Learning goals: follow instructions to complete a chemistry experiment, collect experimental data

Introduction

Models provide a useful way of visualizing the arrangement of electrons in a molecule. As you learned in class, there are several types of structure representations used by the chemist at different times to explain chemical phenomena. Today, we will use 3 of these models to explore chemical structures: Lewis dot structures, a chemical modeling kit, and chemical modeling software.

Lewis dot structure

Lewis dot structures allow you to predict molecular arrangements based on the formula. You should be familiar with Lewis dot structures. See an example of a Lewis dot structure on the right.

3D model

Lewis dot structures give no 3D information, so modeling kits can help visualize molecules in 3D. For example, the 3D model makes it obvious why dichloromethane is polar, but this is not obvious looking only at the Lewis dot structure.

ChemDraw is chemical modeling software. It calculates 3D structures and molecular orbitals, which help chemists predict the chemical and physical properties of theoretical molecules. In this lab, you will use ChemDraw to analyze the structures of molecules in 3D.

Background

This infographic can help you review Lewis dot structures. Note: step 6 is optional and only applies to some molecules. Also note: below step 6 are some tips on when to follow the octet rule and when the octet rule may have an exception.

Procedure

For each given molecule, first, construct a Lewis dot structure as your pre-lab. Then, in lab, make a 3D model in ChemDraw. Using the model in ChemDraw and your textbook, identify the electron geometry or molecular shape of the central atom (if it is not obvious which is the central atom, just pick one of the central atoms). Answer the discussion questions using the models and the software. Complete the worksheet in lab. Hint: consider revising your Lewis dot structure if it disagrees with the 3D model, but trust your brain more than the computer.

To make a molecular model

1. Open Chem3D 15.1
2. Click on the white panel to the right of the main window. It is titled “ChemDraw-LiveLink.”
3. Click on the “A” icon to type, then type the formula of your molecule of interest. For example, you enter H₂O as H2O. When you hit “enter,” the molecule should appear in the big blue window. Note: For more advanced structures or when this method doesn’t produce the structure you were expecting, you will need to type it differently (for example: CH2NH instead of CH3N) or use the tools to draw the molecule by hand. Your instructor and/or TA can show you show you how to do this.
4. Optimize the structure by hitting “control-m.” When you do this, the calculated total energy of the model will appear in an output window.
5. There are several things you can do to get a better look at the molecule.
   1. You can click the third button from the left on the top toolbar to rotate the molecule. It looks like a circle with an arrow on it. After you click it, you can use the mouse to rotate the molecule.
   2. You can click “View” on the main menu, then click “Model Display.” This will present you with many options to change the display of the molecule. For example, “Display Mode” gives you more modes. The “Ball & Stick” mode is most common, but “Wire Frame” is convenient for a very complicated molecule, and “Space Filling” is helpful for visualizing atom size differences.

To measure bond lengths and bond angles on the model

After making the model and optimizing the structure, click “Structure” on the main menu, then “Measurements,” and then select “Generate All Bond Lengths” or “Generate All Bond Angles.”

To measure formal charge of atoms in the model

Hover over the atom of interest. If it has a formal charge, it will say “Formal charge: -1,” for example. It will also display delocalized charges.

To predict the shape of the molecular orbitals

After making the model and optimizing the structure, click “Surfaces” on the main menu, then “Choose Surface,” then “Molecular Orbital.” The HOMO (highest occupied molecular orbital) will automatically be shown, but you can choose another molecular orbital from “Select Molecular Orbital.” You will also see the energies associated with each orbital. The labels are with respect to the HOMO and the LUMO (lowest unoccupied molecular orbital). Notice that the HOMO energy is usually negative, indicating a favorable state, but the LUMO energy is usually positive, indicating an unfavorable state (which is why it is UNOCCUPIED!) The difference between the HOMO and the LUMO energies is called the band gap.

Report

Fill out this worksheet and turn in a paper copy. Optional survey.

Learning Goals

1. Build confidence in yourself as an independent experimentalist;
2. Perform chemical degradation analysis, then use the results to calculate an empirical formula;
3. Use a gravity filtration apparatus;
4. Use a drying oven;
5. Use a Bunsen burner to heat a solid;
6. Observe a redox reaction and a precipitation reaction;
7. Identify sources of uncertainty associated with the procedure and the analytical balance.

Introduction

John Dalton

John Dalton was a chemist, physicist, meterologist, and teacher who lived 1766 – 1844 in England. (Can you believe that he was 34 years old when Middlebury College was founded?!) He was born into a poor Quaker family, so he started supporting himself by working as a teacher at age 12 and continued throughout his life. You can thank Dalton for much of the atomic theory that you are learning in Chem 103.

Over the next two weeks, your laboratory experiments will use Dalton’s atomic theory. In today’s lab, you will determine the molecular formula of a copper compound.

Background

Copper is normally in the 0, +1, or +2 oxidation state. This means copper can combine with 0, 1, or 2 chloride ions. Sometimes water molecules cling to cations and form hydrates. It is difficult to predict the number of water molecules that cling to a salt, but it must always be an integer value.

Today, you will determine the empirical formula of a compound that contains copper, chlorine, and water (Cu_xCl_y zH₂O). You will decompose the compound in order to solve for x, y, and z. The decomposition is in two steps. First, you will heat the compound to dehydrate the compound. Then, you will reduce the copper cations to copper metal to drive off the chlorine. By analyzing how the mass changes after each step, you will be able to solve for the mole ratios in the original compound. A tutorial of this lab is found here.

Procedure

Note: do this lab individually, but share a hood with a “hood mate” and consult with them as needed.

Dry the hydrate

1. 1. Set up a ring stand with a clay triangle on the ring in a fume hood. Place a Bunsen burner underneath with a really small blue flame.
   2. Record the mass of a clean, dry crucible in Table 2.
   3. Place about 1 g of the unknown salt in the crucible. Record the mass of the combined crucible and salt in Table 2. Break up any big chunks with a spatula.
   4. Place your salt and crucible on the clay triangle. Heat super gently. Like giving a baby a warm bath. Stir carefully to prevent burning. Observe the color change from blue-green to brown. If your salt begins to spatter or turn black, lower the heat either by turning down the gas or by raising up your sample from the burner. Don’t hesitate to start over if your salt spatters.
   5. Once the salt is completely brown with no signs of yellow or green, remove the heat and let it cool for about 5 minutes.
   6. Record the mass of the cool crucible containing the dehydrated salt in Table 2. If the mass is not stable, let it cool longer.

