MODULE 2 Chemical Bonding and Molecular Geometry
CHEM-1112 M2 Explore: Molecular Geometry and Polarity
Learning Objectives
LO2.1.1 Predict molecular structures using VSEPR Theory
LO2.1.2 Interpret atomic dipoles to identify molecular polarity
LO2.1.3 Determine crystal structure of several molecules
What You Will learn
In this module we will review topic pertaining to molecular geometry and polarity
- Practice Drawing Lewis Structures.
- View Introduction Video for the experiment and view the Electronegativity Chart
- Practice identifying Molecular Geometry from Lewis Structures.
- View the Introduction Video and view the Geometries of Molecules handout.
- Practice identifying the polarity of the molecule from the molecular geometry.
- View the Introduction Video and view the Geometries of Molecules handout.
What You Will Do
- Read/watch content and do the practice activities in Canvas.
- Read the Modeling Geometry and Polarity handout on labflow.
- View the Writing Lewis Structures on Labflow.
- Submit Pre-lab worksheet: Pre-Lab Quiz - Modeling Geometry and Polarity
- Submit Mastery questions: Report - Modeling, Geometry and Polarity
- Help each other and have discussions in the Discussion Forum if you have questions.
Electronegativity
Electronegativity generally increases as you move from the bottom left of the periodic table to the top right. The image below shows this visually. You can also explore this sonically using the accessible periodic table at IndepenentScience.com. Just select electronegativity from the drop down. Then use tab to move through the melements and the space or enter button to hear the sound.
Why is it important to know: Molecular formation, shape, and behavior are strongly influenced by the push and pull involved with electrons as they distribute throughout a molecule.
Valence Shell Electrons and the Octet Rule
Think of the element number as the count of how many electrons an atom of an element has. These have an orderly arrangement in the atom. The outermost electrons are special because those determine many of the behaviors of the atom including behaviors that result in the molecular geometry when atoms combine to form molecules. The periodic table arranges elements so that the column number tells you how many electrons are in the valence shell.
Atoms like to have a full valence shell. For most elements, this is going to have 8 electrons. Octa- is a prefix that means 8, so this rule is called the octet rule. For elements that ar in column 8, they already have a full valence shell with 8 electrons - just what they want. So with that satisfied full valence shell, they are not looking to gain any more electrons, so they are un-reactive, or they don't combine with other elements. Column 7 elements are almost there. They want one more to be full. Those elements are typically very reactive and combine with other elements that will give them that extra electron that they want to have. Columns 3A through 6 also want electrons to fill their valence shell, but they need more than one. Column 4 elements have 4 electrons and want to get to 8 so they are looking for four more electrons. So look at the other columns for how many electrons are in their valence shell and figure out how many electrons more they need to get to 8.
Columns 1 and 2 Elements
Column 1 elements have one electron in that outermost shell, the valence electrons. Column 2 elements have two. They have a long way to go to get to 8, so they use a different strategy to get to a full valence shell. They can give away electrons and end up with the full shell right below. These have low electronegativity because they have so little 'hunger' for electrons and give them away.
Columns 3B Through 2B
Visually, it is easy to see that lower, flat area in the middle. We will not consider that area at this time. It plays by different rules that you will get into in higher levels of chemistry.
Practice: How Many Valence Electrons?
Use the periodic table to identify how many electrons are in the valence shell for the following atoms.
B
Boron
3
3 valence shell electrons
Cl
Chlorine
1
1 valence shell electron
Mg
Magnesium
2
2 valence shell electrons
Lewis Dot Diagrams and Lewis Structures
Drawing Lewis Dot Diagrams
Don't confuse Lewis Dot Diagrams with Lewis Structures. They are related, but different. Lewis Dot Diagrams focus on the atom itself. Lewis Structures are for working out how those atoms will combine into molecules and where the electrons tend to hang out when they do.
So, let's focus on the Dot Diagram side of things. When you are working with the last column on the periodic table, it is really easy. Those have all the electrons that can fit in the valence shell. Remember the octet rule? This means they have a full valence shell, so all of those except He (helium) will have 8.
