heterogeneous, no matter how finely ground (powdered) the solids are. Portland cement, which is mentioned above, is a heterogeneous mixture. Black gun powder is a heterogeneous mixture of finely ground potassium nitrate (KNO3; saltpeter), sulfur (S), and carbon (C) in the form of charcoal. Although black gun powder might look homogeneous to the naked eye, under a microscope one can easily see that it is a mixture of white microcrystals of saltpeter, yellow sulfur particles, and black charcoal particles.

Liquid heterogeneous mixtures are also common. Many liquids form homogeneous mixtures (solutions) with other liquids, for example water and rubbing alcohol (isopropanol). In such cases, we say that the two liquids are miscible. There are also liquids that are not soluble in other liquids. For example, cooking oil and water are immiscible (not miscible). Such liquid heterogeneous mixtures are used in liquid motion toys. While milk might look homogeneous, it is in fact a heterogeneous mixture, as can be clearly seen under a microscope (Figure 1-10).
Figure 1-10. Raw milk under a microscope, magnified 800 times (source).

1.4.5. Separation of Mixtures and Purification Methods: Distillation and Recrystallization. Mixtures of substances often need to be separated. Many separation methods exist. One of the oldest ones is distillation. In the process of distillation, a liquid is heated to produce vapors that are then condensed back to a liquid on a cold surface. In this way, seawater that is not drinkable can be converted to fresh water, as schematically shown in Figure 1-11.
Figure 1-11. Distillation as a method to make fresh water from seawater (source).

The boiling point of water is 100 oC, whereas sodium chloride (NaCl) and other salts present in seawater boil at much higher temperatures. For example, sodium chloride (NaCl), the main salt component of seawater boils at 1,413 oC and is therefore not vaporized at 100 oC, the temperature at which water boils. This video illustrates the simple idea of distillation, showing how to make distilled water at home.

While distillation is used to separate and purify liquid substances, the most widely employed technique for purification of solid substances is recrystallization. Every substance has a specific and finite solubility in a particular liquid at a given temperature and pressure. For example, the solubility of table salt (NaCl) in water at room temperature (~23 oC) is approximately 360 g in 1 liter (L). This means that 1 liter of water can dissolve only up to 360 g of NaCl at 23 oC. A solution that contains the maximum possible amount of a given substance at a particular temperature is called a saturated solution. If we leave a saturated solution of table salt in a bowl to sit uncovered for a period of time, some of the water will evaporate. The salt that was originally dissolved in the amount of water that has evaporated precipitates out to give crystals of NaCl. This event is called crystallization from solution.

The phenomenon of crystallization is employed in the purification method called recrystallization. For example, recrystallization is used to produce food-grade table salt from mined salt. Naturally occurring NaCl is often contaminated with sand (silicon dioxide, SiO2) that could scratch and damage the teeth. To remove the sand, the salt from the mine is treated with water. The NaCl dissolves in the water, whereas the insoluble SiO2 does not and is filtered off. After the filtration, the solid-free solution is evaporated to leave crystals of sand-free sodium chloride.

I reemphasize that the terms crystallization and recrystallization have different meanings. Crystallization is a process of the formation of a crystalline substance from its solution. Recrystallization is a purification technique that involves two processes: (1) dissolution of a solid substance to be purified and (2) crystallization of the solid from the resultant solution.

Here is another example of purification by recrystallization. Suppose someone put, by mistake, a teaspoon of white sugar (about 4 g) in your table salt dish containing 100 g of table salt. Can the table salt be purified from the accidentally added sugar? It can. Sugar (C12H22O11) is much more easily soluble in water than salt (NaCl). At room temperature, the solubility in water is 360 g per liter for salt (see above), whereas for sugar it is much higher, about 2,110 g per liter. To purify the salt, we would dissolve the entire mixture in a minimal amount of water and allow the water in the resultant solution to evaporate. As the volume of the water in the solution diminishes, the salt will crystallize out, whereas the much more easily soluble sugar will remain in solution. After some amount of the salt has crystallized out, it can be easily separated from the solution containing both the salt and sugar. Knowing the amounts of NaCl (100 g) and C12H22O11 (4 g) in the mixture and their solubility in water (see above), we can calculate how much pure table salt could be recovered. If you wish, you may try this calculation. [Answer: in theory, approximately 99.3 g of NaCl could be recovered]. Our goal at this point, however, is not to perform such calculations, but rather to understand the very idea of recrystallization as an efficient method to purify soluble solid substances.
Digression. For performing the calculation in the example above, we assumed that the presence of NaCl does not affect the solubility of sugar and that the presence of sugar does not affect the solubility of NaCl. That was a reasonable assumption to make in order to solve the problem. In fact, however, the presence of a substance always either enhances or diminishes the solubility of another substance. In many cases, the difference is very small and can be neglected. In some other cases, however, the change in solubility caused by the presence of another substance is significant. For example, many organic compounds are much more easily soluble in pure water than in a concentrated aqueous solution of NaCl. This effect is widely used in separation, purification, and isolation technologies for some organic compounds.
There is another recrystallization technique that is used most frequently in laboratories and in manufacturing. The solubility of most solids as well as liquids (but not gases!) increases with temperature. Therefore, a solid substance can often be recrystallized by dissolving it in a minimal or close to minimal amount of hot water (or some other solvent) and then cooling the solution. If the substance to be purified by this method is contaminated with an insoluble impurity, such an impurity can be removed by filtering the solution while hot.

