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Physical Properties and Intermolecular Bonding in Solids.

Physical Properties and Intermolecular Bonding in Solids.

Introduction: (Initial Observation)

 

Most chemicals in their pure form have no use for household consumers. Other manufacturers purchase pure chemicals and make certain compositions that are useful for consumer markets. Before doing that, chemists need to test physical properties of all ingredients to determine if such a composition is possible and what should be the proper order of mixing the ingredients.

For every solid ingredient, chemists may need to know the solubility, melting point, hardness, vapor pressure and conductivity. By learning about the relation between physical properties of a substance and its intermolecular bonding, you may be able to predict the physical properties of different substance just by knowing their chemical formula.

In this project we study the relation between the intermolecular bonding and physical properties of a substance.

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This project guide contains information that you need in order to start your project. If you have any questions or need more support about this project, click on the “Ask Question” button on the top of this page to send me a message.

If you are new in doing science project, click on “How to Start” in the main page. There you will find helpful links that describe different types of science projects, scientific method, variables, hypothesis, graph, abstract and all other general basics that you need to know.

Project advisor

Safety Requirements: Experiments of this project require goggles, protective clothing and a fire safe environment.

Information Gathering:

Find out about different types of intermolecular bonds and the strength of each type of bond. Read books, magazines or ask professionals who might know in order to learn about the effect of bond types on physical properties of substances. Keep track of where you got your information from.

Following are samples of information that you may find.

This experiment was designed to illustrate the properties of the five major categories of solids, namely, ionic, metallic, polar molecular, nonpolar molecular and covalent network. Since observable properties depend upon the type of bonding within the solid, the characteristics of bonding within these solids can also be discussed.

Compounds with IONIC BONDS, potassium chloride and sodium nitrate were our examples, will generally show properties of conductivity in the molten state or in water solution, somewhat high melting points, solubility in water (polar solvent), insolubility in mineral spirits (nonpolar solvent), and a low vapor pressure (no odor). These properties are explained by the presence of strong ionic bonds in the compound, formed by the attractions of oppositely charged ions which are very strong over short distances. These bonds can be broken by melting or dissolving, both of which free the ions from their crystalline structure, thus accounting for their conductivity.

Substances with METALLIC BONDS, iron, copper, lead and copper-zinc alloy were our examples, will generally show properties of conductivity in the solid and liquid states, insolubility in both types of solvents, high melting points and low vapor pressures. The atoms in a metal are arranged in a regular pattern or lattice. The metal is held together by the valence electrons that are free to move through orbits which extend over the entire lattice thus accounting for conductivity in both the solid and liquid states.

A compound described as MOLECULAR, can consist of either polar or nonpolar molecules. Our polar molecular examples were ascorbic acid, sugar and hydroquinone. Nonpolar molecular compounds were p-dichlorobenzene and paraffin. The properties of these two types of molecular solids differ because of the difference in their molecular polarity. The nonpolar solids are not soluble in water (a polar solvent), whereas the polar solids are water-soluble. Since there are no charged units in either type of solid they do not conduct electricity. Nonpolar molecular solids usually have high vapor pressures and low melting points. (See remarks about ascorbic acid under Tips.) The vapor pressure of most polar molecular compounds is lower, so most do not have an odor. Nonpolar molecular compounds are held together by very weak London dispersion forces, while polar molecular compounds are held together by the somewhat stronger dipole-dipole forces or hydrogen bonds.

Compounds which are described as COVALENT NETWORK, silicon dioxide and silicon carbide were our examples, generally have very high melting points and very low vapor pressures; they do not conduct electricity. In a covalent network solid, all of the individual atoms making up the solid are held together in a giant lattice by covalent bonds. This structure creates a very stable system.

Question/ Purpose:

In this experiment you will investigate the physical properties of solid substances. These properties can be correlated to bond type within the solid.

The purpose of this project is to identify the relationship between properties of solids and their intermolecular bond types.

Question:

Can you determine the solubility or estimate melting point of a substance by knowing its intermolecular bonds?

