HYBRIDIZATION AND SHAPES OF SIMPLE MOLECULES INCLUDING CARBON COMPOUNDS

Contents

• Introduction
• Definition of Hybridization
• Atomic orbitals
• Hybridization in chemistry
• New hybrids formed by hybridization
• Rules of Hybridization in Atomic Orbitals
• Sigma bonds (σ-bond) and pi bonds (π-bond)
• Types of Hybridization
• Hybridization of Carbon Compounds
• Shapes of simple molecules
• Summary
• Key Terms
• Recommended Videos
• Test Questions
• Discuss and Explain

Learning Objectives

By the end of this section you should be able to:

• Explain hybridization of atomic orbitals.
• Explain how hybrids are formed.
• Explain the types of hybrids.
• Explain the structures of molecules using lewis dot.
• Explain the shapes of molecules using VSEPR theories.

Introduction

I've always noticed that whenever peeps hear the word hybridization in chemistry for the first time, they always find somewhere to hide and I always laugh because the concept is not actually as hard as it sounds and you too can get the concept today; right now; if you would pay close attention and read slowly through this post.

What is Hybridization

Firstly, let me deviate a bit off chemistry. This is most probably not your first time to hear the word 'hybrid' right? I mean you must have heard of 'hybrid computers' in ICT, 'hybrid offsprings' in Biology and Breeding, 'hybrid cars' in Automobile Technology. 

All these hybrids have something in common. 

What is that? 

All hybrids are made up of distinct (or different) materials.

Different materials are combined to give rise to a hybrid. The hybrid will then have the characteristics of each of the materials that formed it.

You can see that a hybrid offspring is gotten from interbreeding varieties or different species of plants or animals.

A hybrid computer is developed from combination of both analog computer and digital computer.

Hybrid cars are engineered from combination of both diesel (or fuel) engine and an electric engine. 

Just as hybrids in all of these various fields are made up of several components, so is a hybrid in chemistry also made up of several components.

By now, the question that should be in your mind is "what components in chemistry give rise to a hybrid?"

This will bring us to a concept called Atomic Orbitals. Many of you know this already, but for the sake of those who don't or may have forgotten, let me discuss briefly on it.

What is/are Atomic Orbitals

We all know that an atom is the smallest unit of an element, and an atom (though it is very small, yet it still) possesses the full chemical properties of its element.

PRO TIP: No two elements have exactly the same chemical properties because atoms of different elements are never the same.

We also know that what make up every atom are protons, electrons and neutrons (neutron is not found in hydrogen -). Protons and neutrons fuse together to make up a dense material (in the center of the atom), this is called a nucleus, you remember that right? 

Structure of an atom

Where are the electrons then? 

The electrons are actively revolving around this dense mass of neutron and proton (nucleus) in what we call orbitals. 

Orbitals are regions where electrons can be found around the nucleus of an atom.

The movement of these electrons are somewhat not stable or predictable. Why?

I will be using our solar system as a perfect example. Several planets move around the sun right?

You can plot a real or definite path for the movement of a planet around the sun and this is called an orbit. 

A simple view of an atom looks similar and you may have even pictured electrons as perfectly orbiting around the nucleus - well, you're in the uni now and that notion has to change right now because the fact is that it is impossible to draw orbits for electrons. 

You can draw an orbit for a planet circling round the sun but you cannot do so for electrons. 

To plot an orbit (or path) for something, you need to know exactly where it is and where it will be in the next second or instant. 

You can't do this for electrons and this is backed up by Heisenberg's uncertainty principle that freely explains that you cannot know with certainty both 'where an electron is' and 'where it is going next' - [the principle actually says that it is impossible to precisely define both the location (or position) and momentum of an electron at the same time].

That makes it impossible to plot an orbit for an electron around a nucleus. Is this a major problem? No! If something is impossible you have to accept it and find a way around it - Jim Clark.

Oh I've digressed a bit, I just needed you to get that because it may not return to my head again  and I hope you do?

Yes? Let's ride on.

I was talking about atomic orbitals and I believe you know it now - regions where electrons can be found around the nucleus of an atom.

What are electrons doing in the orbital? 

Someone once randomly asked this question. 

Well, a professional answered the question, he said "We don't know, we can't know, and so we just ignore the problem!" "All you can say is that if an electron is in a particular orbital it will have a particular definable energy"

We basically have 4 electron orbitals or atomic orbitals:

We have s orbitals, p orbitals, d orbitals and f orbitals.

s orbitals appear to be somehow spherical.

p orbitals are dome shaped, like you inflate and tie two balloons together, or more less like figure 8.

d orbitals are found in transition elements and they have rather complex shapes - 

while f orbitals have most complex shapes - found in Lanthanides and Actinides series: you may not be needing them at this stage.

I hope you're following me, because you really need to get the concept of atomic orbitals for you to fully understand hybridization. 

Wanna skip explanation? You can watch the videos in the recommended Videos section.

How electrons are arranged in orbitals?

Only 2 electrons can occupy an orbital.

s orbitals can carry only two electrons - anything more than that is an abomination.

p orbitals occur in three subdivisions - x, y and z. These three subdivisions are arranged at right angles to one another. Each of these orbitals (x or y or z) can take only two electrons.

d orbitals occur in more complex subdivisions - 5 subdivisions and each subdivision can contain only 2 electrons.

Now, if you recall the atomic principles you learned during your secondary school times, electrons start filling up orbitals from a lower energy level before moving to higher energy levels (we have energy levels 1, 2, 3...).

Energy level 1 has just a s orbital.

Once this orbital gets filled up, electrons move to energy level 2

Energy level 2 has a s orbital and px, py, pz orbitals.

This means that if Hydrogen, being the first element, has 1s¹ electronic configuration, helium will have 1s² configuration.

