For each of the following molecules, write the Lewis...

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 $$ \text {a} C F_{4} \quad \text { e. BeH }_{2} \quad \text { i. } \operatorname{KrF}_{4} $$ $$ \text {b} \mathrm{NF}_{3} \quad \text { f. } \mathrm{TeF}_{4} \quad \text { j. SeF }_{6} $$ $$ \text {c} \mathrm{OF}_{2} \quad \text { g. AsF_ } \quad \text { k. } \mathrm{IF}_{5} $$ $$ \text {d} \mathrm{BF}_{3} \quad \text { h. } \mathrm{KrF}_{2} \quad \text { L. } \mathrm{IF}_{3} $$

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
$$
\text {a} C F_{4} \quad \text { e. BeH }_{2} \quad \text { i. } \operatorname{KrF}_{4}
$$
$$
\text {b} \mathrm{NF}_{3} \quad \text { f. } \mathrm{TeF}_{4} \quad \text { j. SeF }_{6}
$$
$$
\text {c} \mathrm{OF}_{2} \quad \text { g. AsF_ } \quad \text { k. } \mathrm{IF}_{5}
$$
$$
\text {d} \mathrm{BF}_{3} \quad \text { h. } \mathrm{KrF}_{2} \quad \text { L. } \mathrm{IF}_{3}
$$

Key Concepts

Lewis Structures

Lewis structures provide a visual representation of the valence electrons in a molecule, illustrating how atoms are bonded together (through shared electron pairs) and the location of lone pairs. This is a foundational step in understanding molecular structure and reactivity.

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the three-dimensional arrangement of atoms in a molecule. By considering the repulsion between electron pairs (both bonding and non-bonding), VSEPR theory helps in determining the molecular geometry and bond angles.

Hybridization

Hybridization refers to the mixing of an atom's atomic orbitals to form new hybrid orbitals with a specific geometry. These hybrid orbitals explain the observed bond angles and geometrical arrangement of atoms around a central atom in molecules.

Molecular Polarity

Molecular polarity is determined by the differences in electronegativity between bonded atoms and the overall geometry of the molecule. The spatial arrangement of polar bonds and any lone pairs dictate whether the molecule has a net dipole moment, which influences its physical properties and interactions.

Transcript

00:01 Okay, for each of these molecules, we'll draw the lewis structure and then determine the expected hybrid orbitals that will be used for bonding around the central atom, and then predict whether or not these molecules are polar.

00:18 For the first one, we'll draw the lewis structure and recognize that it is a symmetrical molecule with the tetrahedral geometry, therefore it will not be polar.

00:29 The bond angles will be about 109 .5 degrees, and the central atom will be sp3 hybridized.

00:37 And we'll go to nitrogen trifluoride.

00:39 Nitrogen trifluoride has a tetrahedral electron group geometry, but the molecular geometry, with one of them being a lone pair, is going to be trigonal parametal.

00:49 The bond angles will be something just under 109 .5 degrees, and the molecule will be polar, because all of the nitrogen -fluorine bond polarities do not cancel.

01:01 Because it's not all of the electron groups are the nitrogen fluorine bonds, so their polarities don't, it's not symmetrical, and their polarities don't cancel.

01:15 Then we look at oxygen diphluoride, and we see there are four electron groups around oxygen.

01:21 Two of them are bonding, so this is bent.

01:23 It's something just under 109 .5 degrees, and because it's bent, the bond polarities of oxygen and fluorine do not cancel, so it's polar.

01:35 Boron trifluoride and with boron trifluoride we see that there are only three electron groups around boron it only has six valence electrons but boron is one of the exceptions or six valence electrons for some molecules is acceptable so this is triginal planar at 120 degrees sp2 hybridized and because of the symmetry all of the boron fluorine bond polarities cancel so the molecule is non -polar then we go to beryllium with two hydrogens.

02:09 Barillium only has four valence electrons surrounding it, two electron groups.

02:14 Burrillion is another one of the exceptions where some molecules of beryllium are fine with just four valence electrons surrounding it.

02:23 This would be one of them.

02:24 So the molecule then with only two electron groups is linear at 180 degrees and it has sp hybridization around beryllium and because of the symmetry the molecule is non -polal then we look at trilium tetrafluoride, and we end up placing another lone pair around trillium because we have the an extra set of electrons and that extra set of electrons typically is placed around the central atom.

03:00 So now we have five electron groups around the central atom with five electron groups around the central atom.

03:07 One of them being a lone pair.

03:10 This has a seesaw geometry.

03:12 There's actually two bond angles.

03:14 One is 120 degrees.

03:16 The other is 90 degrees.

03:18 If you go back and review the seesaw geometry, with five electron groups around the central atom, this is d.

03:25 Sp3 hybridized.

03:26 The molecule is not symmetrical, although all the bonds are the terrillion fluorine bonds, because it's not symmetrical, the molecule is polar.

03:41 Okay, then for the next one, we've got arsenic pentafluoride.

03:49 If we take all the valence electrons that are available to us, we'll see that our arsenic is just going to have five bonding groups surrounding it.

03:59 Five bonding groups correspond to a trigonal bi -paramidal geometry with bond angles of 120 and 90 degrees, a dsp3 hybridization, and because of the symmetry of the molecule, this molecule is also non -polar...

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