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Lewis Structure of SF4 and Hybridization of SF4 (sulfur tetrafluoride)

Sulfur tetrafluoride (SF4) is four fluorine atoms surrounding a central sulfur atom. This is an example of a molecule that does not follow the octet rule.

Sulfur brings six valence electrons, and normally this means it has two unpaired electrons to share in covalent bonds. However, in sulfur tetrafluoride (SF4) there are four bonding pairs, and so four of those valence electrons are involved in single bonds with fluorine atoms.

Lewis structure of SF4 (sulfur tetrafluoride). Four bonding pairs and one lone pair on the sulfur atom.

How is it possible for sulfur to violate the octet rule? It’s because sulfur’s valence electrons on in the third energy level (shell). This means that the 3d orbitals can be involved. The 3s orbital, all three 3p orbitals AND one of the 3d orbitals all combine together to make five equal-energy (degenerate) sp3d hybrid orbitals.

Five atomic orbitals hybridize to allow sulfur to expand it octet.

This gives sulfur a hybridization of sp3d. The VSEPR notation here is AX4E, which corresponds to “sawhorse” geometry.

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Lewis Structure of SF2 and Hybridization of SF2 (Sulfur Difluoride)

Sulfur difluoride (SF2) is one sulfur atom connected to two fluorine atoms. They are both non-metals, so they share electrons to form covalent bonds.

Sulfur brings 6 valence electrons with it, and so needs two extra for have a full octet.

Fluorine brings 7 valence electrons with it, and so needs one extra to complete its octet.

This means that sulfur can share ONE electron with EACH of two fluorine atoms, completing all of their octets simultaneously.

Lewis Structure of SF2 (sulfur difluoride). Sulfur is single-bonded to each of two fluorine atoms, and has two lone pairs as well.

In the end, sulfur is single-bonded to each of two fluorine atoms (this is two bonding pairs) and has two lone pairs on it as well. This gives it a VSEPR notation of AX2E2, which is angular / bent / non-linear geometry.

What is the Hybridization of S in SF2?

The sulfur atom has no double bonds, which means that no pi-bonds are needed. This means its hybridization is sp3.

What is the Hybridization of F in SF2?

The hybridization of the fluorine atoms is sp3 as well, since they also do not have any double or triple bonds.

What is the Bond Angle in SF2?

Sulfur has two single bonds and two lone pairs around it, and this is four things, so the electron pair geometry is tetrahedral. Due to the lone pairs, most teachers want to hear that the bond angle is “less than 109.5 degrees”, since the lone pairs repel the bonding pairs and push the single bonds together more than they do in a tetrahedral molecule like CH4. In the case of SF2, the actual bond angle is just 98 degrees.

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Lewis Structure of Iron (II) Oxide, FeO

Iron (II) oxide’s Lewis Structure is among one of the easiest to draw. The iron atom, because it has a +2 charge in this compound, is drawn with two valence electrons – and since it is a metal, it wants to give them away (“lose them”).

Oxygen, by contrast, is a non-metal with six valence electrons – that’s just two short of a full valence shell. The high electronegativity of oxygen attracts the electrons that iron wants to give away, and the two bond together.

“Bond together” in this case means making an ionic compound. There are no MOLECULES of iron (II) oxide – instead there is a crystal lattice of alternating positive and negative ions.

The iron atoms, which each lost two electrons, have a +2 charge and become cations. The oxygen atoms, which each gained two electrons, have a -2 charge and become anions.

This ionic compound has the same number of cations and anions in the crystal, since they occur in a 1:1 ratio.

This is the Lewis Structure of Iron (II) Oxide
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Lewis Structure of iron (III) oxide, Fe2O3, step-by-step

The Lewis Structure of iron (III) oxide, Fe2O3, consists of five ions: Two iron ions with a +3 charge each, and three oxygen ions with a -2 charge each.

Iron III oxide is an ionic compound, because it consists of a metal and non-metal. These types of atoms have a big enough electronegativity difference that electrons are *transferred* from one atom to another, rather than being shared.

To begin, we note that the iron atoms need a charge of +3 … this is evident in the chemical formula (Fe2O3), since the “3” on the O had to have been criss-crossed down from the iron. It is also obvious in the name: The Roman numeral (III) after iron indicates that its charge in this compound is +3.

So we draw two iron atoms with three valence electrons each.

Each oxygen atom brings 6 valence electrons (Oxygen is in Group 16 and is two electrons short of a full octet in its outer shell).

This image shows the transfer of electrons from Iron atoms to Oxygen Atoms, and the complete lewis structure of Iron (III) Oxide, Fe2O3.

One iron atom gives two electrons to an oxygen, but then still has one electron left. So it gives that electron to another oxygen, but that oxygen requires one more as well. So another iron atom must come into play; it gives one electron to complete the second oxygen’s octet and then gives away both of its leftover electron to a third oxygen.

This is an ionic compound, so there is no “hybridization of oxygen in Fe2O3” – it is instead a lattice of alternating positive and negative ions.

This structure likely reminds you of the lewis structure of Calcium Bromide (CaBr2), which was also ionic.