CRYSTALLINE SOLIDS

This experiment was developed and written by M.L. Jezl, C.R. Landis, D.W. Pollock, and J.L. Scott.

INTRODUCTION

Most solids you encounter on a day to day basis are mixtures of many different kinds of pure substances. Although mixtures are common, they are also more difficult to study analytically. Most of the chemistry you will be learning focuses on pure substances. Pure solids can be classified as crystalline or amorphous. Solids which have the same repeating unit throughout the entire structure are called crystalline. You are probably familiar with some crystalline substances like diamonds, table salt, and sugar. The network of repeating units is often referred to as the "extended lattice." The lab will focus on crystalline pure solids because they often have physical properties which can be related to their repeating atomic structures.

Lab Schedule

In the first session you will prepare the solutions for crystal growth and explore the solubility properties of different classes of crystalline solids. In the second session you will harvest seed crystals and suspend them for larger crystal growth. You will also build models of diamond, graphite, and MoS2 to investigate their `glide plane' properties. In the third session you will study the morphologies of your crystals and find their glide and cleavage planes. A model of the NaCl structure will be used as an aid for identifying cleavage planes. Finally, some of the material properties of the substances used in lab will be tested.

SESSION #1

Objectives:

* Explore solubility as a tool for solid classification.

* Start growing seed crystals for the next session.

Solid Classifications

There are three important subdivisions of crystalline solids. The first two are covalent and ionic networks. Network solids have an undefined size and all atoms are indirectly connected to one another. That is, you can start at any atom and trace a path to any other atom through bonds in the structure. Networks are classified by the kind of bonds that are formed between the atoms in the structure. The bonds in covalent networks are formed by the equal sharing of electrons between two atoms (called covalent bonds). When atoms of different elements form bonds, they don't share the electrons equally. Some elements have a strong pull for electrons when they form bonds resulting in an uneven distribution of electrons around the two atoms. In an extreme case, atoms can sometimes remove electrons from their neighbors. When atoms gain or lose electrons, they become charged atoms called ions. The attractions between oppositely charged ions are called ionic bonds.

The final classification of crystalline solids which will be discussed is molecular extended lattices. The repeating unit for this class are molecules which are packed next to each other. It is important to note that there are no covalent or ionic bonds between the individual molecules -- they are held together by weaker attractions called intermolecular forces.

Figure 1. Classification scheme for pure solids

Solubility

In the first part of this session, you will use solubility as a way of learning about some of the basic properties of the classifications above. When a substance dissolves, we say that it is soluble in that liquid. Chemists usually refer to the liquid part of the solution as the solvent, and the solid as the solute. Whenever you dissolve sugar in coffee, you are using solubility to your advantage, but, you may not have considered how dissolving a substance works. Is the solute destroyed when it dissolves? Can you get the solute back out of the solvent? When a substance dissolves, it means that the solvent molecules are able to surround the solute particles and make them disordered. You can only dissolve a certain amount of a substance before the solvent molecules can no longer effectively surround the solute particles. When the maximum amount of solute has been dissolved, the solution is said to be saturated. The extent of solubility will always depend on the characteristics of the solute and the solvent. Can we make any predictions about solubility?

Consider covalent network solids, such as diamond. We know that these are huge extending structures which are held together by covalent bonds. Do you think solvent molecules could surround and disorder these substances?

Ionic network solids are composed of charged atoms which also extend into huge structures. Many common solvents are partially charged because of their shape and the unequal sharing of electrons between different atoms. Water is an example of such polar solvents. The positive end of water can be attracted to the negative ions in ionic networks. Likewise, the negative end of water can be attracted to the positive ions. One by one, the ions in ionic network solids can be "plucked" away and surrounded by polar solvent molecules and dispersed randomly throughout the solution. Do you think that ionic network solids are soluble in non-polar solvents?

Molecular network solids are more difficult to distinguish by solubility because the individual molecules in these networks can either be polar or non-polar. As a general rule, "like dissolves like" meaning that non-polar molecules dissolve in non-polar solvents and vice versa.

