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Lewis Structure and Hybridization of HSCN (thiocyanic acid, hydrogen thiocyanide)

HSCN is the corresponding acid of SCN-, which is the thiocyanide ion. All four of these atoms are non-metals, the bonds between them will be covalent, and it will make a molecular compound.

Carbon brings four valence electrons with it, and so needs four more to complete its valence shell. Nitrogen brings five valence electrons, and so it needs three to completed its octet. Carbon and nitrogen share three electrons each (with each other) and a triple bond forms between them.

Meanwhile, carbon needs one more electron and sulfur can provide it. This also gives sulfur a seventh electron, and it gets its eighth from hydrogen.

HSCN has a single bond between H and S, a single bond between S and C, and a triple bond between C and N.

What is the hybridization of C in HSCN?

Carbon is triple-bonded to nitrogen; this requires two pi bonds and that means it requires two leftover p orbitals after hybridization; so the C is “sp” hybridized.

What is the hybridization of N in HSCN?

N is also triple bonded, so it is “sp” hybridized as well. The sigma bond (first bond between the two is sigma) and the lone pair are where the “hybridized sp orbitals” are used here.

What is the hybridization of S in HSCN?

The sulfur atom has NO double or triple bonds. It has two single bonds (both sigma) and two lone pairs. This means is does not need any leftover p orbitals (which would need to stay unhybridized) and so sulfur is “sp3” hybridized here.

What is the VSEPR shape of HSCN?

The carbon atom is only bonded to two other atoms and has no lone pairs; this gives it AX2 geometry, and so there is a linear arrangement around the carbon atom.

The sulfur, on the other hand, has two sigma bonds (single bonds on either side) AND two lone pairs. This gives it Ax2E2 geometry and so it is V-Shaped/Non-linear/Bent.

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Lewis Structure and Hybridization of HCN (hydrocyanic acid, hydrogen cyanide)

HCN has a hydrogen atom single-bonded to a carbon atom, and that carbon atom is triple-bonded to a nitrogen atom.

These are all non-metals, so the bonds are covalent and HCN is therefore a covalent (aka Molecular) structure.

Carbon brings four valence electrons with it; it needs four more to complete its valence shell. Hydrogen shares one electron with it, and nitrogen shares three. This completes carbon’s octet.

Carbon likewise shares one electron back with Hydrogen (this complete’s hydrogen’s outer shell of two electrons, aka Doublet) and carbon shares three electrons back with Nitrogen. This completes nitrogen’s octet.

You can watch this structure get drawn below, or you can scroll to the bottom of this page for a completed structure.

What is the hybridization of Carbon in HCN?

Carbon is triple-bonded to nitrogen, and so there are two pi bonds (Remember: The first bond between any two atoms is a sigma bond, and the second/third bonds are pi bonds). This means two p orbitals are required to be left over after hybridization.

2 pi bonds = 2 leftover p orbitals.

This means only ONE of carbon’s p orbitals is available to hybridize, and so the hybridization of C in HCN is “sp”.

What is the hybridization of N in HCN?

Nitrogen is triple-bonded to carbon, and so two pi bonds are required here as well. This means only one of nitrogen’s p orbitals is available to be hybridized, and so the hybridization of nitrogen in HCN is “sp”.

What is the molecular shape (VSEPR shape) of HCN?

Because the carbon is connected to two atoms, with no lone pairs on that central carbon, the geometry is AX2, which is “linear”. The bond angle is 180 degrees.

This image shows HCN with its constituent atoms, the sharing of electrons from one atom to another, and the final Lewis Structure showing a single bond to hydrogen and a triple bond to nitrogen.
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Lewis Structure of BCl3, Boron Trichloride (and Hybridization)

Boron Trichloride, BCl3, has three chlorine atoms surrounding a single boron atom.

This is because each chlorine brings seven valence electrons with it, and needs just one more electron to complete its octet.

Boron has three valence electrons to start with, but does not need a full eight electrons to be stable. It is a violation of the octet rule, but this is the way things are.

So, Boron shares ONE electron with each of three chlorine atoms, and each chlorine shares one electron with Boron:

The Lewis Structure of Boron Trichloride (BCl3) has three chlorine atoms surrounding a single boron atom.

This is a trigonal planar arrangement and implies that the boron must be sp2 hybridized. The extra unhybridized p orbital is empty, but its presence is what keeps the three sp2 orbitals separated by exactly 120 degrees.