Crystallization is a fascinating process. Watching crystals grow can be mesmerizing. There are many excellent videos on YouTube showing how to safely grow beautiful crystals for fun, such as copper sulfate and colored and glowing crystals of ammonium dihydrogen phosphate. Our goal, however, is not only to have fun crystallizing compounds, but also learn how crystallization can be used for purification. The next section below describes three crystallization experiments that you may consider doing yourself.

1.4.6. Crystallization Experiments. Please carefully read and understand the following:

DISCLAIMER: Although most of the experiments described in this subsection and elsewhere in this website are regarded as low hazard, I expressly disclaim all liability for any occurrence, including, but not limited to, damage, injury or death which might arise as consequences of the use of any experiment(s) discussed, listed, described, or otherwise mentioned in the free online course Chemistry from Scratch. Therefore, you assume all the liability and use these experiments at your own risk (see Terms of Use).

Should you decide not to do the experiments, please still read this subsection.

To carry out chemical experiments successfully and safely, we should adhere to certain rules.

(1) Know well the materials and chemicals to be used in the experiments.

(2) Having learned about properties and potential hazards of all of the chemicals to be used, decide what protective gear to wear.

(3) Use your brain. If we are Homo Sapiens, which is the Latin for "wise man", we are supposed to act wisely. In order to act wisely, we should think. Think over your plan for a chemical experiment, pay attention to detail, and make sure that the procedure is clear. Think as you carry out the experiment. Think how to dispose of the materials after the experiment.

A brief description of what is needed to perform the crystallization experiments is presented below.

Chemicals. The compounds you will be recrystallizing are NaCl (sodium chloride or table salt), KNO3 (potassium nitrate or saltpeter), and CuSO4•5H2O (copper sulfate pentahydrate). As we all must know very well the chemicals that we work with, some information on these three compounds is provided below.

NaCl. Sodium chloride, also known as table salt or just salt, is nontoxic, cheap, and available from any food store. Almost all brands of food-grade NaCl contain a tiny quantity of an anti-caking agent (sodium aluminosilicate) in the form of a fine, water-insoluble powder. That is why solutions of table salt in water are slightly opaque (cloudy). Two major types of table salt are sold in the stores, plain and iodized. Iodized salt contains a small amount of an iodine compound added to the NaCl as a supplement, in order to prevent iodine deficiency. The iodine additive to produce iodized salt is one of the four iodine compounds: potassium iodide (KI), potassium iodate (KIO3), sodium iodide (NaI) or sodium iodate (NaIO3). Most commonly, iodized salt contains about 0.005% of potassium iodide (KI).

KNO3. Potassium nitrate (KNO3) is a long-known and widely used chemical compound that is sometimes called saltpeter. Saltpeter used to be found in nature in considerable quantities, particularly in India. That is why KNO3 is often called Indian saltpeter to distinguish it from its sodium analog, NaNO3, which is found native in Chile and Peru and is called Chilean saltpeter. In the old times, saltpeter was used on a large scale for curing meat and in the production of gunpowder, also known as black powder. The consumption was so large that the natural deposits of KNO3 have long been exhausted. These days almost all KNO3 on the market is synthetic, meaning not naturally occurring but made from other chemicals. Potassium nitrate is manufactured on a large scale and used mostly in fertilizers. It has many other applications, such as in fireworks and for stump removal.