Identify Variables:

When you think you know what variables may be involved, think about ways to change one at a time. If you change more than one at a time, you will not know what variable is causing your observation. Sometimes variables are linked and work together to cause something. At first, try to choose variables that you think act independently of each other.

Independent variable is the type of bonds in a substance.

Dependent variables are solubility, melting point, hardness, vapor pressure and conductivity.

Controlled variables are environmental variables and experimental procedures.

Hypothesis:

Based on your gathered information, make an educated guess about what types of things affect the system you are working with. Identifying variables is necessary before you can make a hypothesis.

Following is a sample hypothesis for melting point, vapor pressure and solubility of solids, but don’t read that. Write your own hypothesis. In your hypothesis also include hardness and conductivity.

Substances with ionic bonds, have a high melting point, low vapor pressure, dissolve in water (polar solvent) and do not dissolve in nonpolar solvents.

Substances with metallic bonds have a high melting point, low vapor pressure and do not dissolve in any solvent.

Substance with polar molecular bonds have a low melting point, low vapor pressure, dissolve in water.

Substance with nonpolar molecular bonds have a low melting point, high vapor pressure, dissolve in nonpolar solvents.

Substances with covalent bonds, have a high melting point, low vapor pressure and do not dissolve in any solvent.

Note that hypothesis does not have to be correct. Later the results of your experiments may support or reject your hypothesis.

You may also choose or use a table like this for your hypothesis, results, and conclusion.

Type of Bond Solubility in water Solubility in nonpolar solvents Melting point hardness conductivity Vapor pressure
Ionic
Metallic
polar molecular
nonpolar molecular
covalent

Experiment Design:

Design an experiment to test each hypothesis. Make a step-by-step list of what you will do to answer each question. This list is called an experimental procedure. For an experiment to give answers you can trust, it must have a “control.” A control is an additional experimental trial or run. It is a separate experiment, done exactly like the others. The only difference is that no experimental variables are changed. A control is a neutral “reference point” for comparison that allows you to see what changing a variable does by comparing it to not changing anything. Dependable controls are sometimes very hard to develop. They can be the hardest part of a project. Without a control you cannot be sure that changing the variable causes your observations. A series of experiments that includes a control is called a “controlled experiment.”

Experiment:

This experiment can be used while studying solids to show the relationship between properties of solids and bond types. The properties studied are solubility, conductivity, relative melting point, hardness, and vapor pressure.

  1. All chemicals used in this experiment may be obtained locally. Common names and possible sources are:
Chemical Name Common Name Store
iron steel wool or iron nails hardware
copper copper pipe hardware
lead lead sinkers sports
petroleum solvent mineral spirits hardware
ascorbic acid vitamin C grocery or drug
potassium chloride salt substitute (Adolf’s)
(You may use sodium chloride or table salt instead.)
grocery
p-dichlorobenzene urinal and toilet bowel deodorizer blocks grocery
paraffin candles or canning wax grocery
sucrose sugar grocery
hydroquinone photo developer photography
sodium nitrate nitrate of soda garden
copper-zinc alloy brass hardware
silicon carbide carborundum hardware
carbon graphite lubricant hardware
silicon dioxide sand hardware