Where:

1 = Energy Level

S = S Orbital

¹ / ² = Number of electrons present

Filled up 1s orbital will lead to creation of a new 2s orbital - electrons will move to second energy level.

Lithium will then have 1s² 2s¹ electronic configuration.

Beryllium - 1s² 2s² electronic configuration.

Again, filled up 2s orbital will make electrons to climb to 2p orbitals. 

Boron will then have 1s² 2s² 2px¹.

Carbon = 1s² 2s² 2px¹ 2py¹

Why did 2px¹ not become 2px² before electrons enter 2py¹?

Answer= Hund's rule

He said= electrons fill sub-orbitals (of an orbital) one by one first. After all sub-orbitals have been served with one electron round, then the electrons can start pairing to make complete pairs.

Don't forget that s orbital shape appears spherical while each p orbital appears like figure 8.

Now that you have understood concepts of hybridization and atomic orbitals, we can properly enter Hybridization in chemistry

Hybridization in chemistry is the mixing together of two (or more) atomic orbitals (like mixing a s orbital and p orbital[s]) to have new orbitals that are identical (in terms of many things which I will discuss everything below).

That sounds cringy right? Why on earth would a s orbital mix with p orbitals? I mean why? I need you to ask me that question why?

My answer: Yeah! Atomic Orbitals have to mix because it is the only way some feasible chemical bonding can occur (especially covalent bonds).

Let me explain why atomic orbitals hybridise better for you.

Using carbon as an example: we have just seen that carbon has electronic configuration of 1s² 2s² 2px¹ 2py¹ 2pz⁰ in its ground state.

To achieve an octet structure (=a stable state, having total of 8 electrons in energy level 2) carbon needs 4 more electrons.

However, only unpaired electrons can bond - (valence bond theory).

This makes it that the two paired electrons (in the carbon atom) occupying the 2s orbital (i.e 2s² electrons) must become unpaired so there will be room for bonding. 

Since the energy gap (=energy required for electrons to transfer) between the 2s and 2p orbitals is very small, one of the 2s electrons is transferred (we call it 'promoted') to the empty 2p orbital (which is specifically the 2pz orbital which had no electron).

PRO TIP: Energy gap between s and p orbitals of an energy level, say 2s and 2p, are small which is why electrons can be promoted from 2s to 2p. It'll be hard for electrons to be promoted from 2s to 3s or 3p because they are not in the same energy level (a lot of energy would need to be expended).

This promotion will cause the new electronic configuration of carbon to be 1s² 2s¹ 2px¹ 2py¹ 2pz¹.

PRO TIP: Bonding take place by filling up all half filled orbitals (orbitals with only one electron) to give rise to full filled orbitals (which now have two electrons).

Now the four electrons (those ones in 2s¹, 2px¹, 2py¹, 2pz¹) appear to be ready for bonding, but there is a problem! A very big one!

The 2px, 2py and 2pz orbitals (like I said earlier) are oriented at right angles to each other. 

If bonding occur in this state, these 3 orbitals would form 3 equal bonds oriented at 90° to each other.

But the electron in the 2s orbital would form a bond that is of a totally different type and orientation from the other three. 

No such compound exists!!

This is where hybridization has to enter! 

The first step (that hybridization would do) is to take either 2, 3, or all four of these orbitals (the 2s and 2p orbitals) and intermix them = to equalize their energies!!

Let’s say that we take all four orbitals and mix them, we get 4 equivalent new orbitals.

These new orbitals will be of the same energy, which is intermediate between the energy of the original 2s and 2p orbitals.

We cannot name the new orbitals s or p, because they’re neither. 

We have to find a new name that reflects the fact that the new orbitals were created from one s orbital and three p orbitals. 

We will call them sp³ orbitals. The process that leads to their formation is called sp³ hybridization.

I want to believe you're now getting the concept of hybridization in chemistry, if you still don't, start reading this post right from the top again. I've broken it down well enough.

In summary, Atomic orbitals can be combined and reshaped –much like dough– to make other orbitals of different shapes and properties, so that chemical bonding can occur.

New hybrids formed by hybridization

There are two basic forms of new orbitals that can result from hybridization. The new orbitals can be:

HYBRID ORBITALS: They result from combinations of different atomic orbitals within one atom especially before a reaction takes place.

MOLECULAR ORBITALS: They result from combinations of atomic orbitals between (2) atoms as bonding takes place to form molecules.

Rules of Hybridization in Atomic Orbitals

The atomic orbitals that can form a hybrid must be found in roughly the same energy level

s, p, or d orbitals that mix together to give rise to new orbitals must be found within the same energy level (say 1, 2, 3, 4...). 

This means that only 2s and 2p orbitals can mix together, 3s and 3p orbitals can mix together. 

2s orbital cannot mix with 3p or 4p or 4s orbital it's an abomination.

It is not compulsory for all atomic orbitals in an energy level to combine

Yes! This happens a lot of times. It is not a must that a s orbital must combine with all three p orbitals or all available d orbitals. This is usually dependent on other factors which you're going to understand subsequently in this post.

The number of Hybrid Atomic orbitals formed is always equal to the number of simple Atomic orbitals that combined to form it

If we say a s orbital mixed with two p orbitals, 3 new hybrid orbitals will be formed.

They basically follow arithmetic principles:

1s + 1p = 2 (sp hybrid orbitals)

1s + 2p = 3 (sp² hybrid orbitals)

1s + 3p = 4 (sp³ hybrid orbitals)

Sigma bonds (σ-bond) and pi bonds (π-bond)

Since the beginning of this post we've just been discussing Hybridization! Hybridization!! Hybridization!!! What does hybridization lead to, or why do orbitals of atoms need to hybridize? 

Did I hear someone ask me that question?

