Draw the Lewis structure for AsF
6
?
in the window below and then answer the questions that follow. – Do not include overall ion charges or formal charges in your drawing. Draw the Lewis structure for IBr
3
?
in the window below and then answer the questions that follow. What is the electron-pair geometry for I in IBr
3
?
? What is the the shape (molecular geometry) of IBr
The Correct Answer and Explanation is :
Lewis Structures of AsF₆⁻ and IBr₃
Step 1: Drawing the Lewis Structure for AsF₆⁻
- Count the valence electrons
- Arsenic (As) has 5 valence electrons
- Fluorine (F) has 7 valence electrons, and there are six fluorine atoms
- The negative charge (-1) adds one extra electron
- Total valence electrons = 5 + (7 × 6) + 1 = 48 electrons
- Determine the central atom and arrange the atoms
- As is less electronegative than F, so As is the central atom
- Distribute six fluorine atoms around arsenic
- Draw single bonds between As and F
- Each single bond uses 2 electrons
- 6 bonds × 2 electrons = 12 electrons used
- Remaining electrons = 48 – 12 = 36 electrons
- Distribute remaining electrons to satisfy octets
- Each fluorine needs 6 more electrons to complete its octet
- 6 fluorine atoms × 6 electrons = 36 electrons (all remaining electrons are used)
- Check the octet rule and formal charge
- As has 12 valence electrons (expanded octet, which is allowed for Group 15 elements in Period 4 and beyond)
- Formal charges: As has a -1 formal charge, matching the given ion charge
Thus, the Lewis structure of AsF₆⁻ has As in the center with six F atoms surrounding it, each bonded by a single bond, and a negative charge on the ion.
Step 2: Drawing the Lewis Structure for IBr₃
- Count the valence electrons
- Iodine (I) has 7 valence electrons
- Bromine (Br) has 7 valence electrons, and there are three Br atoms
- Total valence electrons = 7 + (7 × 3) = 28 electrons
- Determine the central atom and arrange the atoms
- I is less electronegative than Br, so it is the central atom
- Distribute three Br atoms around I
- Draw single bonds between I and Br
- 3 bonds × 2 electrons = 6 electrons used
- Remaining electrons = 28 – 6 = 22 electrons
- Distribute remaining electrons to satisfy octets
- Each Br needs 6 more electrons to complete its octet
- 3 Br × 6 electrons = 18 electrons used
- Remaining electrons = 22 – 18 = 4 electrons
- Place remaining electrons on the central atom (I)
- The 4 leftover electrons go on I as two lone pairs
- Iodine now has 5 electron groups (3 bonds + 2 lone pairs)
Electron-Pair Geometry and Molecular Geometry of IBr₃
- Electron-Pair Geometry
- The electron-domain (steric) number is 5 (three bonding pairs + two lone pairs).
- The electron-pair geometry is trigonal bipyramidal.
- Molecular Geometry (Shape)
- The two lone pairs occupy equatorial positions to minimize electron repulsions.
- This results in a T-shaped molecular geometry.
Explanation (300 words)
The Lewis structures of AsF₆⁻ and IBr₃ help predict molecular geometry based on VSEPR (Valence Shell Electron Pair Repulsion) theory. For AsF₆⁻, the arsenic atom is surrounded by six fluorine atoms, forming an octahedral molecular shape with 90° bond angles. The molecule is nonpolar because the symmetrical distribution of fluorine atoms cancels out dipole moments.
For IBr₃, iodine is the central atom, forming three single bonds with bromine atoms and having two lone pairs. The presence of lone pairs affects the molecular shape. Although the electron-pair geometry is trigonal bipyramidal, the actual molecular geometry is T-shaped because lone pairs occupy equatorial positions to reduce repulsions. The bond angles are slightly less than 90° due to lone pair repulsions. The molecule is polar because of its asymmetric shape, leading to a net dipole moment.
These structures illustrate key concepts in molecular geometry:
- Lone pairs cause deviations from ideal geometries
- Symmetry determines molecular polarity
- Expanded octets occur in elements beyond Period 2, such as arsenic and iodine
Understanding these structures helps predict chemical reactivity, intermolecular forces, and physical properties like boiling points and solubility.