Experimental Procedures:

<Take notes on a brief video depicting an ionic network solid being dissolved by water molecules>

Your lab instructor will have a series of labeled beakers containing a variety of common solvents. Your instructor will attempt to dissolve sodium chloride, graphite, and sugar in each of these solvents. Record detailed observations of the results, and write an explanation for each observation given your knowledge of solubility and solids. You probably don't know whether sugar (sucrose) is a polar or non-polar molecular network solid. Make sure you describe how you can tell from the solubility results.

Table 1. Solvents used in solubility tests

For the rest of the lab you will prepare solutions for crystal growth. In groups of three students, clean and dry three 250 mL beakers. These should be rinsed with distilled water prior to drying. Prepare and label the following solutions:

Beaker# 1. 40g NaNO3 + 50 mL distilled water

2. 20g K3Fe(CN)6 + 50 mL distilled water

3. 20g KAl(SO4)2^.12H2O + 100 mL distilled water

Place each beaker on a hot plate located in the hoods at each bench. Add a stir bar to each beaker and stir the solution continually while heating it using a medium setting. Don't allow too much water to evaporate as this will reduce the amount of solid that will dissolve. When all the solid has gone into solution, remove the beakers from the hot plates and set them aside to cool.

After the solutions have cooled to room temperature, if no crystals have begun to form, drop about 6 flecks of additional solid in the solution. These will serve as crystallization sites. Cover the beakers with a watch glass making sure there remains a small opening at the beaker spout which will allow the water to evaporate slowly. Set the beakers in your lab drawer so they may sit undisturbed until the following week.

Questions

1) Your conclusions from the solubility tests were dependent on having some previous knowledge of the properties of the solute and solvent. How well do you think solubility would work for determining a completely unknown solute? Which kinds of crystalline solids are the easiest to determine with solubility tests? Which are ambiguous?

2) You know that NaNO3, K3Fe(CN)6, and KAl(SO4)2 all dissolve in water. With what you know about solubility, can you classify these pure crystalline solids into the categories in Figure 1? Can any categories be rules out?

SESSION #2

Objectives:

*Prepare seed crystals for alum and potassium ferricyanide.

*Build models to demonstrate glide planes.

Growing larger crystals

The starting materials you used in the first session were the solids in a powdery form. These powders are still crystals, but only very tiny ones. In order for single crystals to grow in an extended orderly fashion they must be allowed to grow slowly so that each new atom can migrate to its favored position. Last session you started growing crystals with saturated solutions of the various solids. The term saturated means that the maximum amount of solute was dissolved in the solvent. Over the week, the solvent evaporated, leaving the solution supersaturated. That is, there was not enough solvent left for the solute to remain dissolved. Instead, the solute reordered itself and formed a solid known as a precipitate. If this reordering process is too fast, you will just get the powder back out of the water. Instead, you probably formed several small crystals on the bottom of your beaker. The goal of this session will be to start growing one large single crystal of each solid.

When crystals of alum and potassium ferricyanide form on the bottom or sides of containers , they are sometimes distorted because the growth of the crystal is restricted. A solution to this problem is to grow your crystal away from the surface of the container in the center of the solution. You will suspend a seed crystal by a thread in the center of your solution to allow unrestricted growth to occur. The crystals of NaNO3 grow with a flat edge all along the beaker surface, so you will not need to suspend these crystals.

Seed crystals are just normal crystals which trigger further crystal growth on top of themselves. Disordered ions are more likely to reorder themselves on a seed crystal (which already has an extending ordered structure) than they are to start a new crystal on their own. As you might imagine, the success of your large single crystal depends greatly on having a high quality seed crystal as a foundation. Each of the solids you will be growing crystals from will have a different shape according to their arrangement of the atoms and how fast each face of the crystal grows.

Glide Planes

Crystalline substances often consistently break along specific planes in the crystal. You will be using two different kinds of modeling kits to explore the origins of glide planes in crystalline substances. Consider the representation of graphite in Figure 2 below. The vertices of the hexagons represent the carbon atoms, and imaginary planes have been superimposed to enhance the visibility of the graphite layers. In reality, these sheets would extend for miles to represent just a pencil point of graphite. A corner of the sheet is enlarged so you can better see the local structure of graphite. The dotted lines in the blown-up representation are meant as a visual aid; they show which atoms line up between layers.

Figure 2. A representation of graphite.