Reduce the copper salt to elemental copper

1. Transfer the dehydrated salt to a 50 mL beaker. Rinse the crucible twice with approximately 8 mL of DI H2O each time and pour the rinsate into the same beaker as the dehydrated salt to ensure all of the salt is transferred. Swirl to completely dissolve the salt. Note: the salt turns green when it is dissolved in just a little water, and it turns blue when it is fully hydrated.
2. Coil an approximately 10 cm piece of aluminum wire so that it fits into the beaker. Place it in the beaker and ensure it is completely immersed in the solution. Wait about 30 minutes for the copper to deposit on the surface of the wire. It is finished when the solution is colorless and the bubbling slows. While it is reacting, wash and dry the crucible (but keep yours), set up the next part of the experiment, get a drink of water, work on homework, do your pre-lab for next week (you might want to read through the procedure – it’s a tough one, for sure), go ask Prof. Larrabee questions, call someone in your family, run a 5k, take a nap, etc.
3. Use a spatula to scrape off as much copper as you can from the aluminum wire and put it into the solution.
4. Using that filter paper, set up a funnel filtration apparatus on top of a small Erlenmeyer flask similar to the one shown on the right.
5. Slowly pour the copper and the solution onto the filter paper. Use small amounts of water to rinse all of the copper onto the paper.
6. Open the filter paper, label it, and place it on the watch glass. Put it in the drying oven for about 10 minutes until the copper is dry and crumbly and the paper is dry and begins to brown.
7. When your filter paper starts to brown, carefully transfer all of your copper to your crucible. Put it back in the drying oven for 5 minutes.
8. When you check on your copper, break up the large chunks with a spatula to ensure it is dry throughout. Don’t let it turn black.
9. The drying process is complete when the copper is red/brown and the consistency of Grape Nuts cereal. Let the crucible and copper cool, then weigh it. Record the mass in Table 3. Consult your instructor or TA if the mass of copper is greater than 0.40 g.
10. Pour the filtrate down the drain. Rinse, dry, and return the aluminum wire to where you found it. Save your solid copper!

Oxidize elemental copper to copper oxide

1. Return the copper and crucible to the Bunsen heating set up.

2. Heat the copper and crucible over a hot flame and this time: Burn, baby, BURN. Heat until the copper turns completely black.

3. Let it cool until it is cool enough to handle. Weigh the product. Does the mass increase or decrease? What could cause the mass to increase? What could cause the mass to decrease?

Data analysis

To determine the formula, you will first convert the masses to moles, then divide by the smallest number to get the mole ratios. For example, if you were using a hydrocarbon (composed only of hydrogen and carbon) and you found that it contained 0.25 moles of carbon and 0.99 moles of hydrogen, then you would divide both by 0.25:

0.25 moles carbon / 0.25 = 1 mole carbon

0.99 moles hydrogen / 0.25 = 3.96 moles hydrogen, so approximately 4 moles of hydrogen (remember it must be an integer value!)

Therefore, the empirical formula is CH₄.

Report

Fill out this worksheet. Turn in either a paper or digital copy.

Learning Goals

1. Practice using computational chemistry software;
2. Perform flame tests;
3. Use Erlenmeyer flasks to mix solutions;
4. Use a spectroscope to observe atomic emission;
5. Manipulate units of wavelength and energy;
6. Practice writing ground state and excited state electron configurations.

Background

By Comunicaciones CONICYT – María Teresa Ruiz, Premio Nacional de Ciencias Exactas., CC0, https://commons.wikimedia.org/w/index.php?curid=69688240

María Teresa Ruiz is an astronomer born in 1946. In 1997, Ruiz and coworkers discovered a free-floating brown dwarf, which is an unusual celestial body that is halfway between a planet and a star. She used optical spectroscopy to analyze both the light emitted from the brown dwarf and the light it absorbed. This data gave clues about the brown dwarf’s temperature and chemistry. Over the next two weeks, you will learn about optical spectroscopy, both emission and absorption.

Colored chemicals absorb and/or emit light in the visible portion of the electromagnetic spectrum, which has a wavelength of approximately 400 – 700 nm. The color of the absorbed or emitted light depends on the amount of energy the chemical absorbed or emitted. Wavelength and energy are negatively correlated. 

Absorption occurs when an electron in a chemical absorbs energy from the light, temporarily promoting the electron to a higher energy orbital. Light emission can occur when an electron relaxes back to the ground state and produces light, but emission is less common than absorption because there are many of non-radiative ways for the electron to relax. 

Most chemicals are colored because they absorb light and reflect only a portion of incident light. In this case, the color that a chemical absorbs is the opposite of the color that it appears. The color wheel shows you which colors are opposite of one another. The color wheel helps you to predict the color that a chemical absorbs based on the color it appears (and vice versa). For example, beta-carotene, a pigment found in many fruits and vegetables including carrots, absorbs purple and blue light (400 – 500 nm) and reflects all of the other colors, so it appears yellow/orange. It DOES NOT emit yellow/orange light.

If a chemical is colored due to emission, then the color it appears corresponds to the color it emits. Many processes can temporarily promote an electron to a higher energy orbital and induce emission; these include light, heat, electricity, and chemical reactions. Today, you will observe light emission from electricity in part 1, chemical reactions in part 2, and heat in part 3.

Objective

In part 1 of this lab, you will observe emission of atomic hydrogen which will be excited with electricity. Using the Rydberg formula, you should be able to figure out which lines in the emission correspond to which transitions in the hydrogen atom. Refer to Chapter 3.1 in your textbook to refresh yourself on the Rydberg formula. You will also model molecular hydrogen to better understand molecular orbital theory.