Potassium nitrate is a white crystalline solid and is reasonably safe to work with. It is often stated that KNO3 is non-toxic. Indeed, KNO3 is an ingredient of some toothpastes formulated for sensitive teeth. It is a good idea, however, to always keep in mind that "the dose makes the poison". Eating too much of such innocent substances as sugar or table salt can result in poisoning and can be even lethal. Just do not eat KNO3, avoid getting it in your eyes, and wash it off with water if you accidentally spill it on your skin. Keep potassium nitrate in a glass or plastic bottle. Do not store it in tins or metal containers and avoid mixing KNO3 with other materials. You will understand why when we talk about reduction and oxidation processes (Volume 2). Saltpeter for our experiment can be purchased from many sources. An easy and cheap way to obtain saltpeter is to buy Spectracide Stump Remover, which is virtually pure KNO3.

CuSO45H2O. Copper sulfate pentahydrate is a very well-known and widely used compound. First off, what is this strange formula and name? Many compounds crystallize from aqueous solutions in the form of crystal hydrates. A crystal hydrate is not considered to be a mixture of water with the other substance because the molecules of water of crystallization are part of the larger molecular structure.

The number of water molecules constituting such crystals can vary. For example, gypsum is also a crystal hydrate that contains two molecules of water, CaSO42H2O. These water molecules are called water of crystallization and are shown at the end of the chemical formula of the main compound, followed by a middle dot. For many, but not all, crystal hydrates, the water of crystallization can be removed from the solid by heating.

In the old times, copper sulfate pentahydrate was known under a variety of names, including bluestone, roman or blue vitriol, and vitriol of copper. As you may gather from some of these names, CuSO4•5H2O should be blue in color. And, it is beautifully blue indeed (Figure 1-12).
Figure 1-12. Crystals of CuSO45H2O (source).

Heating blue CuSO4•5H2O results in the loss of the water of crystallization and the formation of white CuSO4 (watch this video). On addition of water to white CuSO4 the blue color returns due to the restoration of the pentahydrate.

As we will learn soon, CuSO4 reacts with iron and many other metals to give metallic copper. Therefore, CuSO4•5H2O in solution or as a solid should not get in contact with metal items.

Copper sulfate is broadly used in agriculture to control fungus diseases of plants and correct copper deficiency in soils. The compound is "moderately toxic" and a mild irritant to skin and eyes. If you spill CuSO4•5H2O on your skin or touch it, just rinse the contact areas with water. By the way, touching chemicals with hands is "bad manners" in chemistry not only because some substances can do harm to the skin, but also because by touching a chemical you could contaminate it with compounds on the surface of your skin. Copper sulfate is readily available from many online vendors and from some hardware stores. It is sold under its chemical name or a trade name, such as Zep Root Kill.

Protective gear. One famous chemistry professor once said that a highly skillful research chemist should not wear a lab coat but rather have his best suit on when working in the laboratory. What the professor meant was that a truly good chemist must know how to handle chemicals safely to avoid exposure to hazardous materials, spills, fires, explosions, etc. If a chemist knows how to work safely, he or she needs only very little protective gear, although in certain cases safety glasses, gloves, a lab coat, and a face shield should be worn.

The most important protective gear for a research chemist is his/her knowledge of chemistry and excellent laboratory skills. We will be developing the basics of both in our course. The experiments described in this section are pretty safe to conduct for the beginner. One does not have to wear gloves or goggles to do these experiments, but you surely can and should if you want to and if it makes you feel safer.

Equipment. You do not need any special equipment to carry out these experiments. It would help a lot, though, if you had a balance. In chemical research, the term "balance" is vastly preferred over the conventional "scale". If you do not have a balance, consider spending around 10 dollars to get yourself a small food scale.

Should you use a balance, do not put any materials to be weighed directly on the pan. Place a sheet of clean paper on the pan, tare the balance, then do the weighing. If you do not intend to use a balance, the recommended approximate quantities of the chemicals may be measured using standard kitchen cups and spoons. Keep in mind that measuring solid materials by spoons is not very accurate, and especially so since the site seems to be silent on whether a heaped, rounded, level, or scant spoon is meant.

Experiment 1. Recrystallization of Table Salt. You can use iodized salt or plain non-iodized salt for this experiment. Say you have iodized table salt. How can you make it iodine-free? Prepare a concentrated (saturated or close to saturated) solution of your iodized salt in water. The solution should be prepared in a vessel with a pouring spout. If you have a laboratory beaker, use it. If you do not, you may use a glass or a clear plastic measuring cup. Put about 4 teaspoons (~43 g) of the iodizied salt in one half-cup of water (~0.12 L) and stir with a spoon for a few minutes until all of the salt has dissolved or no more solid salt seems to be going into the solution.

