For each of the following molecules, write the Lewis structure(s), predict the molecular structure (including bond angles), give the expected hybrid orbitals on the central atom, and predict the overall polarity $$ \text {a} C F_{4} \quad \text { e. BeH }_{2} \quad \text { i. } \operatorname{KrF}_{4} $$ $$ \text {b} \mathrm{NF}_{3} \quad \text { f. } \mathrm{TeF}_{4} \quad \text { j. SeF }_{6} $$ $$ \text {c} \mathrm{OF}_{2} \quad \text { g. AsF_ } \quad \text { k. } \mathrm{IF}_{5} $$ $$ \text {d} \mathrm{BF}_{3} \quad \text { h. } \mathrm{KrF}_{2} \quad \text { L. } \mathrm{IF}_{3} $$

#tag_title# f. TeF4#tag_content# Lewis structure: Draw a tellurium atom in the center. Attach four fluorine atoms to the tellurium atom with single bonds. There are two lone pairs of electrons on the tellurium atom. Molecular structure: See-saw (six electron pairs) Bond angles: Approximately 90° and 120° Hybrid orbitals on central atom (Tellurium): sp3d2 hybridized Overall polarity: Polar (due to unsymmetric distribution of polar Te-F bonds and the presence of lone pairs) #tag_title# g. AsF3#tag_content# Lewis structure: Draw an arsenic atom in the center. Attach three fluorine atoms to the arsenic atom with single bonds. There is a lone pair of electrons on the arsenic atom. Molecular structure: Trigonal pyramidal (four electron pairs) Bond angles: 107° (approx.) Hybrid orbitals on central atom (Arsenic): sp3 hybridized Overall polarity: Polar (due to the presence of a lone pair, causing unsymmetric distribution of polar As-F bonds) #tag_title# h. KrF2#tag_content# Lewis structure: Draw a krypton atom in the center. Attach two fluorine atoms to the krypton atom with single bonds. There are three lone pairs of electrons on the krypton atom. Molecular structure: Linear (five electron pairs) Bond angles: 180° Hybrid orbitals on central atom (Krypton): sp3d hybridized Overall polarity: Nonpolar (due to the symmetric distribution of polar Kr-F bonds) #tag_title# i. KrF4#tag_content# Lewis structure: Draw a krypton atom in the center. Attach four fluorine atoms to the krypton atom with single bonds. There are two lone pairs of electrons on the krypton atom. Molecular structure: Square planar (six electron pairs) Bond angles: 90° Hybrid orbitals on central atom (Krypton): sp3d2 hybridized Overall polarity: Nonpolar (due to symmetric distribution of polar Kr-F bonds) #tag_title# j. SeF6#tag_content# Lewis structure: Draw a selenium atom in the center. Attach six fluorine atoms to the selenium atom with single bonds. Molecular structure: Octahedral (six electron pairs) Bond angles: 90° Hybrid orbitals on central atom (Selenium): sp3d2 hybridized Overall polarity: Nonpolar (due to symmetric distribution of polar Se-F bonds) #tag_title# k. IF5#tag_content# Lewis structure: Draw an iodine atom in the center. Attach five fluorine atoms to the iodine atom with single bonds. There is a lone pair of electrons on the iodine atom. Molecular structure: Square pyramidal (six electron pairs) Bond angles: 90° (approx.) Hybrid orbitals on central atom (Iodine): sp3d2 hybridized Overall polarity: Polar (due to the presence of a lone pair, causing unsymmetric distribution of polar I-F bonds) #tag_title# l. IF3#tag_content# Lewis structure: Draw an iodine atom in the center. Attach three fluorine atoms to the iodine atom with single bonds. There are two lone pairs of electrons on the iodine atom. Molecular structure: T-shaped (five electron pairs) Bond angles: Approximately 90° and 180° Hybrid orbitals on central atom (Iodine): sp3d hybridized Overall polarity: Polar (due to unsymmetric distribution of polar I-F bonds and the presence of lone pairs)