In part 2 of this lab, you will observe the emission from a light-producing chemical reaction. You will model the product and run calculations to determine the quantum mechanical source of the emission.

In part 3 of this lab, you will perform flame tests and propose the chemicals you could use to make a fireworks display.

You may perform parts 1, 2, and 3 in any order.

Procedure

Part 1: Hydrogen

Experimental component

Note: perform this section individually.

1. First, look through the spectroscope through the window. You should see a rainbow approximately spanning 4.0 – 8.0 on the scale in the spectroscope. The units on these instruments are in hundreds of nanometers, so 4.0 corresponds to 4.0 x 10² nanometers. If you don’t see a rainbow across this range, ask your TA or instructor for help.
2. Turn on the hydrogen lamp and observe the emission through the spectroscope.
3. Use the Rydberg formula to calculate the energy transition associated with each line. h = 6.626*10^-34 Js, c = 2.9979*10⁸ m/s
4. Turn off light source once you are finished observing the emission. Record your observations.
5. Extra lamps are provided for you to explore, but they are not required for this lab.

Computational component

1. Open Chem3D 18.0
2. Click on the white panel to the right of the main window. It is titled “ChemDraw-LiveLink.”
3. In the text bar, type “Hydrogen,” then hit “enter.” A molecule should appear.
4. Optimize the structure by hitting “control-m.”
5. There are several things you can do to get a better look at the molecule.
   • You can click the third button from the left on the top toolbar to rotate the molecule. It looks like a circle with an arrow on it. After you click it, you can use the mouse to rotate the molecule.
   • You can click “View” on the main menu, then click “Model Display.” This will present you with many options to change the display of the molecule. For example, “Display Mode” gives you more modes. The “Ball & Stick” mode is most common, but “Wire Frame” is convenient for a very complicated molecule, and “Space Filling” is helpful for visualizing atom size differences.
6. After making the model and optimizing the structure, click “Surfaces” on the main menu, then “Choose Surface,” then “Molecular Orbital.”
7. The HOMO (highest occupied molecular orbital) will automatically be shown, but you can choose another molecular orbital from “Select Molecular Orbital.” You will also see the energies associated with each orbital. The labels are with respect to the HOMO and the LUMO (lowest unoccupied molecular orbital). Notice that the HOMO energy is usually negative, indicating a favorable state, but the LUMO energy is usually positive, indicating an unfavorable state (which is why it is UNOCCUPIED!)
8. Draw on your worksheet the HOMO and the LUMO. Indicate color differences either by using color in your drawing or by shading.

Part 2: Luminol

Experimental component

Source: https://www.carolina.com/teacher-resources/Interactive/luminol-glowing-reaction/tr10786.tr

Note: perform this section with a partner.

Create Solution A

1. Measure out the correct mass of NaOH pellets according to your pre-lab calculation. Place into a 250 or 300 mL Erlenmeyer flask.
2. Measure out 100 mL of DI water in a 150 mL beaker. Add to the Erlenmeyer flask with the NaOH pellets. Swirl to dissolve.
3. Measure out ~0.18 g of luminol in a weigh boat. Add to the Erlenmeyer flask. Swirl to dissolve. Pro-tip: an Erlenmeyer flask is a great choice when mixing up a solution from a solid because you can swirl without the solution splashing out of the flask. Always use an Erlenmeyer flask that has plenty of extra space to allow you to swirl.
4. Record the appearance of Solution A.

Create Solution B

1. Measure out 0.03 g of potassium ferricyanide.
2. In a second 250  or 300 mL Erlenmeyer flask, add 100 mL of DI water, 1 mL of 3% hydrogen peroxide, and the 0.03 g of potassium ferricyanide. Swirl to dissolve.

Mix them together

1. Go into the dark room.
2. Slowly pour the two solutions simultaneously into the funnel at the top of the apparatus.
3. Record your observations immediately after the two solutions mix.
4. Dispose of the solution in the waste containers.

Computational component

1. Open Chem3D 18.0
2. Click on the white panel to the right of the main window. It is titled “ChemDraw-LiveLink.”
3. In the text bar, type “3-aminophthalate,” then hit “enter.” A molecule should appear.
4. Optimize the structure by hitting “control-m.”
5. After making the model, click “Surfaces” on the main menu, then “Choose Surface,” then “Molecular Orbital.”
6. The HOMO (highest occupied molecular orbital) will automatically be shown, but you can choose another molecular orbital from “Select Molecular Orbital.” You will also see the energies associated with each orbital. The labels are with respect to the HOMO and the LUMO (lowest unoccupied molecular orbital).
7. You will need to record the difference between the MOs prescribed on your worksheet. Note that ChemDraw reports the band gap in terms of eV, so you will need to convert from those energy units to wavelength units to determine the color that corresponds to each of these potential transitions. This will require using Plank’s constant in different units than you typically use: h = 4.1357 × 10^–15 eV s. I suggest that you use Excel or some other software to do this repeated calculation. Once you know the colors of each possible transition, you will be able to choose the most likely transition that corresponds to the color you observed.

Part 3: Flame tests

Note: perform this section individually in a fume hood.

1. Observe sources one at a time by eye. If you are color-blind, you will need to work with a non-color-blind partner.
2. Light a bunsen burner with a non-luminous blue flame using your amber tubing and your bunsen burner in  a fume hood.
3. For the first metal, obtain a wooden splint that has been soaked in the metal solution.
4. Wave the wet part of the splint in the hottest part of the flame (the top of the inner cone).
5. The color you observe is the first color you see, not the yellow color of the wood burning. If you aren’t sure what color is wood burning, experiment by burning a dry splint.
6. Write down your observations of the color(s) and intensity on your worksheet.

Report

Fill out this worksheet. Turn in either a paper or digital copy.

Learning goals: maintain safety in a chemistry laboratory,  follow instructions to complete a laboratory experiment, collect experimental data, explain likely sources of experimental error

Introduction

A chemical reaction occurs when electrons and/or bonds rearrange. A chemical reaction can be described with a chemical equation. You may notice a chemical reaction has occurred by one of the following indicators:

• A color change (eg. a colorless solution turns red),
• A change of state (eg. a solid precipitates in a solution),
• A change in temperature (eg. a solution gets hot).

