K-12 Teaching and Learning From the UNC School of Education

methanol structure

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Learning outcomes

The goal of this activity is twofold:

  • Computational (chemistry education): Use interactive computing (specifically computational chemistry) to introduce and/or reinforce basic chemical principles including Lewis dot structures, bonding, atomic and molecular orbitals.
  • (Computational chemistry) education: Help students understand how COMPUTATION is used in chemistry education and chemistry research. In this case, students use a high-performance computational chemistry server, running state-of-the-art chemistry software. This server is located in Durham, NC, and is accessible 24/7 to pre-college students who reside in the State of North Carolina.

With this activity, students can calculate and visualize the atomic and molecular structures of bonds and lone pairs in the molecule methanol (methyl alcohol, CH3OH).

Teacher planning

Time required for lesson

Approximately 45 minutes; lab time approximately 30 minutes

Materials/resources

If this activity is used as a demo, the teacher needs an account on the North Carolina High School Computational Chemistry Server. Accounts are free of charge, and instructions for obtaining an account are available at this site. If this activity is a student lab, students must have account on the chemistry server.

Technology resources

  • Access to Java-enabled web browser. Users are encouraged to use Mozilla Firefox, which is downloadable free of charge.
  • An LCD projector connected to an internet-accessible computer is required for demo purposes.

Pre-activities

Activity background

The molecule methanol (methyl alcohol) has the structure CH3OH, and contains fourteen valence electrons (four for carbon, six for oxygen, one each for the four hydrogens). Students should be able to construct a Lewis dot structure, making sure that the carbon and oxygen atoms have enough electrons to satisfy the octet rule (eight surrounding electrons) and that the hydrogens have two electrons. Given that ten of the electrons form bonds, it should be the case that there are two lone pairs (LPs) resident on the oxygen. The Lewis diagram and the bonded structure are shown in the Methanol Lewis dot structure graphic.

Activities

In this quick lab, students will build the methanol molecule and run a basic computational chemistry calculation (using the Hartree-Fock method with 3-21G mathematics, using the Gaussian software) to calculate where the electrons are located. The specific calculation is known as a “natural bond order,” or NBO, determination.

Method

  1. Once logged in to the computational chemistry server, click on New Job → Create New Job. You will now see a Build Molecule window. Click on Open Editor.
  2. Once the editor window is open, click ONCE to deposit a carbon (gray) atom. If you deposit too many, click on Tools→Adjust. Click on the extra atom(s), hit the delete key. Return to building using Tools→Build.
  3. Under the Build menu, click on O to select an oxygen (red) atom. Deposit this atom in the Build window.
  4. Draw a bond between the carbon and oxygen by clicking on one of the atoms, draw a line to the other atom.
  5. Under the Cleanup menu, click on Comprehensive. This should add the remaining hydrogens and clean up the structure so that the bond lengths and angles are within reason. Place the cleaned molecule in the Build molecule window by clicking in that window. Alternatively, click on File→Close in the editor menu bar.
  6. Go to the next window by clicking on the arrow at the bottom right of the window.
  7. In the Choose Computational Engine window, click on the Gaussian option. Gaussian is a commercial computational chemistry software package, and is considered by most to be the “industry standard” in the chemical research community.
  8. You can name the job if desired, something like “Methanol NBO calculation.” Alternatively, you can leave it named CH4O.
  9. For this lab, you will use a specific mathematical theory known as “Hartree-Fock,” supported by a mathematical description (basis set) of the molecule known as “3-21G.” This is a standard approach to molecular calculations.
  10. For the Calculation type, use the pull-down menu and choose Natural Bond Orbitals.
  11. Charge should stay at 0 (neutral molecule) with multiplicity (a description of the spin condition of the electrons) of singlet.
  12. Hit the next arrow in the lower right-hand corner to start the calculation.
  13. The calculation should take about twelve seconds to complete. To follow the progress, click on the Refresh button located to the right of the New Job menu item. If you are doing this experiment with a class of students, it will take a little while for all of their calculations to complete.
  14. Once the job is completed, click on either the name of the job or the magnify tool to see the results.