Notice that there are bonds between atoms in the planes, but there are no bonds between layers. You may wonder why graphite layers are held together at all! As a general rule, you can think of all atoms as being slightly sticky. The more surface area present, the more effective these sticky forces become between atoms. Graphite has huge extending sheets of atoms providing enough surface area for the graphite layers to be held together loosely.

Although the sticky forces hold the sheets together, they are not nearly as strong as the covalent bonds between carbon atoms. The difference between the strength of covalent bonds and "atomic stickiness" (by the way, chemists actually call this "van der Waals forces") has important consequences. First, it explains why there is a greater distance between sheets than within the sheets (the weaker attractions don't hold atoms as close). Second, the sheets are able to slip past each other easily because the attractions are weak all along the layers. For this reason, the imaginary planes that run through the graphite sheets are called "glide planes." Glide planes are present anytime the attractions along one plane are weaker than any other planes in the substance. The extensive glide planes in graphite are responsible for the softness and lubricating qualities of graphite.

Experimental Procedures:

Harvesting seed crystals

Inspect the second and third beakers for the best seed crystal of potassium ferricyanide and alum, respectively. Good seed crystals should be approximately 3 mm long and not be stuck to other crystals. Also, make sure there are no discontinuities or breaks within the crystal as this may be an indication of more than one crystal being present (hold the crystal up to the light to assist you in this process). The crystal does not need to be well shaped at this point.

After applying a small amount of Duco-cement to a thread (enough to cover the entire tip of the thread), touch the end firmly to the seed crystal. Allow the cement to harden between the crystal and the thread. In the meantime you may make up your crystal growing solutions and build some model structures (see further below).

Because we consume a great deal of chemicals while making our crystals, we will attempt to reduce our chemical consumption by reusing as much of our solutions as possible. If there is any solution left in the second and third beakers, transfer these solutions to similarly marked 400 mL beakers. You will need solutions with a final total volume of 150 mL each, so for each additional mL of water added to each solution, you must add 0.5g K3Fe(CN)6 or 0.18g KAl(SO4)2^.12H2O. You may use the solid that has already precipitated at the bottom of the earlier beakers as part of the solid added.

As before, stir and heat your solutions until the solids completely dissolve, then set aside to cool. (Warning: It is very important that the solutions are completely cooled before adding the seed crystals. If they are not, the crystals will redissolve). Proceed to the model building portion of this lab.

When the solutions have been completely cooled, bend a piece of copper wire around the lip of the beakers so the watch glasses can still fit on top. It is all right if there is a small space between the beaker and watch glass due to the wire. Tie the opposite end of each seed crystal thread around one of the copper wires and adjust the length of the thread such that the seed crystal will be suspended as close to the center of the solution as possible (Figure 4). You may want to cut the excess thread to avoid dangling it in the solution and having it become an unwanted crystallization site. Again, cover each beaker with a watch glass and let them sit undisturbed in your lab drawer over the next week. Refer to the diagram below to make sure you have the correct setup. The NaNO3 beaker should have crystals starting to grow already, and these should also be left until the next session.

Figure 3. Growing a large single crystal from a seed crystal

Model building - Solid State Kit

Your group should check out two solid state model kits from the stockroom. With the kits, construct the models for diamond (pg. 44 of the manual enclosed in the kit), graphite (pg. 45), and MoS2 (pg. 53). Use the following steps to create your model:

1) Select the appropriate square cardboard template for each compound as specified in the manual, then place this template over the plastic base such that the colored dots match in the same corner.

2) At this point, the holes of the template should align with the holes in the base. Insert the metal rods into the specified holes.

3) Your model will be built up in layers, where each layer is assigned a number by the model kit instructions. You will start by building the layer with the smallest number which may be "zero" or "one" depending on the structure. Find the listing of sphere and spacer types directly below the diagram of the template in the instruction manual. Below the spheres and spacers there should be all the layer numbers you will be using. Match the type of sphere (or spacer) to the layer number you are building. For example, the instructions direct you to use only the colorless spheres in your models of graphite and diamond. But, note that the NaCl and MoS2 structures call for two different types of spheres.

This model kit does not show chemical bonds, so spacers are used to prop-up some atoms to their correct position. Make sure you are using spacers if necessary or your models won't work out correctly!