Today’s lab activity is an example of a chemical reaction. You will form a golden solid from two colorless solids.

Background

Lead nitrate and potassium iodide are both highly soluble in water. This means:
Pb(NO₃)_{2 (s)} → Pb²⁺_(aq) + 2NO₃^–_(aq)
KI_(s) → K⁺_(aq) + I^–_(aq)

When you mix those two solutions together, all four ions encounter each other, and that’s when the chemical reaction occurs:

Pb²⁺_(aq) + 2I^–_(aq) → PbI_2(s)

The potassium and nitrate ions are not involved in the chemical reaction; they are called “spectator ions.” They are present in the same amounts before and after the reaction (on both sides of the arrow).

Pb²⁺_(aq) + 2NO₃^–_(aq) + 2K⁺_(aq) + 2I^–_(aq) → PbI_2(s) + 2NO₃^–_(aq) + 2K⁺_(aq)

When the solid forms, it first looks like yellow clouds because the tiny solid particles are suspended in the solution (like mud). We have to take advantage of its slight solubility in order to form a sparkling, golden solid. When you first add the lead nitrate to the potassium iodide, you will notice the cloudy precipitate forms and then dissolves. This is because lead nitrate is “slightly soluble.” 0.08 g of it can dissolve in 100 mL of water: the yellow clouds only remain visible when there is more than 80 mg present per 100 mL of water.

Temperature dependence of solubility for several solids. Source: http://wps.prenhall.com

Many solids are more soluble in hot water than cold. Heating the solution will cause the lead nitrate to dissolve completely so that it can slowly crystallize as it cools, which forms large crystals.

Procedure

Note: do this lab individually.

Prepare solutions

1. Tare a 250 or 300 mL Erlenmeyer flask labeled “lead nitrate”. Add approximately 0.3 g of lead nitrate to the flask. Record the exact mass on the worksheet.
2. Measure out 100 mL of deionized water (DI H2O) with a graduated cylinder. Pour the water into the flask containing the solid lead nitrate. Swirl to dissolve.
3. Repeat this procedure with another flask for the Potassium Iodide.
4. Add a few drops of 0.5 M hydrochloric acid to each solution.

Form Crystalline Lead Iodide

1. Add the lead nitrate to the potassium iodide one drop at a time, swirling to dissolve after each drop. How many drops can dissolve? Once the yellow clouds start to form and will not dissolve when swirling, pour in the remainder of the lead nitrate. Record your observations.
2. Prepare a hot water bath with a Bunsen burner, a ring stand, a ring, a wire mesh, and an 800 mL or 1000 mL beaker – similar to the picture to the right. Alternatively, you can use a tripod instead of the ring stand and ring. Put approximately 500 mL of water in the beaker, then heat the water up to boiling. (Hint: If you read through these instructions before coming to lab, you will probably start with this step so you do not have to wait for your water to heat up.)
3. Carefully place the Erlenmeyer flask containing the lead iodide in the hot water bath. Heat the lead iodide solution up to about 80^oC or until the lead iodide dissolves. Record your observations.
4. Once the lead iodide has completely dissolved, remove the Erlenmeyer flask from the hot water bath. Set it on the bench top and watch the crystals form. Wait about half an hour for the crystals to precipitate slowly at room temperature. (The longer you wait, the nicer your crystals will look.) Record your observations. This may take a few minutes, so while you’re waiting, you should work on cleaning up the hot water bath, completing the worksheet, inventorying your glassware, and preparing for the next steps:
   • Prepare an ice bath. With the same 800 mL beaker, fill it about halfway with ice and water.
   • Obtain a Büchner funnel, a filter flask, and an appropriately sized piece of filter paper. Record the mass of the filter paper.
   • Set up a filter system like the one shown to the left. Wet the filter paper with a few mLs of water so that it sticks to the bottom of the funnel.
5. After you have waited about half an hour, put the flask into the ice bath for about ten minutes.
6. After that, turn on the vacuum source and pour the crystals and solution into the Büchner funnel. Use a few mLs of ice water to rinse away any crystals that may have stuck to the sides of the flask.
7. Turn off the vacuum source once the crystals and paper appear dry.
8. Weigh the filter paper and crystals on a tared watch glass. Record the mass on your worksheet and calculate the final mass of crystals.
9. Ask your instructor or TA to check your data and observations before you clean up or leave.

Clean up

Pour the filtrate into the waste beaker labeled “lead iodide waste liquid.” Take a picture of your pretty crystals to show your roommate, then put the crystals and filter paper into the waste beaker labeled “solid lead iodide waste.” Pour hot water and ice baths down the sink. Rinse all glassware and return to them to where you found them.

Report

Complete the worksheet. Turn in either a paper or digital copy. This article will help you with the conclusion question. Optional survey.

_{by M. J. Simpson}

Note: Plan on dressing in old clothes for this lab. These dyes are powerful and may permanently stain your clothes if you splash or drip them on yourself.

Tie dye is a classic color chemistry lab. There are many different “recipes” for tie dye, but we will be following the procedure provided by Colorado Wholesale Dye Corp to ensure the most vibrant colors possible.

The wash and dyes will be prepared for you, so these are the instructions you need to know:

Before coming to lab: Prewash your 100% cotton, white garment before coming to lab.

In lab:

1. Use a permanent marker to write your name on the tag of your garment.
2. Soak your garment in the Sodium Carbonate wash for at least 10 minutes. This solution is very basic, so use gloves when in contact with the wash. Wring out the garment to dry it as much as you can.
3. Fold and tie with waxed sinew your garment however you like.  Here are some designs you might like to try. This video shows you some techniques:

    

4. Working on a surface covered in newspaper, apply the dye. Be sure to saturate the fabric with dye, but don’t put so much on that the dye pools underneath the fabric. Again, wear gloves when working with dye or it will dye your skin!
5. Wrap your garment in newspaper and put it in a sealed plastic bag. Write your name and section on the bag, then give it to the instructor.