I have a video where I draw this Lewis Structure, if you’re a visual learner:

Now, in reality, solid BCl3 is more complicated. The lone pairs on each chlorine atom are attracted to the wide-open slightly-positive charge on the Boron atom, and a Lewis Acid-Base reaction happens: That’s fancy chemistry talk for chlorine sharing its lone pair with boron.

This gives the boron atoms in solid BCl3 a tetrahedral geometry; since each of them actually connect with FOUR chlorines each. It makes the entire structure more like a lattice, rather than being individual molecules.

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Lewis Structure of BF3 (and hybridization)

BF3, boron trifluoride, is a tricky molecule to draw because Boron is an exception to the octet rule. It does not need eight electrons in its outer shell, although it can hold eight just like most other non-metals.

The Lewis Structure of BF3, boron trifluoride, has one boron atom in the centre, and three fluorine atoms surrounding it. Each of the fluorine atoms shares ONE electron with the boron, and the boron in turn shares ONE electron with each of the fluorines.

This gives it a trigonal planar shape, and its hybridization looks like it’s sp2:

This is the simplest Lewis Structure for Boron Trifluoride and it’s probably the one your teacher is looking for you to draw.

You can watch this structure get drawn here:

HOWEVER, in the real world things are not so simple.

At low temperatures, boron trifluoride forms a solid, and the arrangement of the atoms change. One fluorine from a BF3 molecule will donate one of its lone pairs to give a different Boron atom a full octet .. and then a fluorine connect to that Boron will complete the octet on another… etcetera.

In the end, each boron has fluorine atoms tetrahedrally arranged around it. This would imply its hybridization has changed to sp3, but also makes the substance seem more like a covalent network (also known as giant covalent) … or even like an ionic compound, if you wanted to bold.

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Lewis Structure of Al2O3, Aluminum Oxide

Aluminium oxide is a solid ionic compound, made from atoms of one metal (Aluminum) that have lost three electrons each to become +3 cations, and atoms of a non-metal (oxygen) which have gained two electrons each to become -2 anions.

The numbers 3 and 2 have a lowest common multiple of 6. This means you need TWO aluminum atoms (giving away 2×3=6 electrons total) and THREE oxygen atoms (accepting 3×2=6 electrons total).

Because the electrons are LOST by the metal, and GAINED by the non-metal, there has been a transfer of electrons and this makes it an ionic compound by definition.

Here, two Aluminum atoms are giving away 3 electrons each, and three oxygen atoms are accepting two electrons each.

You can watch this Lewis Structure get drawn below.

I created a video for this post! You’re welcome !

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Lewis Structure of Magnesium Fluoride (MgF2)

Magnesium fluoride (MgF2) is not a molecular/covalent compound; magnesium has a low electronegativity (it is a metal, after all) and fluorine has a high electronegativity (it is a non-metal, halogen, and has the highest electronegativity of all atoms on the table). So, a transfer of electrons occurs.

Magnesium, when it bonds to a non-metal, always loses two electron to become a +2 ion. This is because the neutral atom has two valence electrons in its outer shell, and to satisfy the octet rule, it wants to lose those two electrons. This makes it a cation.

Fluorine atoms, on the other hand, carry seven valence electrons in the outer shell of their neutral atoms. So their preference is to gain one electron each, and get a -1 charge; this negatively-charged particle is called an anion.

We require one magnesium atom, giving away its two electrons, giving one electron each to two fluorine atoms:

And you can watch this process occur here:

This ionic structure is similar to MgO, which is also ionic. It is not to be confused with other three-atom compounds, made entirely of non-metals, like SF2, which is a covalent compound.

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Lewis Structure of MgO, Magnesium Oxide

Magnesium oxide is an ionic compound. Its formula unit (MgO) is made from one magnesium atom, which loses two electrons to become a +2 cation, and one oxygen atom, which gains those two electrons to become a -2 anion.

The Lewis Structure can show the transfer of electrons from metal (low electronegativity) to non-metal (high electronegativity). Observe below.

You can watch this getting drawn in the video below.

Magnesium oxide does not hybridize to do this, so if you’re looking for the molecular shape, it does not exist. Remember, ionic compounds are crystal lattices, which means the ions (Mg+2 and O-2) surround each other and hold together with ionic bonds. It crystallizes in a face-centred cubic structure, which has an octahedral coordination geometry – each Mg+2 is surrounded by six O-2 ions, and each O-2 ion is likewise surrounded by six Mg+2 ions.

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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