Why are we unsure if all of the amount of NaCl will dissolve? Do we not know the solubility of NaCl and the quantities of NaCl and water used? We do, but since a teaspoon is not a very accurate measure, 4 teaspoons of NaCl can be considerably more than 43 g or less than 43 g. If it is more than 43 g, not all of the salt will dissolve. If it is less, then our solution will not be saturated and we will have to wait longer for the crystallization to begin. So, if all of the NaCl has dissolved, keep adding table salt in small portions (~0.5 teaspoon) and agitating (stirring) with a spoon until no more dissolution takes place and solid NaCl is clearly seen at the bottom when you stop stirring the mixture. Now your mixture consists of a saturated solution of table salt and solid table salt, and it is time to separate the two by filtration.

For the filtration, we need a clean funnel made of glass, plastic, or metal. If you have a coffee filter that fits your funnel, you can use those. If you do not have a coffee filter but have round filter paper, fold it and place in a funnel, as shown in Figure 1-13.
Figure 1-13. Making a filter using round filter paper (source).

If you do not have special round filter paper, you can make a paper filter cone using sheets of a thick paper towel, as shown in Figure 1-14. If you wish, you may use a different folding method shown in this video.
Figure 1-14. Making a filter using regular filter paper or a kitchen paper towel.

Alternatively, you may use cotton wool (cotton balls) in place of porous paper for the filtration. If you decide to use those, you first have to shape your cotton wool in such a way that it has a 'tail' that can go inside the stem of the funnel. Gently push down the cotton tail into the stem using a narrow spatula, or a pencil, or a coffee/tea wooden stir stick, or any rod that is smaller in diameter than the inner diameter of the funnel stem. Continue gently pushing the cotton down the stem in small movements until reasonably tightly packed. Do not pack too tight, though, and do not apply force, especially if the funnel is made of glass because it may break.

Now that you have both your iodized table salt solution and your filter in place, we need a clean and dust-free receiving container for the filtrate. Filtrate is the solution obtained as a result of a filtration process. Figure 1-15 displays the simplest laboratory filtration design (left) and a homemade filtration setup built using a kitchen glass jar, a plastic funnel, and a coffee filter (right).
Figure 1-15. Left: A simple laboratory filtration setup (source). Right: A homemade filtration apparatus.

If you use cotton wool for the filtration, pour your salt solution into the funnel slowly enough to avoid overflow. If you use a filter paper cone, pour the solution into the cone at such a rate that the liquid never goes above the paper edge. Do not be too disappointed if your filtrate is slightly cloudy. This is due to the presence of the aforementioned anti-caking agent, which is insoluble in water and is added to table salt in the form of such fine powder that the relatively large pores in the filter paper may not be able to catch the smallest particles. In my experience, a tightly enough packed cotton wool filter may trap the anti-caking agent quite efficiently, although the filtration might be slow in this case.

Now we need to leave the filtrate to evaporate. Evaporation is the process in which molecules of the liquid (water in this case) leave the surface for the headspace (space above the liquid). The rate of evaporation depends on many factors, including temperature, pressure, humidity, and surface area of the liquid. Water evaporates faster in dry air at higher temperatures, so placing your filtrate in a dry warm place will produce faster crystal growth. Also, the larger the surface area of the solution, the faster it evaporates. Consider transferring the filtrate to a shallow bowl if you want crystals of NaCl to appear faster. In any case, cover the dish containing the filtrate with a sheet of clean paper towel and consider securing it with a rubber band. This is needed to protect the filtrate from dust as the water evaporates. Why is it a good idea to protect the filtrate from dust? You want larger crystals of NaCl to form from your filtrate because larger crystals are more beautiful and purer than smaller crystals. Dust particles can serve as so-called crystallization centers, which prompt the formation of a larger number of smaller crystals rather than fewer larger crystals.

Now you just have to be patient. Wait for a sufficient amount of the water to evaporate and crystals of NaCl to appear and grow. It is hard to predict how soon crystals will be produced in your experiment. Usually, it takes at least hours or, more often, days before first small crystals of NaCl appear. After the crystallization has occurred (to any extent), the remaining solution over the crystals is called mother liquor.