4) Place all of the spheres in each layer before proceeding to the next. Repeat the first layer (1') in order to finish the crystal structure.

Leave your graphite and MoS2 models constructed so you can refer to them in the next section.

Model Building - Molecular Modeling Kit

Due to limited time and resources, half of the class will build graphite, and the other half will build diamond. Match up with another group of students so you can pool the resources from two model kits. The black spheres in your kits represent carbon atoms, and you will be using them exclusively. The shorter connectors represent single bonds , while two of the longer connectors are used for double bonds. You will need to use parts from more than one kit to build your models, so, please make sure that all of the pieces are returned to your kit at the end of the lab. Each kit should contain 14 black carbon atoms, roughly 40 single bonds, and 12 double bonds.

Graphite

Start by building a ring of carbons in which every other bond is a double bond. Note that your ring of carbons all lie in the same plane. The students who are working on diamond are also building rings of six carbons, but, their carbon atoms lie in different planes. Continue building rings off of each other until you have a structure containing seven fused rings like one of the figures below. Dotted lines have been added to show that the actual structure of graphite continues far beyond the section you are building. The third structure in Figure 4 is blank, and you may use it to construct a blueprint for your structure if you wish.

Figure 4. Examples of bond layouts for the graphite model.

The entire structure should lie in the same plane. Remember, each carbon should have four bonds (except the ones on the edge of your structure). You will quickly realize that you can't continue placing alternating single and double bonds. However, each carbon must still have one double bond, and two single bonds. There are many possible ways you can build your model. It is unlikely that the model you have built is identical to other models in the lab. There are arbitrary decisions you make which lead to different ways of arranging the single and double bonds. Because the arrangement of bonds is arbitrary, your model, and the structures in Figure 4 are called resonance structures. Even though you can come up with many different structures, it has been experimentally determined that all of the bonds in graphite are equivalent! This observation suggests that the best representation of graphite is a blending of resonance structures. If you were to combine all the possible resonance structures, the bonds would all be the same, and would effectively be "4/3 bonds." This situation occurs frequently in chemistry, and oftentimes chemists notate average bonds in graphite and similar structures by placing a circle inside each ring of carbon:

Figure 5. Graphite representation which shows that each carbon is equivalent.

While this kind of representation may seem less precise, it is the most accurate description of how the bonding works in graphite. Find another group which has built a sheet of graphite and using Figure 2 as a reference, overlap the two layers correctly (there is a specific way the layers overlap!).

Diamond

Diamond is a network solid in which each carbon is bonded to four other carbon atoms. While this sounds like a simple arrangement, without some strategy for building your model, it will be difficult to come up with the correct representation of diamond. Start by building a ring of six carbons. Notice that unless you apply a great deal of force, the atoms in the ring do not lie in the same plane, as they do with graphite. There are two basic arrangements of carbon atoms which give the minimum strain on the ring:

Figure 6. The "chair" and "boat" forms of six-membered carbon rings.

If you look at your structure from the side you may see one of these arrangements. The first is commonly called the "chair" form due to its resemblance to a lawn chair. This is the most stable way to arrange a ring of six single-bonded carbon atoms. The other form is called "boat" and is not quite as stable as the chair form. It is easy to switch between the two forms by flipping the edge of your structure (if you can't see this ask your lab instructor for help). For the diamond structure you want to start by building the chair form (in fact, all of the rings of carbon in graphite use the chair form). If you place your chair structure flat on a table, you should see that there are three positions pointing straight up. Continue by building up carbon atoms in these positions. Place a final carbon atom bonded in the middle of the three you just built. Your structure should now look like Figure 7.

Figure 7. Fragment of the diamond structure.

This is a basic building block of diamond. Unfortunately, to link these units together involves removing some of the bonds temporarily. So, disassemble the top four atoms from your model, leaving just the ring of carbons and three bonds sticking straight up. Now, pool parts with another groups working on diamond and fuse together three rings in a triangle to produce a larger base that looks like one the following overhead views in Figure 11 (the o's represent an open place for a carbon to bond straight up):

Figure 8. The two possible overhead views of the base for a larger diamond model. The o's represent locations where bonds will extend upwards.