Well, atomic orbitals hybridize for one purpose = formation of covalent bonds.

Covalent bonds are formed when two orbitals overlap (like they join together - but not actually joining together, they only share electrons to complete their half pairs). 

Covalent bonds are formed by:

• Overlap of an atomic orbital with another atomic orbital (example: bonds in H2).
• Overlap of a hybrid orbital with an atomic orbital.
• Overlap of a hybrid orbital with a hybrid orbital.

You will understand these better by the end of this post.

Now, I will be using only simple terms here; when orbitals (from different atoms) overlap (to form covalent bonds), they either form a sigma bond or a pi bond - 

A sigma bond is the simplest, most common, most compulsory and first type of bond which 'two different atoms' that combine must carry out.

When two different atoms form a bond (we're talking about covalent bond), they form a bond because they are compatible - they both need to complete their half filled orbitals.

Each atom donates one electron (their unpaired electrons) to give rise to two electrons (paired electrons) making the atoms overlap as a result. This type of bond is a sigma bond. Very easy.

A pi bond on the other hand walks alongside sigma bonds, this means a pi bond cannot happen without a sigma bond. 

In this sense, what result into a pi bond?

I said earlier that it is not compulsory for all orbitals in an energy level to combine (or hybridize). Orbitals that did not take part in hybridization are the ones responsible for forming pi bonds.

In other words, pi bonds occur because of presence of unhybridized orbital(s).

If hybridization take place and there's no unhybridized p orbital, all bonds in the molecule that will be formed will be sigma bonds, but if there is/are leftover p orbitals, those leftovers will form pi bonds.

PRO TIP: It is necessary to know that a pi bond rarely take place alone, it must go along with a sigma bond. Pi bonds are weaker than sigma bonds but a combination of sigma and pi bonds is always stronger than a single sigma bond.

PRO TIP: In organic chemistry, you would need this: All single bonds are sigma bonds. All double bonds comprise of one sigma bond and one pi bond. All triple bonds comprise of one sigma bond and two pi bonds.

PRO TIP: Double bonds are formed because there's presence of one unhybridized p orbital, triple bonds are formed because there are presence of two unhybridized p orbitals.

You can also check out this short table for differences between sigma bonds and pi bonds.

Differences between sigma bonds and pi bonds

Sigma Bond	Pi Bond
Pure or hybrid orbitals can form sigma bonds	Only unhybridized orbitals can form pi bonds
Sigma bonds have high and strong energies	Pi bonds are relatively weak
Sigma bonds can exist independently	Pi bonds must exist along with a sigma bond
Shapes of molecules are determined by sigma bonds along with other factors (explained in this post)	Pi bonds have no impact on shapes of molecules
Two atoms linked by a Sigma bond can rotate freely about the bond	Two atoms linked by a Pi bond cannot rotate freely

Types of Hybridization

sp³ hybridization

In sp³ Hybridization, s orbital mixes with the 3 p orbitals {px orbital, py orbital and pz orbital} to give rise to 4 identical orbitals (called sp³ hybrid orbitals). 

Each of these 4 identical orbitals have 25% resemblance to s orbital and 75% resemblance to p orbital. 

When sp³ hybridization take place to yield 4 sp³ hybrid orbitals, each hybrid orbital is directed towards the 4 corners of a regular tetrahedron (a 3-D triangle).

Regular tetrahedron | photo credit: Wikimedia commons

The angle between each sp³ hybrid orbitals is approximately 109.5°.

Basically, Every sp³ hybridised atom can form only sigma bonds (=because there is no unhybridized p orbital).

Example of compounds that have sp³ hybridization are:

• Methane (CH4)
• Ethane (C2H6)
• All Alkanes
• Water (H2O)
• Ammonia (NH3).

Methane (CH4)

During the formation of methane molecule, the carbon atom undergoes sp³ hybridization prior to bonding.

Each of these sp³ hybrid orbitals forms a bond with one hydrogen atom. 

Thus, carbon forms four sigma bonds (single bonds) with four hydrogen atoms.

Methane molecule is tetrahedral in shape with 109.5° bond angle. Reason for this will be explained in shape of molecules.

Molecular structure of methane

Ethane (C2H6)

Just like in methane molecule, two carbon atoms undergo sp³ hybridization in their excited states.

The two carbon atoms form a sigma 'sp³ to sp³' bond with each other due to overlapping of sp³ hybrid orbitals from each. 

Each carbon atom also forms three sigma bonds with other hydrogen atoms.

Thus there is tetrahedral symmetry around each carbon.

Molecular structure of ethane

Water molecule (H2O)

The electronic configuration of oxygen is 1s² 2s² 2px² 2py¹ 2pz¹.

There are two unpaired electrons in oxygen atom (on 2py and 2pz), which could form bonds with hydrogen atoms. 

In this case, the bond angles in the resulting molecule will be equal to 90°, totally different from the rest

(Remember I said from the beginning, bonding does not occur in this manner!)

Experimental reports showed that bond angles in a water molecule were not equal to 90°, instead 104° was recorded.

Chemists then explained why. They said "oxygen undergoes sp³ hybridization before the bond".

They explained that “during the formation of water molecule, the oxygen atom undergoes sp³ hybridization by mixing its '2s' orbital and its three '2p' orbitals to give rise to four sp³ hybrid orbitals (that have a tetrahedral shape like normal sp³ hybrid orbitals).

Out of these four hybrid orbitals, two are half filled and the remaining two are completely filled (they needed no extra electron).

Now, the two 'half filed hybrid orbitals' form two sigma bonds with 2 hydrogen atoms to yield water (hence they form bond pair) while the two fully filled hybrid orbitals remained the same.

That was accepted, okay fine, but there's still one thing, the reported bond angle was 104° instead of regular tetrahedral angle (which is 109.5°) or linear angle that should have been formed. 