2. Metal jar lids 5 to 10-cm in diameter may be substituted for evaporating dishes.

PROCEDURE

  1. Obtain an evaporating dish and a small test tube. Using a spatula place a pea-sized sample of a solid into the dish (1.0 grams max). Use the bottom of the test tube, carefully, to attempt to grind the solid substance. Use slight pressure at first. Test the volatility of the solid (and thereby its vapor pressure) by cautiously smelling it. Be sure to waft the vapors toward your nose with your hand. Do not directly “sniff” the open dish. Record your observations in each step.
    – Material with high vapor pressure will have an odor even without grinding.
    – Material with low vapor pressure will have an odor after some grinding.
    – Material with very low vapor pressure will have an odor after more high pressure grinding.
    This method can be used to determine if a substance has a high vapor pressure or a low vapor pressure.
  2. Use a multi-meter (Ohm meter) to measure the resistance between two different spots (about 1/4″ apart) on the solid. If you don’t have an Ohm meter, use a 9-V conductivity apparatus to test the solid for electrical conductivity. Do the test by simultaneously touching the bare wire electrodes to the sample while observing the light bulb. Record your observations for each solid. (Conductivity apparatus is a simple electrical circuit with a battery and a light bulb. You can make it yourself.)
  3. Place about 3 grams of a solid in a small boat made from aluminum foil and set it on the hot plate. (Several samples can be done at one time.) Turn on the hot plate and adjust the temperature to a medium setting. Observe the sequence as each solid melts. A solid which melts very rapidly indicates a low melting point. For any solid that melts, test the conductivity of the melt with the conductivity apparatus or an Ohm meter (Make sure wire electrodes do not touch the foil). Discard the solids which melted after checking their conductivity.
  4. If a substance does not readily melt using the above procedure, increase the temperature of the hot plate to its highest setting. Test the conductivity of any sample that now melts as you did above.
  5. If any solid fails to melt at the highest temperature setting, place a 1.0 gram sample of the solid in an evaporating dish and place it on a wire gauze on a ring stand. Using a Bunsen burner flame – gently at first heat the evaporating dish. Note, some solids will not melt even under these conditions. Record all observations.
  6. Place about 3 mL of distilled water in a test tube. Use the tip of your spatula to add a pea-sized sample of solid to the water. Stopper the test tube and shake well. Observe the degree to which the solid dissolves in the water. Test the conductivity of the solution. Repeat for each sample. Record all observations.
  7. Place about 3 mL of mineral spirits in a test tube. Use the tip of your spatula to add a pea-sized sample of solid to the mineral spirits. Stopper the test tube and shake well. Observe the degree to which the solid dissolves in the mineral spirits. Test the conductivity of the solution. Repeat for each sample. Record all observations.

Water solutions may be flushed down the drain. Organic liquids should be collected and disposed of in the manner used by your school system. Solids should be put into waste crocks when cool and disposed of with solid waste.

Take care when using the 9-V conductivity tester to avoid touching both wires simultaneously; a slight shock might result. While testing solubility; thumbs should not be used to stopper test tubes. Take care to keep all organic solvents away from open flames and hot plates. Goggles must be worn throughout the experiment.

Experiment Notes:

  1. Unless you have already studied liquids, it will be important to study solubility tests, specifically what happens when water (a polar solvent) and mineral spirits (a nonpolar solvent) are mixed.
  2. Any simple analog or digital multi-meter, can be used to measure the conductivity of a substance. If you don’t have access to a multi-meter, use a 9-V conductivity tester or build s simple electric circuit for this test. Note that the plastic of electrodes may melt if it comes in contact with very hot material. Keep the contact time low or connect nails as extension to the electrodes.

3. Potassium chloride melts only under extreme conditions, therefore its melting point and conductivity is best illustrated by the following demonstration. Place approximately 3 grams of KCl in a crucible. Set the crucible on a triangle on a ring stand. Place the hottest part of the flame of a Fisher burner against the bottom of the crucible. Heat the crucible until the KCl melts. (This may take up to 15 minutes.) Conclude the demonstration by testing the conductivity of the liquid. When heating the potassium chloride, black spots may appear. These are the decomposition products of any tartaric acid which might be in the sample. This will have no effect on the result and can be ignored.
4. Hydroquinone (m.p. 285°C) is used as an example of a polar molecule because this compound differs from p-dichlorobenzene only in the type of functional group present. The presence of the two -OH groups leads to hydrogen bonding.
5. Ascorbic acid, commonly purchased as Vitamin C tablets at the drug store, can be used to illustrate the properties of a polar molecular compound. This substance shows the solubility and vapor pressure of this group, but since Vitamin C decomposes when heated it cannot be used to illustrate melting point or conductivity.

