In this lesson, students use tools to predict the shapes of simple molecules and discuss factors that cause molecules to adopt certain shapes. These concepts are then applied to real-world examples of how geometry impacts the functions of important molecules.
This lesson will help prepare your students to meet the performance expectations in the following standards:
- HS-PS1-3: Plan and conduct an investigation to gather evidence to compare the structure of substances at the bulk scale to infer the strength of electrical forces between particles.
- Scientific and Engineering Practices:
- Developing and Using Models
By the end of this lesson, students should be able to:
- Determine the geometry and bond angles of simple molecules.
- Describe why molecules form in specific shapes.
- Explain how molecular geometry can affect real-world applications of a molecule.
This lesson supports students’ understanding of:
- Molecules and Bonding
- Molecular Geometry
- Molecular Structure
- Lewis Dot Diagrams
Teacher Preparation: 5 minutes
- Projector connected to a computer with internet access
- Balloons (4 per class period, up to 6 if building structures with expanded octets as an extension)
- Some students may have allergies to latex. Latex-free balloons are strongly recommended.
- This lesson was created as a 5E lesson. The notes below outline the 5 stages (Engage, Explore, Explain, Elaborate, Evaluate) and indicate specific slides in the accompanying PowerPoint presentation for teachers to reference.
- Image credits/copyright information are noted below, where relevant, and on each PowerPoint slide in the notes section.
- Engage: (Slide 2) Ask students how light photons hit our eyeballs and get turned into pictures and colors by our brain. Explain to them that there is a molecule in our eyes called “Retinal” that is responsible for gathering light and interpreting it for our brain.
- The thing that allows retinal to collect light is a very specific structure or isomer of the molecule that we call 11-cis-Retinal (structure A below).
- When light or photons of certain wavelengths hit this molecule, depending on how it is bonded to proteins in your eye, the energy from the light will cause it to snap straight, generating all-trans-Retinal (structure B below).
- The transformation is shown here:
Image © RicHard-59, Retinal cis and trans, CC BY-SA 3.0
- When this molecule snaps straight it changes its orientation with proteins in the eye. This activates a nerve impulse that sends a signal to the brain and results in our ability to see.
- Ask the students to find the difference between the two molecules. Encourage students to determine the difference with no additional instruction, and even have them figure out which bond changed/rotated. Students may need some guidance in their thinking. Consider having them count the carbons and hydrogens and, possibly, the number of bonds (students may need assistance understanding where the carbons and hydrogens occur in the line structure portion of the structure).
- (Slide 3) They should identify that both molecules contain the exact same atoms bonded in the same order; all that differs is the orientation of one of the double bonds causing a different 3D structure. This simple process is responsible for our vision!
- Tell the students that in today’s lesson we are going to learn more about how the structure of molecules can change how they behave.
- Explore: (Slide 4) Now show the students pictures of diamond, graphite (pencil lead), and pure charcoal (in this order). Ask them what element or elements they think each material is made of.
- (Slide 5) After a few guesses share that all these materials are made from exactly one element, carbon.
- Ask the class to discuss with a neighbor or group the following prompt: “Based on what we learned about retinal, develop an explanation of what makes carbon able to form diamond, graphite, and coal. What difference is there between these 3 materials?”
- After a discussion, ask a few students to share their thoughts, and then show the molecular structures for each of the compounds.
- (Slide 6) Have a class discussion about the similarities and differences between these substances and their structures. For example, a property to discuss may include hardness: diamond is the hardest natural material on earth and nearly impossible to scratch or break, graphite spreads out onto paper when rubbed against paper, and charcoal is crumbly. Examples of differences in structure may include that the carbons in diamond all form bonds of equal strengths in a lattice structure that has support and resistance to pressure in every direction. Relate these properties to the molecular structure by asking the students why it might be very easy to break coal or to scratch off chunks.
- Explain: Students should be able to clearly see that structure can have a huge impact on the behavior of a substance. Thus, today we are going to be learning how to describe basic molecular structures and why they form the shapes they do.
- (Slide 7) Students should be familiar with Lewis dot diagrams. You should explain that while Lewis dot diagrams do a great job of describing the atoms and bonds of a molecule, they don’t show us the 3D structure. Tell the students that we are going to learn how to determine the 3D structures of basic molecules from their Lewis dot diagrams.
- Have the students observe a Lewis dot diagram of methane and ask them what shape they think this molecule would be if we were looking at it in 3D or real life.