_{adapted by M. J. Simpson from a lab originally by S.F. Sontum, S. Walstrum, and K. Jewett}

Introduction

Theories of chemical bonding allow us to understand the electronic structure and geometrical arrangement of atoms in a molecule or ion. Models provide a useful way of visualizing the arrangement of electrons in a molecule. As you learned in class, there are several types of structure representations used by the chemist at different times to explain chemical phenomena. These include Lewis structures (including formal charges) based on a simple counting rule; valence bond models (including hybridization, and resonance) based on orbital overlap; valence shell electron pair repulsion theory (VESPR theory) based on electron repulsion, to predict overall shapes of molecules/ions; and molecular orbits to predict certain electrical and magnetic properties.

The figure above is a ball and stick model for an early precursor to one of the anti-HIV drugs known as Protease inhibitors developed at Abbott Laboratories, Inc. We need models like this because actual molecules are too small to see. Modeling of molecular structures allows us to make predictions about the behavior of these invisible molecules so that we can design new chemicals with beneficial properties such as these inhibitors. The ball and stick model kits you use in this chemistry laboratory are very useful for visualizing chemical structures but they do not give quantitative information needed to design and build new molecules with specific properties.  ChemDraw is a computer “modeling kit” which we use to augment our visualization of the molecule with quantitative information.

Procedure

To illustrate how molecular geometry can be obtained from Lewis structures and valence shell electron pair repulsion, we will use molecular models. With models, it is relatively easy to see both geometry and polarity, as well as to deduce Lewis structures. You may want to initially generate your Lewis structures before you come to the laboratory. When you come to the laboratory use the molecular models to check and refine your Lewis structures. In this exercise you will assemble models for a number of common chemicals and interpret them in the ways we have discussed.

The models consist of plastic bonding centers and bonding tubes. The bonding centers represent the hybrid valence electron distribution about the atom. The tubes represent the bonding electron pairs or valence bonds. As you build the models also draw the three dimensional structures on paper, so that you can develop skills at representing these three dimensional structures on paper.

Most of the structures will obey the octet rule. Often the central atom will have four electron pairs and will have a tetrahedral bonding center. In your model kits, multiple bonds are modeled by the long, flexible bonds.  We will see from our molecular orbital studies that the bent “banana” p bonds of our model kit are not good representations of the electron distribution but never the less correctly represent the Lewis structure.

Methane: Construct a model of methane, CH₄.

Place the model on the desk top and note the symmetry of this tetrahedral molecule. The molecule looks the same regardless of which three hydrogens are resting on the desk. All four hydrogens are said to be equivalent. Is this molecule polar? Draw a three dimensional picture of this molecule.

Dichloromethane: Replace two of the hydrogen atoms in methane with a chlorine to make dichloromethane CH₂Cl₂. Is dichloromethane polar? Which Lewis structure best represents dichloromethane or are they the same?

Ethyl Alcohol and Dimethyl Ether: Construct ethyl alcohol, CH₃CH₂OH, and dimethyl ether, CH₃OCH₃.  The following is a chart of the physical properties of these two isomers.

Which isomer is more soluble in water? Which would have a higher boiling point? Why?

Constructing Lewis Structures

Now that you have some experience building molecular models from a Lewis structure you should be ready to fill in a geometrical structures chart starting only from the molecular formula.  In assembling a molecular model of the kind we are considering, it is possible, indeed desirable, to proceed in a systematic manner. We will illustrate the recommended procedure by developing a model for the first molecule on the chart CH₂O. Please work along with your model as we describe the procedure.

1. Choose a skeleton with the least electronegative atom for the central atom.
   For CH₂O, this is carbon. Remember, ring structures may be possible with six member rings being the most stable and 3 member rings the least stable.
2. Determine the total number of valence electrons.
   For CH₂O this total would be 12 valence electrons. Now collect 6 bonding links representing the 6 pairs of bonding electrons.
3. Assemble a single bonded skeleton structure.
   In CH₂O the model can be assembled by using two links to form s bonds between C and H, and then using a third link to form a s bond between C and O.
4. Add lone pairs.
   Use the remaining tubes to fill all the unfilled sites on the bonding center and to satisfy the octet rule. In the model, three links may be used to fill the O site leaving one unfilled sites on C.
5. Form multiple bonds as needed to satisfy the octet rule.
   Move one of the lone pair electrons off of the oxygen to generate a double bond with carbon and satisfy the octet rule on carbon.
6. Interpretation of the Lewis Structure.
   The tubes and spatial arrangement of the bonding centers will closely correspond to the electronic and atomic arrangement in the molecule. Be sure to check that you have a valid Lewis structure (octet rule) and the right total number of electrons. The Lewis structure of the molecule is given below:

Formal Charge: Since all of the octet atoms satisfy the 8n rule there is no formal charge.

Hybridization: The angle between VSEPR groups is 120° so the hybridization is sp² on both the carbon and the oxygen.

Resonance structures: The formation of other resonance structures would entail the moving of p bonds or the formation of new p bonds from lone pairs. For this molecule, there are no other ways to rearrange the electrons over the molecular skeleton and still satisfy the octet rule. This is the only possible full octet resonance structure for this molecular skeleton.

Atom Geometry: Given our model, we would describe the CH₂O molecule as being planar with single bonds between carbon and hydrogen atoms and a double bond between the C and O atoms.

Polarity: Since all bonds are polar and the molecular symmetry does not cancel the polarity in CH₂O, the molecule is polar.  (The compound having molecules with the formulas CH₂O is well known and is called formaldehyde. The bonding and structure of CH₂O are correctly given by this model.)

Isomerism: A possible alternate skeleton and Lewis structure would be H–C=O–H but this structure results in a formally positive oxygen next to a negative carbon. Because of the unfavorable charge separation this molecule does not exist.

Introduction to ChemDraw

ChemDraw is another “modeling kit” which we use to augment our information about molecules, visualize the molecule, and give us quantitative information about the structure of molecules.

The following is a brief introduction on how to use ChemDraw.

