Lewis Structures Part II, Study notes of Geometry

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Lewis Structures Part II

Molecular Geometry – the Valence Shell Electron Pair Repulsion (VSEPR) Theory The electron groups around the central atom repel each other and therefore prefer to be as far apart from each other as possible. This is the main idea of the VSPER theory. We can apply the VSEPR theory to predict the molecular shape/geometry of a molecule.

1. Draw the Lewis structure for the molecule in question.
2. Count the total number of electron groups on the central atom. Add the number of atoms bonded to the central atom and the number of lone pairs on the central atom – this is the total number of electron groups. Note that multiple bonds to one outer atom still count as one electron group.
3. The arrangement of the electron groups is determined by minimizing the repulsions between them.
4. Remember that lone pairs require more space than bonding pairs. Therefore, choose an arrangement that gives lone pairs as much room as possible. The table attached shows the relationship between the number of electron pairs and the molecular geometry. Polarity of a molecule A covalent bond is polar if there is a difference in electronegativity between the bonded atoms. An entire molecule will be polar if the bond dipoles do not cancel. Polar molecules have a positive and a negative end and behave like tiny magnets. The shape of the molecule determines if dipoles cancel or not. Polar molecules – bond dipoles do not cancel. Have a lone pair on the central atom OR different outer atoms. Non-polar molecules – bond dipoles cancel. Have no lone pairs on the central atom AND have all outer atoms the same. Formal Charges The formal charge of any atom in a molecule is the representation of electron distribution on the atom. Remember that the formal charge does not represent the real charge on the atom. It is a fictitious charge assigned to each atom that helps in finding the best Lewis structure for a molecule. The best Lewis structure will  Have the lowest possible formal charge on each atom  Put the negative formal charge on the most electronegative atom (and a positive formal charge on the least electronegative atom) Calculating formal charges

• Sum all the electrons in the lone pairs belonging to that atom
• Add to this half of the bonding electrons
• Subtract this total from the number of valence electrons for that atom to get its formal charge. In the above example for CO 2 , the Lewis structure on the Left is the better one as it places a formal charge of zero on each atom. Recall that the formal charges sum to zero for a molecule and to the charge for an ion.

Resonance Sometimes one Lewis structure is not enough to describe a molecule completely. For example, ozone (O 3 ) can be represented by the following two structures: One might expect ozone to have one single bond and one double bond based on either of the above Lewis structures. However, experimentally it is found that both the bonds are equivalent and intermediate between a single and a double bond. Thus, the true structure of ozone is a resonance hybrid of the above two Lewis structures. We represent resonance by drawing the two structures with a double headed arrow between them. Remember that although we may draw two resonance structures, there is actually only one structure that is a hybrid of the two drawings – the molecule does not "flip" back and forth, but rather is permanently somewhere between the two structures. Each bond in ozone is a combination of one single and half a double bond. ↔ In the above example, both the Lewis structures are equivalent and contribute equally to the true structure of ozone. However, sometimes some of the Lewis structures are preferred over others (see the discussion on formal charges above). In that case, the preferred Lewis structure(s) contribute more to the overall structure of the molecule. Hybridization To explain molecular geometries, we assume that orbitals mix together to form new orbitals. This process of mixing atomic orbitals is called hybridization. The new orbitals are called hybrid orbitals. The number of hybrid orbitals formed will be equal to the total number of orbitals that are mixing together.

Name_________________________

Team Name ______________________

CHM111 Lab – Lewis Structures (2) – Grading Rubric

Criteria Points possible Points earned

1. Correct Lewis Structure drawn with correct geometry, bond angles and polarity.

2. Correct FC and favored structure 2
3. Resonance, FC and favored structure 3
4. Correct dipoles 1.
5. Dipoles drawn and explained. 1 Total 20 Subject to additional penalties as per the instructor

CHM 111 Lewis Structures Part 2 Name: ___________________________ Molecular Formula Lewis Structure Number of electron groups on central atom Number of lone pairs on central atom Molecular Geometry Bond angle(s) Polar? Yes or No Hybrid- ization NBr 3 CH 3 Br XeF 2 IF 5 CH 2 S