Here is a question for you. If you used iodized salt for the crystallization experiment and let your solution evaporate to dryness, will the crystals of NaCl produced be iodine-free? No, of course not! The iodine-containing additive will not evaporate or otherwise "disappear". That is why to purify a substance from a soluble impurity by recrystallization, the filtrate must not be taken to dryness. At least some sufficient amount of the mother liquor should be maintained, so that the impurity stays dissolved in it.

There is no need to explain how to dispose of your solutions and crystals from this experiment. Put them down the drain and the filter in the garbage container – it is just table salt!

Experiment 2. Recrystallization of Potassium Nitrate, KNO3. As mentioned above, solubility of solids and liquids, but not gases, usually increases with temperature. (There are some exceptions such as lithium sulfate, Li2SO4, which is slightly less soluble in hot water than in cold water.) The effect of higher solubility at higher temperatures can be modest for some substances, such as NaCl, or can be huge, as is the case with KNO3. Potassium nitrate is vastly more soluble in hot water than in cold water. The solubility of KNO3 in boiling water (100 oC) is 2,440 g/L, over 10 times that at room temperature, 240 g/L.
Figure 1-16. Temperature dependence of solubility of KNO3 (saltpeter) and NaCl (table salt) in water (source).

Figure 1-16 points to a huge difference in the temperature dependence of solubility for NaCl and for KNO3. We have previously used the g/L unit for solubility. The solubility curves in Figure 1-16 are based on a different unit, the amount of KNO3 or NaCl in grams that dissolves in 100 g of water. As density of water is 1 g/mL (mL = milliliter), 100 g of water = 100 mL. Consequently, solubility in 100 g of water = solubility in 100 mL of water. Both units, g/L and g/100 mL, are broadly used to express solubility. Figure 1-16 shows that in a temperature range 0 oC – 20 oC, KNO3 is less soluble in water than NaCl. At 25 oC, both salts have about the same solubility. Above 25 oC, however, the solubility of KNO3 increases dramatically, whereas for NaCl this effect is rather weak. If we prepare a saturated solution of NaCl in boiling water and then allow it to cool to room temperature, only a small amount of NaCl will crystallize out. Therefore, this warming-cooling technique is not very efficient for NaCl. In contrast, this method is great for KNO3 because of the huge difference in its solubility in hot water and in colder water.

To recrystallize KNO3, we do not have to use boiling water. Just hot water from the tap is hot enough to perform the experiment nicely. Run hot water from the faucet for a minute or so to ensure the highest temperature. The temperature of tap hot water in many countries is approximately 50 oC, so do not scald yourself! Use a glass cup with a handle. Rinse the cup a few times with the running hot water to warm it up. In this way, the hot water for the recrystallization would not be cooled down significantly by the colder walls of the cup. Pour the running hot water into the cup to approximately one-half of its volume (about 100-120 mL). Right after that, scoop KNO3 in the hot water in the cup and stir the solution. Do it until the solid saltpeter stops dissolving. The solubility graph above suggests that about 70 g of KNO3 dissolves in about 100 mL of water at 45 oC. As one teaspoon of KNO3 is roughly 10 g, you will need about seven teaspoons of saltpeter to be added to the hot water in the cup. Again, this is a very crude estimate, of course, but it gives you an idea of the amount of KNO3 to be put in the hot water until it no longer dissolves. Once this point is reached, start adding small portions of hot water from the faucet to the cup while stirring or swirling until all of the solid KNO3 has gone into the solution. Now your solution is saturated or close to saturation at the current temperature. Let the cup sit at room temperature and watch the formation of white crystals as the temperature of the solution drops and the solubility of KNO3 decreases accordingly.

To dispose of your crystals and solutions, simply put them down the drain without any safety or environmental concerns: as mentioned above, KNO3 is used in some brands of toothpaste, as a fertilizer, and for stump removal.

Experiment 3. Recrystallization of Copper Sulfate Pentahydrate, CuSO45H2O. First and foremost, let me remind you that you must not use anything that is made of metal to agitate your CuSO4 solutions. Moreover, make sure that your CuSO4•5H2O, in solution or as a solid, does not get in contact with any metal items. As we will learn later in our course, CuSO4 reacts with iron and many other metals to give metallic copper.