Make sure that all three rings are in the same plane. If one of the rings points downwards, you haven't connected them correctly. As a group, build the next layer up to complete three structures like the one in Figure 7. Complete the "pyramid" by adding a top layer. Your structure should be large enough that you can continue the repeating pattern on your own. Extend your model by as many carbon atoms as you have left in your kits. Have your lab instructor check your structure when you are done.

Glide Planes in MoS2 (both groups)

Compare the structure of graphite (if you didn't build it, find a group that did) to that of MoS2. Locate the glide plane in MoS2. Remember that the solid state kits do not show bonds between atoms, so you may want to consult with your lab instructor before making up your mind. Draw a diagram of the glide plane in MoS2 in your notebook.

Questions

1) Both the graphite structure and the diamond structure contain rings of six carbon atoms. In graphite, the carbon atoms in each ring all lie in the same plane, while in diamond, the carbon atoms are in different planes (they are staggered). Explain why these rings have different shapes.

2) The distance between two bonded carbon atoms in graphite is about 1.4x10^-10 meters. The distance between the carbons in your graphite model is approximately 0.04 m. So, the scaling factor for the model is roughly 3x10⁸ larger than graphite on the true atomic level. Find the diameter of the model necessary to represent a pencil-point of graphite, if the dot is 1mm in diameter (0.01m) If there are about 1600 meters in a mile, convert your value to miles. Describe a distance that you are familiar with that is comparable to this amount.

3) Diamonds are routinely cut into various shapes to form jewelry pieces. Do you think the cutting process is an example of gliding or cleaving? Defend your choice.

4) Diamonds have the unusual property that they fluoresce (give off light) when subjected to x-rays. This property is exploited in the diamond mining process to detect diamond deposits. Unfortunately, zirconium is another substance which fluoresces in the same way under x-rays, and is often found in diamond mines as well. Speculate on some ways that the zirconium and diamond could be separated after being detected by the x-ray process. You can just brainstorm, you don't need to look up anything about zirconium.

5) Considering diamond and graphite, which structure seems to be more dense? Suppose you are given samples of graphite and diamond. Describe a method that you could use to determine their densities from these samples. Describe another method you could use to determine their densities just from the models you have built. Have your lab instructor give you feedback on your ideas.

SESSION #3

Objectives

* Gather the large single crystals and compare their shapes

* Build models to demonstrate cleavage planes

* Study the properties of some substances studied in the lab

Crystal Shapes

Of the crystals you have grown, NaNO3 has the simplest shape on the atomic level. It is similar to the NaCl structure in that the chloride (Cl^-) ions of NaCl are "replaced" by nitrate (NO3^-) ions. Nitrate is an example of a polyatomic ion, which means that the ion unit is composed of several atoms, not just one atom like chloride (Cl^-). Nitrate ions have a triangular shape which causes the repeating units in NaNO3 to distort from a cube-like shape to a rhombohedron. The diagrams below show the repeating units for NaCl and NaNO3 for comparison. Overhead views of the top and second layer of the repeating pattern are shown as a visual aid (note that the bottom layer is the same as the top layer). The shape of the repeating pattern is also given at the far right. The macroscopic shapes of the crystals you have grown are related to the repeating pattern of atoms, and also on the growth rate of each of the crystal faces.

Figure 9. NaCl structure and shape of its single crystal.

Figure 10. NaNO3 structure and shape of its single crystal.

The shapes of KAl(SO4)2^.12H2O (`alum') and K3Fe(CN)6 (potassium ferricyanide) have more complicated atomic structures (too complicated to explore in this lab) and also more complicated shapes as seen below.

Figure 11. Shapes of 'ideal' crystals of alum and potassium ferricyanide.

Cleavage Planes

Recall that unlike graphite and diamond, sodium chloride is an ionic solid. When you break sodium chloride with an object, it breaks along specific planes known as cleavage planes. These are similar to glide planes, except cleavage planes occur along planes of maximum repulsion. The diagram below shows a cross-section of a simple ionic substance being cleaved.

Figure 12. The cleaving process in a simple ionic substance.