1. Open Chem3DPro
2. Click on the white panel on the right labeled ChemBioDraw – LiveLink
3. Make sure the Text tool is selected when the Tools menu pops up
4. Type the formula. For example, SO₃^2- is entered as SO3-2
5. If you want a different isomer, you can enter the formula more specifically. For example, instead of C3H4, you might enter H2CCCH2
6. Hit enter. The molecule will appear. If it is outlined in red, that means something is wrong with your molecule. Consult the lab instructor or TA if you cannot figure out the problem
7. Hit control-M to minimize energy. The total energy will be displayed at the bottom under Output
8. To measure bond lengths, hover the mouse over the bond
9. To explore the structure in different modes, go to View → Model Display → Display Mode, and then pick another mode. Try a few different ones
10. To show or hide lone pairs, go to Structure → Lone Pairs, and make your selection
11. To save a structure as an image, go to File → Save As, and then change the file format to whatever you like (such as .png or .jpg)
12. Hit control-N to start over
13. Another way to build a molecule is to build a backbone with the solid bond building tool and then change the atom types manually. After building the backbone, just click on the intersection you want to change and type the symbol of the atom you want to change. Double click the bond to make a double bond. This is especially useful for building the P₄ isomers

Report:  Fill out the worksheet for the report.

_{by S. F. Sontum and K. Jewett (edited by R. Sandwick)}

Introduction

The oxides of carbon could not be more chemically different. Carbon monoxide (CO) is a polar basic gas that binds strongly to metals while carbon dioxide (CO₂) is a non-polar gas that reacts with water. Carbon monoxide is produced when hydrocarbons are burned in a limited amount of oxygen while carbon dioxide is produced when hydrocarbons are burned in excess oxygen. Carbon monoxide is a deadly poison while carbon dioxide is essential to the metabolism of plants and animals.

Carbon dioxide is the fourth most common gas in our atmosphere and its concentration is increasing. Atmospheric carbon dioxide plays an important role in maintaining surface temperatures. It is one of the “greenhouse” gases responsible for global warming. The biological processes of metabolism, which produces carbon dioxide, and photosynthesis, which consumes carbon dioxide have functioned to maintain the levels of carbon dioxide in our atmosphere. The industrial revolution has shifted this balance. The concentration of carbon dioxide has increased more than 30% to the present day value of 400 ppm. Global temperatures are rising.

This experiment focuses on a physical property of CO₂ – its energy content, and a chemical property of CO₂ –  its ability to dissolve in water to make the solution acidic. Like all molecules CO₂ stores potential energy in the molecular bonds that hold the oxygen to the carbon and in the intermolecular forces between molecules. We will measure the heat of sublimation of dry ice to investigate the intermolecular forces between CO₂ molecules.

The phase diagrams below indicate the stable forms of compounds at different pressures P and temperatures T. The lines describe when two phases are in equilibrium and the triple point is the place where all three phases can coexist in equilibrium. For example, the triple point of water occurs at 0.01 ^oC and 0.006 atm. In order to convert ice directly into steam the partial pressures of water would have to be lower than 0.006 atm. Because we live at higher pressures (i.e., at 1 atm), we always see ice melt before it vaporizes to steam. The process of converting a solid to its liquid is called melting (aka, fusion). The triple point of carbon dioxide on the other hand occurs at a much higher pressure and lower temperature, -57 ^oC and 5.11 atm.  We call solid carbon dioxide “dry ice” because at normal atmospheric pressures it converts directly from the solid into the gas. The process of converting a solid into a gas is called sublimation. Both melting and sublimation require the input of heat.

Phase diagrams

Calorimetry

When a solid substance is heated, it absorbs thermal energy and its temperature increases.  When the melting point of the substance is reached, addition of further thermal energy breaks up the forces holding the solid together, and a liquid begins to form. A balance is established between solid and liquid, and as more thermal energy is added the temperature remains constant while the amount of solid decreases and the amount of liquid increases. The heat absorbed when one mole of solid is melted at constant pressure is called the molar enthalpy of fusion (Δ H_fus). When heat is absorbed, the sign of Δ H_fus is, by convention, positive. Reactions that absorb heat are said to be endothermic reactions. Those which produce heat are said to be exothermic.

On the molecular scale, many processes occur as the ice absorbs heat: the ordered array of molecules in the crystal lattice of the solid is broken down into a collection of mobile, disordered liquid-phase molecules. The water molecules in ice are also losing potential energy associated with intermolecular forces between molecules and gaining kinetic energy of motion.  Meanwhile, the temperature of the water, associated with the average kinetic energy of motion, rises as the molecules move about and vibrate more rapidly.

The change in internal energy for a reaction Δ E_rxn can be measured by running the reaction in a constant volume bomb calorimeter. By designing the calorimeter so that no heat leaks out to the surroundings, the heat absorbed by the reaction should equal the heat lost by the calorimeter. The heat absorbed by the calorimeter can be measured by the heat capacity of the calorimeter (C_calorimeter) times the change in temperature.

Δ E_rxn  = – C_calorimeter Δ T

The melting of ice or the sublimation of carbon dioxide is not a constant volume process but rather a constant pressure process. To investigate the heat flows in a constant pressure process we will have to define a new form of internal energy called enthalpy, H. At constant pressure the change in enthalpy is related to the change in internal energy minus the work done on the system due to volume changes (Δ V).

ΔH  = ΔE  + Δ(PV) = ΔE  + PΔV

At constant pressure, enthalpy changes are a direct measure of the heat absorbed by the system. Enthalpy changes are easier to measure than internal energy changes because it is easier to maintain the constant pressure than constant volume.

Diagram of a constant pressure calorimeter

Δ H is measured with a calorimeter where the amount of heat flowing is reflected in a temperature change of a known mass of water. In fact, the unit of energy, the calorie, has been defined to be that amount of heat that will raise one gram of pure water one degree Celsius in temperature. Although the calorie is convenient energy unit, the scientific community now uses the Joule as its standard unit of energy. (1 cal =  4.184 J) Assuming no heat loss during the transfer, the heat of a reaction will be equivalent to the heat absorbed by g grams of water when the temperature raises Δ T degrees Celsius, or

ΔH_rxn = – g H₂O (4.184 J/g H₂O ^oC ) ΔT

This is the basic equation describing a solution calorimeter that is intended to measure the change in enthalpy during a constant pressure process.