The solubility of CuSO4•5H2O in water at room temperature is just slightly above 300 g/L. Place 2 teaspoons (made of plastic, not metal!) of CuSO4•5H2O (about 35 g) in ½ cup of water (about 0.12 L) at ambient temperature and agitate for a few minutes with a plastic spoon or wooden popsicle stick. If all of the copper sulfate has dissolved, add more in small portions. Keep adding and stirring until the dissolution stops. Filter the solution into a clean glass jar using a filter, as described in Experiment 1. Tie a piece of thread or string to a popsicle stick or a pencil. Cut the thread so that its end would be hanging approximately in the middle of the jar. Tie a knot at the end of the thread, immerse it in the filtrate, cover the vessel with a sheet of paper towel, and leave undisturbed at room temperature. Inspect your solution from time to time. As the water evaporates, blue crystals of CuSO4•5H2O will grow. You will see something that looks like what is shown on the photo in Figure 1-17.
Figure 1-17. A simple setup for CuSO45H2O crystal growth (source).

Growing crystals is an art. If you are extremely lucky, the crystallization will occur mostly at the knot rather than on the glass walls of the jar to produce larger crystals like the one shown in Figure 1-18. Usually, however, to obtain such large beautiful crystals, one needs to use some crystallization tricks. One such trick is called seeding. A larger, well-shaped crystal is selected from the previous growth, which is then tied to the string and immersed in a saturated solution of the same compound. In this way, further growth occurs on the surface of the seeding crystal rather than on the walls of the crystallization vessel. The walls of the vessel used for such crystallizations should be smooth and not only dust-free, but also desirably free from scratches, chipping, and other defects that can serve as crystallization centers. Details can be found on many web pages such as here and here and in this excellent video.
Figure 1-18. A homegrown single crystal of CuSO45H2O (source).

Your copper sulfate solutions after the experiments may be safely disposed of by putting them down the toilet or sink. In fact, the major application of CuSO4•5H2O in households is to prevent clogging of sewer systems.

Crystal growth is used to make beautiful crafts, and you should feel free to go wild exploring various crystallization methods, inventing your own techniques and tricks to grow nice crystals, and coming up with new ways to decorate crafts with your crystals. The goal of our course, however, is to familiarize you with crystallization as an efficient method to separate and purify chemical substances.

1.4.7. Unconventional Separation Techniques. Distillation and crystallization are undoubtedly the most common and widely used separation and purification methods. There are, however, many other ways to separate and purify compounds. An interesting example is the separation of a mixture of sulfur and iron using a magnet (Figure 1-19). Magnets attract iron but do not attract sulfur. A mixture of sulfur and iron powders cannot be separated by distillation or recrystallization. However, placing a magnet close enough to the mixture separates the two quite nicely, as shown in Figure 1-19 and in this short video.
Figure 1-19. Separation of a mixture of iron (Fe; black particles) and sulfur (S; yellow particles) using a magnet (source: left and right).

In Figure 1-19, note the watch glass placed in between the magnet and the mixture. The watch glass was used for easy separation of the iron powder from the magnet. Lifting up the magnet while holding the watch glass removes the magnetic force from the iron filings, which makes it easy to discharge them into a storage container placed underneath. A sheet of paper or plastic film in place of the watch glass would also do the job.

1.4.8. Exercises.

1. Is a mixture of nitrogen gas (N2) and argon gas (Ar) homogeneous or heterogeneous? Answer

2. A few sugar cubes are dropped in a glass of water. Before all of the sugar has dissolved, the mixture is [homogeneous or heterogeneous]. After thoroughly stirring the mixture, all of the sugar has dissolved to form a uniform clear solution. The resultant mixture of water and sugar is [homogeneous or heterogeneous]. After standing uncovered for a period of time, some of the water evaporated and a small amount of sugar crystals formed, which made the mixture [homogeneous or heterogeneous]. Answer

3. As you heat water to bring it to a boil, long before it starts boiling bubbles appear on the inner surface of the container. These bubbles (a) are the air that was dissolved in the water and that bubbles off as the water gets hot. Like all gases, the air components are less soluble in water and other liquids at higher temperatures; (b) cannot be air because air is insoluble in water. It is steam due to the slow boiling of water at temperatures below its boiling point of 100 oC. Answer

4. There was a sample of NaCl contaminated with sand and a small quantity of KNO3. To purify the NaCl, a technician shook the sample with excess water, then filtered the mixture, and finally left the filtrate to evaporate to dryness. What did the technician do right and what wrong? Answer

5. What do rainwater and distilled water have in common? Answer

6. Suggest a simple method to separate a mixture of iron (Fe) and copper (Cu) filings. Answer

7. A filtrate is (a) a solution to be filtered; (b) a solid to be filtered off; (c) a solution after filtration. Answer

8. What is the difference between a filtrate and a mother liquor? Answer