In this diagram, you can think of the large spheres as being chloride ions and the smaller spheres as the sodium ions. So, what happens when you apply a cleaving force? STEP 1 in the diagram is immediately before the cleaving force has been applied. STEP 2 shows atoms shifting downward as a result of the applied force. Note that this position puts the same charges opposite each other. Similar charges repel each other, leading to STEP 3 where the solid breaks along the cleaving plane. You should note that for the diagram above, there is also a cleaving plane horizontally through the substance. What happens if the applied force is not exactly along the cleaving plane? As you might expect, forces not directly along cleaving planes are not as effective. But, if cleaving does occur it will always occur down cleavage planes.

Some solids can contain both cleavage planes and glide planes. Both types of planes are convenient places for the substance to break. The important difference between the two types is what occurs on the microscopic level. Glide planes result due to weak attractions along a plane, while cleavage planes occur when strong repulsions result through motion along the plane.

Materials Testing

You will test the hardness of graphite and diamond which you have been investigating in lab. A convenient (though informal) scale for measuring the hardness of materials is Moh's hardness scale. Some common materials and their Moh hardness value are listed below.

Table 2. Moh's Hardness Scale

Higher numbers correspond to harder substances. When two different substances are rubbed together, the substance with the lower hardness value will be scratched by the other. Diamond is currently accepted as the hardest known material.

Experimental Procedures:

Gathering the Crystals

Remove the crystals you have grown from their solutions. Note their shapes and compare them with the crystals grown by your classmates. Check out a solid state model kit from the stockroom and build the structure for NaCl (pg. 24). Locate a cleaving plane in your model.

Practice crystal cleaving on a piece of rock salt (NaCl). Place a metal spatula or razor blade atop a piece of rock salt and press hard on it (or hit the end with something slightly more massive than the spatula) and note the clean break on each edge that it produces. Now try this with one of the NaNO3 crystals that you grew. To do this you must make sure you use a `single' crystal which has no disruptions within the whole. Search for a couple of the best crystals you grew in order to do the cleaving and gliding for this experiment. Cleave the crystal by laying it on one of its large flat faces, then align the spatula parallel to the slanted side edges and pressing. How do the pieces compare to the original crystal?

Now turn the crystal (or a different NaNO3 crystal) on its edge and hold it firmly with a pair of tweezers. Place a razor blade against one of the obtuse angles (these edges are parallel to the glide planes) and press gently. The planes should gently glide across one another until they form two separate crystals. This may take some practice!

Materials Testing

Try to scratch the surface of a glass slide (quartz) using both a piece of graphite and a diamond-tipped scribe. Use the Moh's hardness scale to rationalize your results. Try to gauge the Moh index for your fingernails.

Rub two ends of tygon tubing together, and note the strong friction. Now, use a piece of graphite to coat the same areas of tygon tubing with a thin graphite layer. Rub these sections together again, and note your observations. Using the graphite spray available, find something in lab which needs lubrication (door hinges, your lab drawer, etc...) and describe how effectively the spray works. Make sure you check the spray bottle for safety hazards and directions.

Place a sample of MoS2 onto a piece of cellophane tape and stick another piece of tape on the other side of the solid. Peel the two pieces of tape apart, and rationalize the results given last sessions discussion of MoS2.

Crystal Growing Questions

1) It was stated that different crystals will have different shapes depending on "... the repeating pattern of atoms, and also on the growth rate of each of the crystal faces." Describe the kind of shapes that you would expect if all the faces of a crystal grew at the same speed. Draw a picture to illustrate several of these shapes. What kind of shapes would you expect if one pair of faces (take two faces on opposite sides) grew twice as fast as the other faces? Draw more pictures to illustrate these shapes.

2) Your lab partner decides that there are two basic crystal types that you have grown in a fictional substance called seaborgium chloride. One kind is rod-like. The other kind has an "x" shape. Your lab partner concludes that the substance grows two fundamentally different shapes of crystals. What do you think? Give an alternate explanation for the two shapes.

3) Your lab partner has somehow broken a piece of seaborgium chloride, but, doesn't have any recollection of the process. Your lab partner believes that since the two pieces of seaborgium chloride look identical that the breaking process must have been due to gliding. Explain if this is necessarily true. What should you do to prove or disprove your lab partner.

Reference (1) Values obtained from the CRC Handbook of Chemistry and Physics, 65th Edition, Robert C. Weast, Melvin J. Astle, and William H. Beyer (editors), CRC Press Inc., Boca Raton, 1984.