- (Slide 8) Click the image on the slide will redirect to the PhET simulation, Molecule Shapes. After a short discussion, use Methane as an example from the “Real Molecules” option on the simulation to show the students the space-filling model. Explain that electron pairs want to stay as far away apart from each other as possible and that the electron clouds take up space. The simulation allows you to pull atoms to other angles, but the structure always returns to the predicted molecular shape as the electron pairs repel each other. Explain that there are various patterns of structures that we can learn to predict if we have a Lewis structure to guide us.
- (Slide 9) Draw the Lewis diagram of acetylene and then ask students to guess what the molecule actually looks like. Remind them that every group (atom or lone pair of electrons) wants to be as far apart from the other groups as possible.
- (Slide 10) Using the image on the PowerPoint, or select the “model” option in the PhET simulation to show students the 3D model of acetylene and let them discover if their prediction was correct.
- (Slide 11) Present the chart to the students and briefly go through the geometries with no lone pairs. Ask a few students to come up to the front of the class to volunteer for a demonstration. Blow up balloons and have the students tie them to each other or a focal point one at a time. As you add balloons, you should see the 3 shapes described.
- Possible Extension: Two more balloons can be used to make shapes with expanded octets.
- Explain that lone pairs of electrons have a similar repulsive effect as the atoms in a molecule, and so they act like atoms when determining the geometry or 3D structure. Explain to the students that memorizing all of these geometries is good but not necessary if they remember that the “electron-dense areas” repel each other. The electrons in the bonds or electrons in unshared orbitals cause atoms to be as far apart as possible.
- Possible Extension: Explain that scientists use a variety of methods to describe and sort these geometries. One way we talk about individual atoms in a molecule is by describing them by their steric number. The steric number is the total number of “groups” attached to a central atom. “Groups” are either other atoms or lone pairs. For instance, in the molecule CH4, the steric number of carbon is 4, because carbon is attached to 4 hydrogen atoms. Go over a few more examples and have the students guess the steric number.
- (Slides 12-17) Go over a few practice questions showing students basic molecules either by their Lewis structures or just their formulas and ask students to determine the molecular geometry.
- Elaborate: (Slide 18) Use the molecular structure for Penicillin as another opportunity for students to understand the importance of molecular geometry in the real world. Students can describe what they see that might be unusual or “against the rules” of geometry that they have learned so far.
- (Slide 19) If they can’t find anything then ask them what the ideal bond angles for the carbons in the “square” or 4-membered ring would be based on the chart used earlier in class.
- (Slide 20) The lower-left carbon of the ring would ideally have 120-degree angles due to the presence of only three groups, but instead, it is constrained to 90 degrees. Explain that forcing molecules into constrained shapes with less than ideal bond angles is like packing a spring into a tin can; the bonds strain against each other, wanting to expand to their ideal structure. This square in the middle of penicillin is like a packed spring waiting to burst free.
- Penicillin uses this action-packed square shape to kill bacteria. The 4-membered ring is highly reactive and when it encounters an enzyme that constructs the cell wall of bacteria, it will spring open and attach itself to the active site (see figure below). This prevents the enzyme from building the cell wall. Without a proper cell wall, the bacteria die.
- Penicillin has no effect on human cells because they don’t have cell walls or the enzymes that build them. So, the strained structure of penicillin allows it to directly target and destroy bacterial cells that are harmful to the body without impacting our own cells.
- Evaluate: Have students complete the attached practice worksheet that uses that PhET simulation.
- An Answer Key has been provided for teacher reference.
For the Student
Directions: Constructing Models with PhET
- Open the PhET simulation, Molecule Shapes
- Select “model” on the left.
- You should see an unnamed molecule with three atoms in a line. Click on the single bond button in the upper right corner of the page; this should add an atom with a single bond to the central atom.
- Select “molecule geometry” and “electron geometry” in the box under Name.
Use the model and the simulation to answer each of the following questions:
- When you added an atom, what happened to the shape of the molecule?
- Add another atom. How does the shape change this time?
- Is this the same as what we would expect from a Lewis diagram of one central atom bonded to 4 others? Why is this molecule the same or different from its corresponding Lewis structure?
- Add single bonds until you can’t add any more. What shape is formed? Why do you think the molecule forms this shape as you add atoms?
- Select “Show Bond Angles” from the menu on the bottom right. Try to pull two atoms closer to each other. Notice what happens to the bond angle. Let go of the atom and note the change in the bond angle. What are the atoms doing when they rearrange themselves?
- Click the “x” next to the single bond button until you are back to only 2 atoms bonded to one central atom, now click on the lone pair button as many times as it will allow, paying attention to the changes in shape each time. Do the lone pairs change the molecular shape in the same way that adding atoms does? Explain your reasoning.
- What does this tell us about the reasons why atoms assemble around the central atom in the shapes that they do?