So why is He (helium) an oddball? Well, its atomic number is 2. It only has 2 electrons. It cannot get to 8. Hydrogen and helium play by slightly different rules. They are focused on getting 2 electrons, or in the case of H (hydrogen) it might even give away the one electron that it has, or perhaps you cvan think of it as hydren is the little guy that bullies know they can just take an electron from.
Drawing Lewis Dot Diagrams
Now that you have had a little introduction, watch the video to learn procedures that are followed to create Lewis Dot Diagrams and it will wrap up by showing you why dot diagrams help you to figure out how elements combine to make molecules.
Time:
Practice Lewis Dot Diagrams
Use this interactive activity to practice crafting Lewis Dot Diagrams. If you prefer, you can go directly to the activity (best option for accessibility).
Practice Lewis Structures
Practice using the interactive activity below. Go to the activity directly if the interactive activity below needs more room - Molecule Builder.
VESPR Theory
VESPR - Valence Shell Electron Pair Repulsion
In this video Paul Andersen explains how you can use Lewis Diagrams and VSEPR Models to make predictions about molecules. The Lewis diagrams are a two-dimensional representations of covalent bonds and the VSEPR models show how the molecule could exist in three dimensional space. Pi bonding and odd valence electrons require an extension of this model.
Time: 12:28
Ionic, Covalent (Polar vs. Nonpolar)
Ionic Bonds: Ionic bonds are formed between non-metal atoms and metals through an electron transfer. The first two columns on the periodic table are the metals. They want to get rid of an electron to get to a full valence shell bellow. When one gives up the electron it becomes a positively charged ion because it has one less electron than it has protons. The atom that wants the electron, a non-metal on the right side of the periodic table, will become a negative ion. The resulting ions are held together by the electrical attraction of opposite charges.
Covalent Bonds: These bonds form between two or more non-metal atoms. they don't give electrons, but they will share them.
Polar: When the covalent bond results in an equal sharing of electrons distributed around the molecule.
Non-Polar: When there are atoms that are more electronegative making them greedy for electrons to stay near them causing one area of the molecule to be more negatively charged than other areas of the molecule. You can see this easily in molecular geometry because you will have unbound electrons on one side pushing the now positively charged atoms to the far side of the molecule.
Molecular Geometry Practice
Identify the covalent molecular configuration. Name it, but also identify if it is an example of polar or non-polar covalent bond and for an extra challenge see if you can figure out an example molecule example. Click the play button on the answer side of a card to view the molecule in 3D.
Linear
Polarity: non-polar
Example: CO2
A central atom has three atoms attached each at 120 degrees from each other.
Trigonal Planar
Polarity: non-polar Example: SO3
A central atom has two atoms attached with unbound electrons pushing them toward the other side at less than 120 degrees.
Bent
Polarity: polar Example: H20
Tetrahedral
polarity: non-polar
Example: CH4
A central atom has three atoms and two unbound electrons attached each at less than 109 degrees from each other.
Trigonal Pyramidal
polarity: polar
Example: NH3
A central atom has five atoms attached. though there are no non-bound electrons, two angles are present: 90 and 120 degrees.
Trigonal Bipyramidal
Polarity: polar
Example: PF5
Limitations of VSEPR Theory
While VSEPR theory is a powerful tool for predicting molecular geometry, it has several important limitations that students should understand:
- Transition metals: VSEPR theory works best for main group elements and may not accurately predict geometries for molecules containing transition metals, which often have more complex bonding patterns.
- Multiple resonance structures: Molecules with significant resonance (like benzene) may not follow simple VSEPR predictions due to delocalized electron density.
- Very large atoms: For atoms in periods 4 and beyond, the theory becomes less reliable due to the increased importance of d-orbital participation in bonding.
- Molecules with odd numbers of electrons: Free radicals and other species with unpaired electrons may not follow typical VSEPR geometry predictions.
- Weak interactions: VSEPR doesn't account for weak intermolecular forces that can influence molecular shape in crystalline solids or concentrated solutions.
- Limited information: It gives no information about bond lengths or bond energies. The theory only predicts shape and bond angles.
Despite these limitations, VSEPR theory remains an excellent starting point for understanding and predicting molecular geometry in most common chemical situations.
Lab Report
You will watch the videos and complete the lab in Labflow.