Why?

Chemists defended again and said that this is due to repulsions caused by the two lone pairs (the ones that were completely filled) on the bond pairs (the ones that took in hydrogen to fill them). 

Thus, water molecule take the form of an angular shape (V shape). This will be explained properly in shapes of molecules using VSEPR theories.

Molecular structure of water

Ammonia (NH3)

The ground state electronic configuration of nitrogen atom is: 1s² 2s² 2px¹ 2py¹ 2pz¹

This gives us another scenario like water.

Experiments reported that the bond angle in ammonia is 107°.

It was therefore proposed that; the Nitrogen atom undergoes sp³ hybridization.

Among the four hybrid orbitals, three are half filled and one is fully filled.

The three half filled hybrid orbitals form 3 sigma bonds with three hydrogen atoms.

The observed decrease in the bond angle is again due to the repulsion caused by the lone pair (full filled hybrid orbital) on the bond pairs(the half filled hybrid orbitals that bond with hydrogen) - to be explained better in shapes of molecules below.

Molecular structure of ammonia

sp² hybridization

In sp² hybridization, one s orbital mixes with two p orbitals (px and py) to give rise to three identical orbitals (called sp² hybrid orbitals). 

Each of these 3 identical orbitals take 33.3% resemblance to s orbital and 66.7% resemblance to p orbital. 

sp² hybridization can also be called trigonal hybridization.

All three sp² hybrid orbitals maintain an angle of 120° apart from each other.

When an atom undergoes sp² hybridization and attaches its 3 hybrid orbitals to other atoms, a trigonal planar shaped molecule is formed.

PRO TIP: trigonal planar shape is different from trigonal pyramidal shape, reason for this will Explained better later in this post.

Examples of sp² Hybridization are:

• All the compounds of Boron (BF3, BH3)
• All the compounds of carbon containing a carbon-carbon double bond e.g: Ethene (C2H4), all Alkenes.

Boron trichloride (BCl3)

The electronic configuration of Boron in ground state is 1s² 2s² 2px¹ 2py⁰ 2pz⁰.

It has only one unpaired electron and two empty orbitals.

For maximum bonding, just like carbon, one electron is promoted from 2s² into 2py to form an excited state of 1s² 2s¹ 2px¹ 2py¹ 2pz⁰

Good! But Remember bonding cannot take place in this form?

Then sp² hybridization takes place. 

The 2s orbital mixes with 2px and 2py orbitals to give three sp² hybrid orbitals. 2pz⁰ is not needed at all.

Remember these three hybrid orbitals are half filled and they are oriented in trigonal planar symmetry.

In this form, Boron can now react.

The three half filled hybrid orbitals form three sigma bonds with three chlorine atoms.

Thus the shape of BCl3 is trigonal planar with bond angles equal to 120°.

Trigonal planar structure in boron trichloride

Ethene (C2H4)

During the formation of ethene (or ethylene) molecule, carbon atom undergoes sp² hybridization in its excited state by mixing its 2s with only two 2p orbitals (2px and 2py orbitals) to give three half filled sp² hybrid orbitals oriented in trigonal planar symmetry.

There is also one half filled unhybridized 2pz orbital on the carbon atom which does not follow the same shape as the hybridized orbitals because it is completely different from them.

Two carbon atoms undergo this process (of sp² Hybridization) and bond together to form a sigma bond with each other by using one of their sp² hybrid orbitals each.

At the same time, a pi bond is also formed between them due to lateral overlapping (side by side overlapping) of unhybridized 2pz orbitals.

Thus there is a double bond (sigma and pi bonds) between the two carbon atoms.

Each carbon atom also forms two sigma bonds with two hydrogen atoms.

Thus ethene molecule is planar the way you know it as.

Trigonal planar structure in ethene

sp Hybridization

In sp hybridization, one s orbital mixes with one p orbital (px) and give rise to two identical orbitals (called sp hybrid orbitals). 

Each of these 2 identical orbitals take 50% resemblance to s orbital and 50% resemblance to p orbital.

sp hybridization is also called diagonal hybridization.

It forms linear molecules with an angle of 180°.

Examples of sp Hybridization:

• All compounds of beryllium (e.g: BeF2, BeH2, BeCl2).
• All compounds of carbon, containing triple Bonds like ethyne (C2H2), infact all alkynes.

Beryllium Chloride (BeCl2)

The electronic configuration of Beryllium in ground state is 1s² 2s² 2px⁰ 2py⁰ 2pz⁰

2px, 2py and 2pz orbitals have 0 electrons.

Now, since there are no unpaired electrons, there's lesser likeliness for bonding to take place; 

This makes Beryllium undergo excitation by promoting one of its 2s electron into an empty 2p orbital (let's say 2px). 

Thus in the excited state, the electronic configuration of Beryllium is 1s² 2s¹ 2px¹ 2py⁰ 2pz⁰

There's still a problem because this atom cannot yet form a bond like this, therefore, Beryllium mixes (hybridizes) its '2s' orbital with the only available '2p' orbital to give rise to two sp hybrid orbitals.

The new hybrid orbitals are still half filled and they are arranged linearly (like along a straight line) - reason for this will be explained below.

These half filled sp orbitals form two sigma bonds with two 'Cl' atoms.

Thus BeCl2 is linear in shape with the bond angle of 180°.

Linear Structure of beryllium chloride

The Same exact thing happens to most Beryllium compounds.

Ethyne (C2H2)

During hybridization here, carbon only mixes its '2s' orbital with only one '2p' orbital to give rise to two 'half filled sp orbitals' which are arranged linearly - 

Two carbon atoms undergo the above process and then they form a sigma bond with each other by using (one-one of their) sp hybrid orbitals.

