Activity Formal Charge and VSEPR Theory for Expanded Octets

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1 Activity Formal Charge and VSEPR Theory for Expanded Octets Directions: This Guided Learning Activity (GLA) goes over formal charge and the structures of molecules with expanded octets. Part A introduces expanded octets and formal charge. Part B discusses drawing Lewis structures of expanded octets and Part C demonstrates how to predict the electron and molecular geometries of expanded octets using VSEPR Theory. This worksheet is accompanied by an indepth key. See for additional materials. Part A Formal Charge Formal charge is defined as the charge an atom in any compound would have if the electrons were being shared equally between the atoms in a covalent bond. It helps us determine whether a central atom will form an expanded octet. (We will discuss expanded octets in Part B below.) We use the following equation to calculate formal charge: Formal charge (F.C.) = #valence electrons #nonbonding electrons #bonds For example, we can assign a formal charge to each atom in HF: F.C. (H) = = 0 F.C. (F) = = 0 So the formal charge on H is 0 and the formal charge on F is 0. Keep in mind that we assign a formal charge to each individual atom in a molecule, treating individual atoms in a molecule as separate atoms. The sum of each of the formal charges on the atoms in a molecule or an ion will be equal to the total charge of the molecule or ion. Practice: Determine the formal charge of each element in the molecules below: a) ClF 3 We must calculate the formal charge of the chlorine and each of the fluorine atoms. Since the fluorine atoms are equivalent, they will all have the same formal charge. F.C. (Cl) = = 0 F.C. (F) = = 0 We can also label the molecule with the individual formal charge on each atom: Activity Page 1 of 7

2 b) NO 3 We must calculate the formal charge of the nitrogen atoms and each of the oxygen atoms. In this example, two of the oxygen atoms are equivalent (the single bond O) and the other is not (the double bond O=). The equivalent oxygen atoms will have the same formal charge, but the double bond oxygen will have a different formal charge: F.C. (N) = = +1 F.C. (O) = = 1 F.C. (O=) = = 0 We can once again label the molecule with the individual formal charge on each atom: Example #1 Determine the formal charge of each element in the ions below: a) PO 4 3 b) OCN Part B Lewis Structures and Expanded Octets Atoms in the third period and below on the periodic table are able to form expanded octets, which means they are able to have more than 8 electrons in their valence shell. This is because atoms in Period 3 and beyond, the dorbitals can be occupied. We just saw one example of an expanded octet in Part A, the ClF 3 molecule. Although the chlorine atom has 10 electrons in its valence shell, it is stable because it has a formal charge of 0. A molecule is the most stable when the formal charge on each of its atoms is closer to or equal to 0 and/or when any negative formal charge is assigned to the most electronegative atom. Keep in mind that the formal charge on a molecule or ion must be equal to the total charge on the molecule or ion. We also assign formal charges to Lewis structures to help us determine which structure is most stable when there is more than one way to draw a Lewis structure. For example, if we draw the Lewis structure for the sulfate ion, SO 4 2, by following only the octet rule, we would get the following structure: Activity Page 2 of 7

3 Although it does follow the octet rule, the most stable Lewis structure of the sulfate ion actually has 12 electrons in the valence shell of sulfur, as shown below: The reason why the second Lewis structure is more stable is because of the rules of formal charge. The figure below shows the formal charge of each atom on each molecule: Notice that the all five of the atoms in the structure on the left have nonzero formal charge and sulfur has a very high formal charge of +2. Once we form double bonds to the sulfur from two of the oxygen atoms, it significantly reduces the formal charge on the molecule, shown in the structure on the right. Example #2 Draw the Lewis structure for each ion. Expand the octet on the central atom to reduce formal charge if necessary. a) ClO 4 b) SO 3 2 Part C VSEPR Theory for Expanded Octets The Valence Shell Electron Repulsion (VSEPR) Theory states that electron groups position themselves around the nucleus to minimize interactions between the groups. You may recall discussing molecular shapes for molecules that have 2, 3, and 4 electron groups around the central atom in Chemistry 151. When there are more than four electron groups around the central atom, as in the case with expanded Activity Page 3 of 7

4 octets, more kinds of electron and molecular shapes are introduced by the Valence Shell Electron Pair Repulsion (VSEPR) Theory, as shown in the table below. Refer to GLA VSPER Theory for a review of additional molecular and electron geometries. A complete list of the electron and molecular geometries is included at the end of this GLA. Number of electron groups around the central atom Number of bonding electron groups Number of lone pair electrons Electron geometry Molecular geometry Bond Angle Example 5 0 Trigonal Bipyramidal SeeSaw Trigonal Bipyramidal 3 2 TShaped 120, 2 3 Linear 6 0 Octahedral Octahedral Square Pyramidal Square Planar Activity Page 4 of 7

5 3 3 TShaped 2 4 Linear Practice: Draw the Lewis structure of each of the following molecules or ions. Determine the electron and molecular geometry of each. a) ClF 3 The chlorine atom has five electron groups around it. Three of the electron groups are bonds and the other two are lone pairs: Electron geometry: trigonal bipyramidal Molecular geometry: Tshaped b) XeF 4 The xenon atom has six electron groups around it. Four of the electron groups are bonds and the other two are lone pairs: Electron geometry: octahedral Molecular geometry: square planar Example #3 Draw the Lewis structure of each of the following molecules. Determine the electron and molecular geometry of each. a) ICl 5 b) SF 4O Activity Page 5 of 7

6 Part D Extra Practice Draw the Lewis structure of each of the following molecules or ions. Determine the formal charge on each atom and expand the octet on the central atom, if necessary, to reduce formal charge. Determine the electron and molecular geometries. 1. BrF 5 2. SF 6 3. IBr 4 4. SO 2 5. ClO 2 6. XeF 2O 3 7. PO 4 8. SF 4 9. NNO 10. NNN Activity Page 6 of 7

7 Table of Electron and Molecular Geometries Source: Openstax Chemistry Textbook (available free at at Activity Page 7 of 7

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