Analysis of Results

  1. Scroll down to Natural Bond Orbitals.
  2. Observe that there are five “bonding (BD)” orbitals, corresponding to the appropriate bonds. Observe also that there are almost (but not quite) two electrons in each of these orbitals.
  3. Observe that there are two lone pairs located on the oxygen atom, and that each lone pair also contains two electrons. This is confirmation that the Lewis dot structure constructed by the students is indeed correct.
  4. You can visualize these orbitals by clicking on the Magnifying Glass tool under Actions. This should open a new Java window known as “MOViewer.” This application will not work with older web browsers, and sometimes does not work with Microsoft Internet Explorer. Mozilla Firefox typically works just fine.
  5. If you look at NBO Orbital #8, you should see a great deal of electron density around the oxygen atom. Slightly more than half of the orbital is an “s” orbital, with slightly less than half residing as a “p” orbital. Regardless, you should be able to a great deal of electron activity around the oxygen. This is the lone pair.
  6. Likewise, if you look at NBO Orbital #9, you should again see electron activity around the oxygen atom. Unlike Orbital #8, the electrons in this lone pair are 100 percent “p,” and you can see the two lobes (one red, one blue) of the “p” orbital.

Extension

  1. If students have studied hybrid orbitals, you can also look at the Hybrid Orbitals section. This section shows the hybridization contribution of each of the electrons. For example, the first electron that binds Carbon 1 (C1) to the oxygen (O2) is partly an “s” orbital (18. 92 percent) and 81.08 percent a “p” orbital. Visualizing the orbital (shown at right), you can see that this is primarily a “p” orbital. The different colors represent the two nodes of the “p” orbital. Notice that they are not symmetrical: the blue node is bigger, this coming from the additive effect of the “s” orbital. Compare this with the 100 percent “p” orbital seen in NBO Orbital #9 in the previous activity.
  2. Students can complete the chart below if desired. This chart asks students to record the data for each of the hybrid orbitals. Students should observe several things:
  3. The extent to which most of the electrons are hybridized. Only a few (four, six, eight, and ten) are completely “s” orbitals. Notice, also, that one of the lone pairs is completely “p” in configuration.
  4. The fact that each orbital doesn’t really “see” an entire electron. For example, in the first orbital, there is only 69 percent of an electron that tends to be found in that orbital, whereas 131 percent of an electron resides in the second orbital.

Note on other nomenclature:

  • CR: core electrons, that is, non-valence electrons, are all “s” orbital electrons.
  • RY: Rydberg electrons, orbitals into which excited electrons can go. Typically we tell students that these orbitals are empty unless energy is applied, but notice that a very little bit of an electron can be found in these orbitals. Electron distribution through a molecule is not as clean as the textbooks tend to describe it!

Assessment

The teacher should use this activity to assess student’s ability to construct Lewis dot structures, and to assist them in visualizing what these structures look like chemically. The teacher can also use this activity to introduce and/or reinforce concepts such as “s” and “p” orbitals, and hybridization.

Supplemental information

Associated Files

Alternative assessments

Students should have an opportunity to describe, either verbally or in writing, the visualized orbitals generated by these computations.

Critical vocabulary

Important definitions are available via Wikipedia and are linked to this lesson plan.

  • Lewis dot structures (notation)
  • Covalent bonding
  • s and p orbitals
  • Lone pairs
  • Computational chemistry
  • Hartree-Fock method
  • Basis sets
  • Molecular orbitals
  • Valence electrons
  • Core electrons
  • Rydberg electrons

Comments

It is the belief of all of the major scientific organizations, including the National Science Foundation, that students should develop and demonstrate competence and competence in the following areas:

  1. Observational science
  2. Experimentational science
  3. Theoretical science
  4. Computational science

This activity makes use of observation, theory, and computation. Students can, on their own, build and calculate other molecules at school or at home. Students can also request a research account on the NC High School Computational Chemistry server, for use in science fair or other research programs.

  • North Carolina Essential Standards
    • Science (2010)
      • Chemistry

        • Chm.1.2 Understand the bonding that occurs in simple compounds in terms of bond type, strength, and properties. Chm.1.2.1 Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds. Chm.1.2.2 Infer the type of bond and chemical formula...

North Carolina curriculum alignment

Science (2005)

Grade 9–12 — Chemistry

  • Goal 2: The learner will build an understanding of the structure and properties of matter.
    • Objective 2.07: Assess covalent bonding in molecular compounds as related to molecular geometry and chemical and physical properties.
      • Molecular.
      • Macromolecular.
      • Hydrogen bonding and other intermolecular forces (dipole/dipole interaction, dispersion).
      • VSEPR theory.