However there are also two unhybridized orbitals (2py and 2pz) on each carbon atom.

These unhybridized orbitals are not in the same orientation as the sp hybrid orbitals. 

These orbitals form two pi bonds between one another during bonding of the two carbon atoms.

Thus a triple bond (comprising of one sigma bond and two pi bonds ) is formed between carbon atoms.

Then each carbon also forms a sigma bond with hydrogen atoms.

Thus ethyne molecule is linear with 180° of bond angle.

Structure of ethyne

PRO TIP: Ethyne is also known as acetylene, if you see acetylene perhaps in your exams or anywhere do not let it confuse you.

sp³, sp², sp hybrid orbitals are the basic types of hybrid orbitals.

There are other complex ones such as sp³d, sp³d², sp³d³, sp³d⁴ and sp³d⁵ orbitals.

sp³d hybridization

sp³d hybridization involves the mixing of one 's', three 'p' (px, py and pz) orbitals and one 'd' orbital to form 5 sp³d hybrid orbitals of equal energy. 

They have trigonal bipyramidal geometry.

Three of these hybrid orbitals lie in the horizontal plane inclined at an angle of 120° to each other known as the equatorial orbitals while the remaining two orbitals lie in the vertical plane of the equatorial orbitals (known as axial orbitals).

Example:

• Hybridization in Phosphorus(V) chloride (PCl5).

sp³d² Hybridization

sp³d² hybridization has one 's', three 'p' (px, py and pz) and two 'd' orbitals, that undergo intermixing to form 6 identical sp³d² hybrid orbitals.

These 6 orbitals are directed towards the corners of an octahedron. Hence they have an octahedral orientation.

They are inclined at an angle of 90° to one another.

Example: 

• Sulphur hexafluoride (SF6).

sp³d³ Hybridization

A perfect example of this type of Hybridization is iodine heptafluoride (IF7)

What happens here?

The electronic configuration of Iodine atom in ground state is: [Kr] 4d¹⁰ 5s² 5px² 5py² 5pz¹

Since the formation of IF7 requires 7 unpaired electrons in iodine, the iodine atom promotes three of its electrons (one from 5s orbital and two from 5p sublevel) into empty 5d orbitals. 

This will form electronic configuration of 

[Kr] 4d¹⁰ 5s¹ 5px¹ 5py¹ 5pz¹ 5dx² - y²¹ 5dz²¹ 5dxy¹

This state is referred to as third excited state. 

The essence of this promotion is to basically make up 7 unpaired electrons (so to make reaction with fluorine feasible).

But then, even Iodine cannot react after electron promotion. It needs to hybridise its orbitals.

Iodine then undergoes sp³d³ hybridization to give rise to 7 half filled sp³d³ hybrid orbitals.

The shape of these seven sp³d³ hybrid orbitals form pentagonal bipyramidal symmetry. 

These 7 hybrid orbitals will then go ahead to form 7 sigma bonds with fluorine atoms.

Other hybridization patterns like sp²d, sp³d⁴, sp³d⁵ - explained in higher classes.

This marks the end of types of Hybridization.

Hybridization of Carbon Compounds

We have discussed this several times so this will be brief:

Carbon's ground state electronic configuration is 1s² 2s² 2px¹ 2py¹ 2pz⁰

Carbon's excited state electronic configuration is 1s² 2s¹ 2px¹ 2py¹ 2pz¹

This makes carbon ready for bonding, but bonding is not feasible. This will lead to hybridization.

Carbon is a very flexible element and it can hybridise its atomic orbitals anyhow it likes (especially the exact way it will be needed for bond formation).

Carbon can undergo various Hybridization to suit the kind of molecule it wants to make

sp³ hybridization in Carbon

When carbon undergoes sp³ hybridization, it mixes its '2s' orbital with '2px, 2py and 2pz' orbitals and this will give rise to four new hybrid orbitals (specifically called sp³ orbitals).

Its electronic configuration will now be looking like 1s²   2(sp³)¹   2(sp³)¹   2(sp³)¹   2(sp³)¹ 

where 2sp³ orbitals are hybrids and are identical.

Each of these sp³ orbitals are half filled. They all need one more electron to become stable.

Therefore Hydrogen (which has 'just the one electron' required) comes to bind to every of these sp³ hybrid orbitals. That is 4 Hydrogen atoms bind to carbon.

The resultant bonds are 4 sigma bonds (single bonds) and the shape of the molecule is tetrahedral.

All Carbon atoms in every alkane exhibit sp³ Hybridization.

sp² hybridization in Carbon 

An example is ethene that has been discussed above. Carbon mixes its '2s' orbital with only '2px and 2py' orbitals giving rise to three new hybrid orbitals (specifically called sp² hybrid orbitals).

Its electronic configuration will now be looking like 1s²   (2sp²)¹   (2sp²)¹   (2sp²)¹   2pz¹

Each of these sp² orbitals are half filled. 

Reactions will lead to formation of sigma bonds by two sp² hybrid orbitals; and the last sp² hybrid orbital, although it also forms a sigma bond, will be complemented with a pi bond as a result of the leftover 2pz orbital. 

Therefore, sp² hybridization in carbon will lead to formation of two single bonds and one double bond.

All carbon atoms in Alkenes exhibit sp² hybridization.

sp hybridization in Carbon 

An example is Ethyne that has been discussed above. Carbon mixes its '2s' orbital with only '2px' orbital giving rise to two new hybrid orbitals (specifically called sp hybrid orbitals).

Its electronic configuration will now be looking like 1s²   (2sp)¹   (2sp)¹   2py¹   2pz¹

Each of these sp orbitals are half filled. 

Bonding with the sp orbitals will lead to formation of sigma bonds. However, one sp hybrid orbital will be complemented by two pi bonds as a result of the leftovers: 2py and 2pz orbitals.