Enthalpy of Reaction (see Hess’s law in the textbook)

Sometimes it is difficult or even impossible to measure an enthalpy of formation directly.  In this experiment we will be determining the enthalpy of formation for CO₃^2-_(aq).

C_(s) + 3/2 O_2(g) + 2 e^– →  CO₃^2-_(aq)

Direct measurement by burning C(s) in oxygen would not give us carbonate, so we will be using an indirect method. According to Hess’s Law (the conservation of enthalpy), if two or more reactions can be added to give a net reaction, Δ H for the net reaction is simply the sum of the Δ H’s for the reactions which are added (energy is additive). Consider the following four reactions:

(1) 2H+(aq) + 2e- → H₂(g)                             Δ H1

(2) CO₃^2-(aq) → C(s) +3/2 O₂(g) + 2e-       Δ H2

(3) H₂(g) + 1/2 O2(g) → H₂O(l)                    Δ H3

(4) C(s) + O₂(g) → CO₂(aq)                           Δ H4

(5) 2H+(aq) + CO₃^2-(aq) → H₂O(l) + CO₂(aq)

Δ H5 = Δ H1 + Δ H2 + Δ H3 + Δ H4

You will measure Δ H₅ directly. We will combine this with the literature values for the heats of formation of H⁺_(aq), H₂O_(l), and CO_2(aq).

H2(g) + 1/2 O_2(g)(g)  → H₂O(l)

Δ H3 =   Δ H_f^o (H₂O(l))= -285.83 kJ/mole

C(s) + O_2(g)(g)  → CO₂(aq)

Δ H4 = Δ H_f^o (CO₂(aq))= -413.80 kJ/mole

H₂(g) → 2 H+(aq) + 2 e-

-Δ H1 = 2 x(Δ H_f^o (H+(aq))=  2 x 0 kJ/mole

You will thus be able to calculate Δ H₂, which is the enthalpy of formation of CO₃^2-_(aq) reaction written backwards.

Bomb Calorimetry

Not like this bomb

As we said above, the daily caloric intake of an adult is about 3000 kcal/day. To maintain an energy balance an average adult consumes six tons of solid food over a 40 year period, which amounts to about 12.5 g/hr. Typically we obtain one third of our energy from each primary type of food substance — carbohydrates, proteins and fat. The average enthalpy content of carbohydrates and proteins are about the same (4 kcal/g) while fats average twice as much (9 kcal/g). We will measure the caloric content of food directly by using a bomb calorimeter. We will run a one-gram sample of food in the bomb calorimeter, determine its heat of combustion, and compare the value with what the package says is its caloric content.

Experimental Procedures

Part 1   Enthalpy of Sublimation

Clean, dry, and weigh the calorimeter (2 styrofoam cups, one inside the other – the inner cup will have a series of holes, plus the cover with one hole taped closed).

Fill a 100 mL volumetric flask to the mark with distilled water adjusted to room temperature and weigh it. Place a split-holed rubber stopper carefully over the top of the high precision thermometer.  Carefully insert the thermometer into the volumetric flask and stabilize the thermometer by attaching the split-holed stopper to a ring stand via a Bunsen burner clamp. Record the stable temperature; this is the starting temperature of the water.

Once you have weighed the calorimeter and measured the stable temperature reading, take a piece of CO₂ (between 2 and 10 grams) with tweezers and put it between the two cups and weigh the entire assembly. The mass will keep rolling down as it sublimes, therefore after a few seconds record the mass, not bothering with the last digit, immediately take the assembly back to your station where your partner has the thermometer ready to put through the cover. Pour the 100 mL of water into the cups, replace the cover, and gently push down, trying not to spill too much water as it seeps into the lower cup. Reweigh volumetric flask to determine the mass of water transferred.

Even though you see your system “smoking”, the “smoke” shouldn’t be cold to the touch as the CO₂ gas has to absorb the heat of the water as it travels through it. Swirl the cups gently, again making sure to hold the two cups together as much as possible and to keep the cover on. Watch the temperature drop. When it stabilizes and you don’t feel the bubbles in solution, swirl again and look in the calorimeter. If it’s cloudy, return to swirling. If it’s clear, record the stabilized temperature.

Empty your calorimeter and thoroughly dry the cups, cover, and thermometer.  Repeat the experiment.  Calculate enthalpy.

Calculations:  The heat gained by the CO₂ (Δ H_sub) is equal to the heat released by the water.  Knowing that the specific heat of water is 4.184 J g^-1 C^o, calculate the average ΔH_sub/g CO₂ and the average Δ H_sub/mol CO₂.  Compare versus the literature value.

Part 2   Enthalpy of Reaction for formation of H₂O and CO₂ from CO₃^2- and H⁺

The calorimeter and thermometer setup will be two small cups, the same size, with no holes.  Rinse the inner cup of your calorimeter with distilled water and dry the cup, cover, thermometer and stirrer thoroughly.

1. Obtain 100 mL each of the 1.00 M solution of K₂CO₃ and the 3 M HCl in clean, dry, labeled beakers.
2. Add exactly 50.0 mL of your solutions to each of two 50 mL volumetric flasks. Weigh both. Load the calorimeter with the 50.0 mL of 3 M HCl , then reweigh the volumetric flask. Insert the thermometer into the 50 mL of K₂CO₃ solution and determine the initial temperature. Rinse the thermometer and assemble the calorimeter (thermometer and cover).
3. Begin measuring the temperature of the HCl solution and record the temperature every 15 seconds. After the temperature stabilizes (three consistent measurements), add the K₂CO₃ solution at a moderate pace so that it does not bubble out (but within 10 seconds), then re-cover the calorimeter. Swirl the solution to release all CO₂ and continue recording the temperature until it is approximately constant for several readings (at the end you may see a slight decrease due to heat leakage). Reweigh the 50.0 mL K₂CO₃ volumetric flask to determine how much was transferred.