Materials and Equipment:

Chemicals:
iron*
copper*
lead*
petroleum solvent*
ascorbic acid *
potassium chloride*
p-dichlorobenzene*
paraffin*
sucrose*
hydroquinone*
sodium nitrate*
copper-zinc alloy*
silicon carbide*
carbon *
silicon dioxide*
Equipment: (Everything can be substituted. Jut use your common sense)
evaporating dish*
spatula
small test tubes with solid rubber stoppers
9-V conductivity tester
hot plate
ring stand and ring
wire gauze
wash bottle with water
crucible tongs
crucible
clay triangle
Fisher burner

p-dichlorobenzene is sold as urinal and toilet bowel deodorizer blocks. Blocks are often placed in screen traps for use; however, you can purchase them separately.

Some similar products are labeled non-para or something similar and do not contain paradichlorobenzene.

Para dichlorobenzene is classified as a cancer causing substance in state of California.

Para-dichlorobenzene is also used as moth balls. Traditionally Naphthalene was used as moth balls, but recently many manufacturers have substituted naphthalene by Para dichlorobenzene. In any case read the box and look for ingredients.

Results of Experiment (Observation):

Experiments are often done in series. A series of experiments can be done by changing one variable a different amount each time. A series of experiments is made up of separate experimental “runs.” During each run you make a measurement of how much the variable affected the system under study. For each run, a different amount of change in the variable is used. This produces a different amount of response in the system. You measure this response, or record data, in a table for this purpose. This is considered “raw data” since it has not been processed or interpreted yet. When raw data gets processed mathematically, for example, it becomes results.

Calculations:

No calculation is required for this project.

Summary of Results:

Summarize what happened. This can be in the form of a table of processed numerical data, or graphs. It could also be a written statement of what occurred during experiments.

It is from calculations using recorded data that tables and graphs are made. Studying tables and graphs, we can see trends that tell us how different variables cause our observations. Based on these trends, we can draw conclusions about the system under study. These conclusions help us confirm or deny our original hypothesis. Often, mathematical equations can be made from graphs. These equations allow us to predict how a change will affect the system without the need to do additional experiments. Advanced levels of experimental science rely heavily on graphical and mathematical analysis of data. At this level, science becomes even more interesting and powerful.

Conclusion:

Using the trends in your experimental data and your experimental observations, try to answer your original questions. Is your hypothesis correct? Now is the time to pull together what happened, and assess the experiments you did.

Related Questions & Answers:

What you have learned may allow you to answer other questions. Many questions are related. Several new questions may have occurred to you while doing experiments. You may now be able to understand or verify things that you discovered when gathering information for the project. Questions lead to more questions, which lead to additional hypothesis that need to be tested.

Possible Errors:

If you did not observe anything different than what happened with your control, the variable you changed may not affect the system you are investigating. If you did not observe a consistent, reproducible trend in your series of experimental runs there may be experimental errors affecting your results. The first thing to check is how you are making your measurements. Is the measurement method questionable or unreliable? Maybe you are reading a scale incorrectly, or maybe the measuring instrument is working erratically.

If you determine that experimental errors are influencing your results, carefully rethink the design of your experiments. Review each step of the procedure to find sources of potential errors. If possible, have a scientist review the procedure with you. Sometimes the designer of an experiment can miss the obvious.

References:

Brown, T.E. and LeMay, H.E.,Jr., Chemistry, The Central Science, Prentice-Hall, Englewood Cliffs, N. J., 1981, p. 313. This work describes the bonding in solids.

Masterton, W.L., Slowinski E.J. and Wolford, E.T., Chemistry in the Lab, Holt, Rinehart and Winston, New York, 1980, p. 97. A similar experiment using typical laboratory chemicals is described.

http://www.unm.edu/~dmclaugh/PrinciplesPDF/14_Intermolecular.pdf

http://dl.clackamas.cc.or.us/ch104-10/(17).htm

http://www.chemguide.co.uk/atoms/bonding/vdw.html

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