Therefore, sp hybridization in one carbon atom will lead to formation of one single bond and one triple bond.

Shapes of simple molecules

Do I need to start by defining shape for you?

No! if you cannot explain shape, authorities are coming to pursue you back to your father's house😏  -

Molecules take the form of various shapes. 

Why?

Molecules take the form of various shapes because of 'the action of paired electrons on one another' - molecules don't just assume any shape because they feel like it's cool.

Let me explain this better.

For example, when drawing methane molecule back in your secondary school, you will draw it like this right?

But in reality, the methane molecule shape is not actually like this: 

The paired electrons (illustrated in the above structure with dash) are actively repelling (or pushing away) one another, and in order to maintain maximum repulsion, they take the form of a 3-D shape, not 2-D shape as drawn above.

3-D shape in the sense that while one is oriented towards the front direction (or front of a plane), another is oriented towards the left direction, then the third is oriented towards the back direction (of that same plane) while the last one is oriented towards the right direction. 

All these paired electrons are trying their best to repel one another to the max level, hence, they are all pushing others further apart from themselves.

This is called electron pair repulsion (recall from electricity: like charges repel).

Therefore it is said that molecules take shape because of electron pair repulsion - 

PRO TIP: every covalent molecule is composed of completely paired electrons in valence shells, hence one electron pair is responsible for actively repelling another electron pair.

Yeah! Well that is just, but the first step actually.

Molecules can take another form of shape if there's presence of lone pair of electrons (lone pair of electrons are electrons that did not take part in bonding with other atoms because they are completely filled or paired already).

A short explanation I developed and I want you to note is that when an atom forms bonds, Lone pair of electrons do not project outwards like bond pairs - but the lone pairs are actively repelling the bond pairs

I'm guessing this looks complex already so let me break it down.

Take methane for example, you see that carbon has four bond pairs (electrons that bonded with hydrogen). Thus all four project upwards out of the central carbon atom and repel one another to take the form of 3-D shape, the shape of a tetrahedral molecule. 

However, if you take ammonia into consideration;

Nitrogen, which is the central atom undergo same process as carbon in methane; sp³ hybridization to form 4 identical orbitals.

But nitrogen contains one lone pair of electrons (full filled hybrid orbital) and three bonding electrons (half filled hybrid orbitals) that require just one more electron each.

Hence: when hydrogen atoms come to attach to these three half filled hybrid orbitals, they form bond pairs, and hydrogen atoms project out of the central nitrogen atom like this:

Now three bond pairs have been formed, and there's one lone pair (on the nitrogen atom) which will still be actively repelling the bond pairs. 

Like I said, lone pairs don't project outwards in space like bond pairs, hence the shape of ammonia molecule will be looking like in this image below: trigonal pyramidal - pyramidal because the projections not evenly distributed.

If not for the action of the lone pair of electrons on the bond pairs, an ammonia molecule would have been looking like a trigonal planar.

Now, if we take water into consideration:

Oxygen is the central atom, and it undergoes sp³ hybridisation to form four identical sp³ hybrid orbitals.

But out of these four hybrid orbitals, two are paired already, they don't need extra electron (= lone pairs) and two are not paired, hence they will need to form bonds. 

Valence electrons in oxygen atom: two 'already paired electrons' at the top and left ; and two 'unpaired electrons' at the bottom and right.

Two hydrogen atoms come to fill these two hybrid orbitals that have one electron each and at the end of the day they finally become paired. With hydrogen atoms projecting outwards from the central oxygen atom.

Like I said, lone pairs do not project outwards like bond pairs, hence the two lone pairs will be there, actively repelling, but they do not show face outwards. 

The other two bond pairs will project outwards and they are actively repelled and also repelling, hence only two outward projections, which accounts for the V structure in water molecule.

I want to believe you are really getting this explanation on shapes of molecules. I've broken it down easy enough so go back and read if you do not get it, or take a break and come back later when your brain becomes cool.

Okay!

So, We've been discussing shapes of molecules gotten from sp³ hybridized atoms, what about sp² and sp hybridized atoms?

When a Central atom undergoes sp² hybridization, we know the shape of the 3 new hybrid orbitals is going to be trigonal planar right? 

That's right!

If that atom forms bonds with all its sp² hybrid orbitals, the resulting molecule will also be trigonal planar. 

But if there's one lone pair that doesn't need bonding amongst the three sp² hybrid orbitals, the resulting molecule will not form a trigonal shape but an angular shape with two projections only. I believe you got that. Lone pairs do not project outwards! But they are actively repelling!

Also, for atoms that undergo sp hybridisation, we know the shape of the 2 new hybrid orbitals is linear.

That's right.

They (the electrons) tried to separate as far as possible.

If that atom forms bonds with its two hybrid orbitals, the resulting molecule will also be linear. 

Let me be sarcastic here LOL (this is definitely not going to happen in covalent bonding) if we assume that one of the two hybrid orbitals is already fully filled (2 electrons inside it hence it doesn't need to form bond = lone pair) and the other is half filled (has only one electron), the half filled orbital forms bond to become fully filled (it becomes a bond pair). The lone pair would not project outwards at all, but the bond pair will project, hence the molecule will look like it's hanging from the central atom (that's funny) but take note that the lone pair will still be actively repelling the bond pair, it won't just project like the bond pair - lone pairs look like they aren't there, but they're acting, repelling other pairs - 😖 

I want to believe you got that.

You can use these explanation to easily predict the shape of any molecule.

All that I've just explained above is a simple summary of the VSEPR Theory or VSEPR model - pronounced as |versper| its full meaning is Valence Shell Electron Pair Repulsion.

I re-translated it in my own words to be 'repulsion of electron pairs in valence shells'.