Carefully empty the calorimeter and clean the cup, cover, and thermometer. Rinse the cup with distilled water. Dry all pieces thoroughly. Repeat the experiment once a second time. Record your data in a table.

Calculations:  Assume the mixture of the two solutions has the same specific heat as water (4.184 J g^-1 C^o) and perform a calculation similar to Part 1 to determine Δ H_rxn and then Δ H_rxn/ mol CO₃^2-.  (This is Δ H₅.) Use Hess’s law to determine Δ H_f for CO₃^2- (Δ H₂ backwards). Compare ΔH₂ to literature values.

Part 3   Bomb Calorimetry

Perform bomb calorimetry on a sample at some point during the lab period. Each pair of students should bring and be prepared to analyze a uniform dry (or chocolate) food sample using the bomb calorimeter. Weigh out one gram +/- .0005 g. Crush (if it’s dry) or break the sample into small pieces (if it’s chocolate) and fill the calorimeter pan. Procedures for using the Parr bomb calorimeter will be demonstrated in class.

Calculations:  The read-out from the calorimeter will give you cal/grams. Remembering that a food calorie (Cal) is actually a kcal, compare your results to the value listed on the package.

Report

Complete the worksheet for your lab report. You should have all calculations clearly shown and your final results compared to literature values.

_{adapted by M. J Simpson from a lab originally by S. F. Sontum, K. Jewett and R. Sandwick}

Learning goals: work collaboratively with a lab partner, follow instructions to complete a chemistry experiment, collect experimental data, formulate logical conclusions based on experimental results

Introduction

Much as a cook uses recipes to guide her/him in preparing dishes, chemists use chemical equations as ways of expressing the results of chemical reactions.  Sometimes it is more descriptive to use the ions involved in a reaction rather than the ionic compounds. For example, to express that hydrochloric acid reacts with sodium hydroxide to give water and sodium chloride, we could write:

Molecular equation: HCl_(aq) + NaOH_(aq) → H₂O_(l) + NaCl_(aq)

The above equation is fine in some ways, but in actuality NaOH and HCl solutions do not contain NaOH and HCl molecules, and there are no NaCl molecules in the solution after the reaction occurs. The better description for this reaction is:

Total ionic equation: H⁺_(aq) + Cl^–_(aq) + Na⁺_(aq) + OH^–_(aq)  →  Cl^–_(aq) + Na⁺_(aq) + H₂O_(l)

Net ionic equation: H+_(aq) + OH-_(aq) → H₂O_(l)

The first equation is called the total ionic equation, while the second is termed the net ionic equation.  A net ionic equation summarizes the changes that have taken place as a result of a chemical reaction.

Background

The solubility of a substance in a solvent is the maximum amount of the substance that can be dissolved in a given amount of solvent. While there is no exact definition for the boundary between soluble and not soluble, a general guide might be:

SOLUBILITY                                      TERM

> 0.1 M                                                 Soluble

0.01 to 0.1 M                                       Moderately soluble

< 0.01 M                                              Slightly soluble

Some guidelines have been prepared to help estimate a compound’s solubility. These are written in a manner where statements above other statements have priority. For example, an alkali metal carbonate is soluble since “all alkali metals (#2) are soluble” have priority over “All carbonates are insoluble.” [There are exceptions that are not included here. (See textbook)]    

Solubility Rules:

1. All nitrates (NO₃-) and acetates (C₂H₃O₂-) are soluble.
2. All salts of alkali metals and NH₄+ are soluble.
3. All common halides (Cl-, Br^–, I^–) are soluble, except those of Ag⁺, Pb²⁺, Cu⁺, Hg₂²⁺
4. Hydroxides (OH-) are insoluble except for Na⁺, Ca²⁺, and Ba²⁺.
5. All sulfates (SO₄^2-) are soluble, except BaSO₄ which is slightly soluble, and CaSO₄ and Ag₂SO₄, which are moderately soluble.
6. All carbonates (CO₃^2-) and phosphates (PO₄^3-) are insoluble.

Procedure

Note: perform this experiment with a partner, but be sure to observe the tests together.

Make up the following standard solutions:

50.0 mL of 0.10 M KI_(aq) from KI_(s)

25.0 mL of 0.10 M CuCl_{2 (aq)} from CuCl₂ ● 2H₂O _(s)

50.0 mL of 0.10  M H₂SO_{4 (aq)} from 1.00 M H₂SO_4(aq)

25.0 mL of 0.10 M NaOH _(aq) from 1.00 M  NaOH_(aq)

50.0 mL of 0.10 M Na₂CO_{3 (aq)} from Na₂CO_3(s)

Three test solutions will be prepared for you:  0.10 M AgNO₃, 0.10 M BaCl₂, and 0.10 M Co(C₂H₃O₂) ₂.

To determine how to prepare a solution from a solution of higher concentration, use the formula: C_i V_i = C_f V_f

where:

C_i = initial concentration of higher concentration solution

V_i = initial volume of the higher concentration solution

C_f = final concentration (of the diluted solution)

V_f = final volume of the diluted solution (volume of the initial solution plus water)

Tests for chemical reactions:

For each of the solutions you prepared, test each solution with the following test solutions:

0.10 M AgNO₃

0.10 M BaCl₂

0.10 M Co(C₂H₃O₂) ₂

0.10 M Na₂CO₃ (the solution you prepared)

A test involves taking 5 mL of your solution and adding 3 – 5 drops of the test solution. Make observations in the table on the worksheet.

Things to note:

(a)  a change of color;

(b)  a separate phase is formed;

(c)  any other change.

Report

Complete the worksheet for your lab report. Optional survey.

The following procedure may help you when writing a net reaction for any chemical reaction.

1. Determine the principal forms present of each reactant in solution.
2. Determine the principal forms present after the reaction has occurred.
3. Write a balanced equation for the reaction.
4. Cross out reactants and products that do not change (the spectator ions).

Example:  Mix HNO₃ solution + KOH solution

Principal substances before reaction:  H+, NO₃-, K+, OH-

Principal substances after reaction: NO₃-, K+

Product formed:  H₂O

Net reaction:  H+_(aq) + OH-_(aq) → H₂O _(l)