VSEPR theory has four assumptions:

• Electrons should only be placed in pairs (2-2) in the valence shell (or outer shell) of an atom.
• Electrons that will form bonds (=they are not paired yet) and those that are already paired (=they won't form bonds) are included.
• All Electron pairs repel each other - maximum separation. 
• Non-bonding pairs (or lone pairs) repel more strongly than bonding pairs, because the Non-bonding pairs are attracted to only one nucleus.

Prediction of the shape of a molecule is easily done by using VSEPR theory. 

You should know that VSEPR theory has so many limitations - you will see that in higher classes.

It is also important to note that while VSEPR theory can be used to predict the shape of a molecule, the actual shape of any solid molecule, or molecules, is gotten by X-ray diffraction- there are several other methods but X-ray diffraction is the most commonly used. On the other hand, the actual shape of gaseous molecules is gotten by electron diffraction.

While VSEPR theory can sometimes be difficult to apply directly, especially by learners like you, a simpler formula was created, AX_mE_n, which serves as a basis for the application of the VSEPR theory, for shape prediction

Where, A = Central atom

X = atoms that bond to the central atom

m = number of X

E = Lone pair of electrons available in the central atom,

n = number of E

m + n gives the steric number or electron domain of the molecule. Steric number refers to how many paired electrons that the central atom of a molecule can have.

All of these numbers can be quickly interpreted to get the shape of a molecule, by using the data in the table provided later below.

However, because you're at your learning stage, you cannot really understand how to use this formula or table without knowing how to draw lewis structures for molecules; so let us take a quick look into that.

What is Lewis Structure?

Lewis structure is also known as Lewis dot structure. 

It makes use of dots to represent electrons in the valence shell of an atom. 

Electrons in the valence shell of an atom are electrons in the outer most shell - outer energy level.

They are the ones that take part in bonding that we have been talking about. If you do not have basic knowledge of electronic arrangements in an atom, you will find this hard to understand, or as if I'm blabbing. 

Every atomic orbital or shell can only take 2 electrons - 

PRO TIP: you should know that energy levels or principal quantum numbers take the summation of all electrons in its various orbitals or shells (we use the formula 2n² for Maximum electron containable but that doesn't concern us here for now)

The first energy level has just a s orbital, the first two electrons fill up this s orbital, the remaining electrons can then occupy the second energy level and beyond. 

Remember from our discussion, the second energy level has s, px, py and pz orbitals so it will take a maximum of 8 electrons because each of these orbitals will take two electrons. 

You should note that the third energy level has s, px, py, pz, and five d orbitals - dz², dx²-y², dxy, dyz, dxz These d orbitals are only filled up in d block elements - we're coming there.

Lewis dot does not take orbitals into consideration, it just arranges electrons into four places around an atom. 

Take carbon for example: carbon has an atomic number of 6. The first energy level takes the first 2 electrons, so the electrons that will be in the second/outer energy level (valence shell) are 4.

PRO TIP: atomic numbers of elements is their number on the periodic table. So when you count the first twenty or thirty elements, the number where you reached an element, that number is also its atomic number - and the atomic number of an element is its number of electrons (and also protons) provided it's in its ground state.

Lewis dot will take these 4 electrons and arrange them with four dots around carbon like this:

Note that Lewis dot tries to obey Hund's rule. Which makes it put the dots one - one first. When the four locations are filled, it then puts the leftover electrons (if available) to pair with the first 'one - one'. But, In normal electronic configurations, the s orbital gets filled first

After Lewis dot has been drawn on carbon, if carbon undergo reaction with hydrogen, hydrogen electrons forms bonds with the single (unpaired) electrons on carbon. So, the structure becomes:

But standards has it that bonds are represented with dash. So it becomes:

You see there are four dashes - indicating four bonds.

The only limitation to Lewis dot is that it lets you view molecules in 2-D format. To view this in 3-D and know the exact shape, you will have to apply VSEPR theories which I've described above. 

Repulsions between these four electron pairs will lead to formation of angles, Hence it can be predicted that the structure of methane (using VSEPR theory) will be tetrahedral.

Also, if we use the formula, AXmEn,

Central atom (A) = Carbon

Bond atoms (X) = Hydrogen only

Number of bond atoms (m) = 4

Number of Lone pair of electrons in carbon (En) = 0

We therefore, have AX4E0

Steric number = m + n

                          = 4 + 0

                          = 4

A steric number of 4, and 0 lone pair of electrons = tetrahedral shape. (Take a look at the table image below).

Let us use Ammonia:

Nitrogen is on number 7 in the periodic table. I was supposed to let you count it yourself.

Hence Nitrogen has 7 electrons

The first energy level will take 2 out of these 7 electrons, and 5 electrons will be left.

The 5 electrons are available in the valence (outer) shell and will be drawn with Lewis dots as:

As you can see, there's already one paired side in nitrogen, this paired side does not need an extra electron, but the remaining three sides need one electron each. Hydrogen atoms come to attach to these unpaired electrons to form:

As you can see, three bonds with hydrogen and one lone pair. Remember I said lone pairs don't project, but they are repelling the other bonds, therefore the structure of ammonia (using VSEPR theory) is trigonal pyramidal.

Also, if we use the formula, AXmEn,

Central atom (A) = Nitrogen

Bond atoms (X) = Hydrogen only

Number of bond atoms (m) = 3

Number of Lone pair of electrons in Nitrogen (En) = 1

We therefore, have AX3E1

Steric number = m + n

                          = 3 + 1

                          = 4

A steric number of 4, and 1 lone pair of electrons = trigonal pyramidal shape.

Let us now use water:

Oxygen in the periodic table is in number ( ) fill that up yourself. You are a big OLODO if you do not know the answer to that in 2-5secs.

This will make the electrons of oxygen ( ).

The first energy level takes 2 electrons, hence 6 electrons will be left in the valence shell.

The 6 electrons available in the valence shell will be drawn with Lewis dots as:

As you can see, there are two already paired electrons. These paired sides do not need extra electron, but the remaining 2 unpaired sides need one electron each. 

Hydrogen atoms come to attach to these 2 unpaired sides to form:

As you see, we can deduce (from VSEPR theory) that the two lone pairs will keep repelling the two bond pairs - which are also repelling each other, hence, there will be angular shape of the water molecular structure not linear shape or tetrahedral.

Also, if we use the formula, AXmEn,

Central atom (A) = Oxygen

Bond atoms (X) = Hydrogen only

Number of bond atoms (m) = 2

Number of Lone pair of electrons in Oxygen (En) = 2

We therefore, have AX2E2

Steric number = m + n

                          = 2 + 2

                          = 4

A steric number of 4, and 2 lone pair of electrons = Angular shape.

I want to believe that you now understand the concept of Lewis structure.

From Lewis structure, you can deduce the shape of a molecule by applying the VSEPR theory.

Below is the table showing:

• number of atoms that bond with central atom (or number of bond pairs), (Xm)

• number of lone pairs, (En)

• electron domains, or steric number = number of bond atoms (m) + number of lone pairs (n), 

• shape of molecule

• ideal bond angles 

• and examples.

Courtesy: Dept of Chemistry, Anugrah college, Gaya

If I were you. I would re-position myself (watch the recommended videos) and read this post again right from the beginning. 

It's pretty lengthy I know, but you would finally get to understand hybridization and shapes of molecules very much better.

Once you're able to understand this lecture, basic hybridization principles in chemical bonding: done and dusted😉

Key Points

💡Hybridization is the mixing of atomic orbitals to give rise to hybrid orbitals that can form bonds.

💡Only atomic orbitals in the same energy level can hybridize.

💡The number of hybrid orbitals formed is equal to the number of atomic orbitals that mixed.

💡It is not necessary that all atomic orbitals in an energy level must participate in hybridization.

💡Hybridization happens only during or prior to bond formation and not in isolated state.

💡The shape of a molecule can be predicted by hybridization of its central atom(s).

💡The bigger lobe of a hybrid orbital always has a positive sign while the smaller lobe on the opposite side has a negative sign.

💡Single bonds are sigma bonds.

💡Double bonds contain one sigma bond and one pi bond.

💡Triple bonds contain one sigma bond and two pi bonds.

💡Pi bonds occur because of presence of unhybridized half filled orbital(s).

💡When one s orbital mixes with three p orbitals, sp³ hybridization is said to happen.

💡The structure of the new hybrid orbitals gotten from sp³ hybridisation is tetrahedral.

💡The shape of molecule formed by an atom that undergo sp³ hybridisation is tetrahedral - with condition that there is no lone pair.

💡If there's presence of a lone pair, the shape of molecule that will be formed by such sp³ hybridised atom will be trigonal pyramidal.

💡If there's presence of two lone pairs, the shape of molecule that will be formed by such sp³ hybridised atom will be angular or V shaped.

💡When one s orbital mixes with two p orbitals, sp² hybridization is said to happen.

💡Sigma and pi bonds are present in sp² hybrid molecules - with condition that there's a leftover half filled p orbital.

💡The structure of new hybrid orbitals formed from sp² Hybridization is trigonal planar.

💡The shape of molecule formed by an atom that undergo sp² hybridisation is trigonal planar  - with condition that there's no lone pair.

💡When one s orbital mixes with one p orbital, sp hybridization is said to happen.

💡The structure of the new hybrid orbitals gotten from sp hybridisation is linear.

💡The shape of molecule formed by an atom that undergo sp hybridisation is linear.

💡VSEPR - Valence Shell Electron Pair Repulsion theories help deduce the shapes of molecules - provided the hybridisation of the central atom is known or Lewis structure of the molecule is known.

Key Terms

Hybridization • Atomic Orbitals • s orbitals • p orbitals • d orbitals • Valence Bond Theory • Energy Gap • Electron Promotion • Hybrid Orbitals • Molecular Orbitals • Sigma (σ) Bond • Pi (π) Bond • sp³ Hybridization • Tetrahedral Structure • Trigonal Pyramidal Strcture • sp³ Hybrid Orbitals • sp² Hybridization • Trigonal Hybridization • Trigonal Planar Structure • sp² Hybrid Orbitals • sp Hybridization • Linear Structure • Trigonal Bipyramidal Structure • Octahedral structure • Pentagonal Bipyramidal Structure • Electron Repulsion • Molecular Geometry • VSEPR Theories • Lewis Dot Structure 

Ensure you watch these videos for Max understanding

Hybridization and Types Explanation

How to determine hybridization (in the central atom) of a molecule - take note that this doesn't always work in advanced chemistry: that's why you need to read through the basics in the main post

Shapes and molecular geometry of simple molecules

Check how well you understand this topic

Discuss and Explain

1. Explain Hybrid Orbitals.

2. Explain why Hybrid Orbitals are better than their parents orbitals.

3. Explain the difference between sp, sp² and sp³ hybridization.

4. Explain the differences between Molecular Orbitals and Hybrid Orbitals.

5. What is the hybridization in BeCl2?

6. What is the shape of methane molecule?

7. Give two examples of sp³ Hybridization.

8. What is the excited state configuration of Carbon?

9. Differentiate sigma bonds and pi bonds.

10. Provide the ground state electronic configuration and number of valence electrons for carbon and silicon. Describe how these atoms are similar. Which bond do you think is stronger, a C - C bond or Si - Si bond? Explain

11. Draw the Lewis structures for: SiH4, PH3 and H2S and provide the molecular geometry for each as predicted by VSEPR Theory.