Chemistry: How to Determine Hybridization
Maybe you’ve been thinking about molecules (I know I do, on an embarrassingly regular basis). If you’ve been thinking about them, perhaps you are also wondering the very important question of: What do those little things actually look like?
Let’s take water as an example. You can’t see an individual molecule of water, because each water molecule is about 2.75 angstroms (approximately 1/100,000,000 of an inch!). We can certainly see a whole collection of water molecules, though, when we drink a glass of water (which contains approximately 6 septillion molecules of water).
Nonetheless, scientists in general (and chemists in particular) have done a lot of research to understand what molecules look like, and particularly how atoms in a molecule are connected to each other. As a reminder, each atom contains neutrons and protons (found in the nucleus), and electrons, which are located in the area around the nucleus.
Electrons and Bonding
The electrons, even though they are small in size (approximately 1/1836 the size of a proton or neutron), are among the most important parts of molecules, because they are the key players in forming chemical bonds with other atoms. These bonds, or connections between atoms, can occur in one of two ways:
- Transferring an electron from one atom to another, where that transfer causes one atom to become positively charged (the atom that lost its electron), and one atom to become negatively charged (the atom that gained an extra electron). This is called ionic bonding, which we are not going to talk about today.
- Sharing electrons between atoms, in which each atom that is involved in this sharing has some kind of “pull” or attraction for the electrons that are located between them. This gentler approach to bonding is called “covalent bonding,” and it can either be “non-polar covalent,” in situations in which the atoms are pulling electrons equally, or “polar covalent,” when the two atoms in the bond are pulling on the electrons with unequal strengths.
In both types of bonding, the active players are the small but mighty electrons! All the other components are kind of just watching to see what the electrons will do.
Let’s talk more about where the electrons are located in an atom before it forms bonds with other atoms, and how that can sometimes change when bonds start to form.
If we thought it was difficult to see a molecule because of how small it is, we can only imagine that trying to see electrons is even more difficult! Nonetheless, scientists were able to use lots of indirect evidence from experiments to decide that electrons are located in regions of space around the nucleus, which is the central area that contains all of the protons and neutrons. Those regions of space where electrons live are called “orbitals,” and each orbital has enough space to accommodate two electrons (kind of like two people sharing a room in a hotel). When one of the “rooms” in the hotel, or orbitals, fills up, then electrons start to fill up the next “room,” or orbital, until every electron has a home, or a region of space that it is assigned to.
This is all a very happy story for electrons that surround an individual atom (i.e., an atom that is not bonding to any other atom), because these electrons are well-behaved and generally stay in the orbitals that they are assigned to.
It is a different story completely, however, when it comes time to form covalent bonds between that atom and another atom. We have already discussed the fact that covalent bonds occur when two atoms share their electrons, but how does sharing work when the electrons are already assigned to their hotel rooms, or orbitals?
It turns out that for a lot of atoms, their electrons have to undergo a rearrangement before they are ready to bond to other atoms. This occurs because the sharing of electrons between atoms that is the definition of covalent bonding occurs when one atom with a half-filled orbital (i.e., an orbital with one electron) binds to another atom with a half-filled orbital, and then the electrons from each half-filled orbital are able to share.
If the atom is in a situation where it doesn’t have as many half-filled orbitals as the number of bonds it wants to form, WHAT CAN IT DO?
It can either refuse to form bonds (always an option, but not always advisable), or it can mix its orbitals into new set-ups that cause each orbital to be half-filled. This is mixing is called “chemistry hybridization,” or “hybridization” for short.
Chemistry hybridization is the phenomenon of orbital mixing, into new orbitals (kind of like new hotel rooms) that are ready to form covalent bonds to other atoms.
Let’s consider a specific example, and talk about how it hybridizes, or undergoes the process of chemistry hybridization.
Let’s say we want to make methane (chemical formula: CH4).
Fun fact: 37% of the world’s methane comes from cow poop! Ew!
Carbon, whose chemical symbol is C, has four electrons that are available for bonding (the other two electrons that it has are held too closely to the nucleus to be of any use for us). When carbon is not forming bonds, those four electrons are found in three orbitals:
- Two electrons are found in an orbital called “2s.”
- Two electrons are found in two “2p” orbitals, with one electron in each of those orbitals.
- There is another 2p orbital that is empty.
Now carbon wants to make four bonds – one bond to each of the hydrogen atoms (chemical symbol: H). What does it do?
It takes the four electrons that are available for bonding, and the four orbitals that are available (one 2s orbital, and three 2p orbitals), and turns all of them into four identical orbitals. Each of these orbitals has 25% of its character from an s orbital (the 2s) and 75% of its character from a p orbital (the three 2p orbitals). Therefore, we call them, “sp3 hybrid orbitals.” Their energy is between the energy of a 2s orbital (lower in energy) and the energy of a 2p orbital (higher in energy). We can put one electron into each of the sp3 orbitals, and now the carbon’s electrons are all set up to form covalent bonds (i.e, sharing) with four hydrogen atoms: each hydrogen atom will share one electron with a half-filled sp3 hybrid orbital.
Sometimes we have a more complicated situation, like if I have ethylene (chemical formula: C2H4). How does chemistry hybridization work for this molecule?
Fun fact about ethylene: it is a gas that is emitted by fruits as a sign of fruit ripening – when you put a banana in a brown bag with an avocado, the ethylene from the banana will make the avocado ripen faster. One of my students decided that the appropriate nickname for ethylene is in fact “banana breath”!
Each carbon in ethylene wants to form bonds to three other atoms, not four! So it does that by taking its 2s orbital, and two of its 2p orbitals, and turning those into three sp2 hybrid orbitals! Each one of these half-filled orbitals can then bond to other atoms, who also have half-filled orbitals to match.
What about the other 2p orbital that we didn’t use, you may ask? Well, we left an electron in that orbital (kind of like saving it for later). In this case, the electron in the 2p orbital that didn’t hybridize overlaps with its twin, an electron in a 2p orbital on the neighboring carbon which ALSO didn’t bind. This kind of overlap (different from our regular bond sharing that causes covalent bond formation) ends up forming a bond that is called a pi-bond (symbolized with the Greek letter π). This overlap doesn’t involve chemistry hybridization at all, so we will leave this discussion for a different post.
Anyway, that’s all for today from the fascinating world of electron hotels, cow poop, and banana breath. Oh, and chemistry hybridization too.
Author: Mindy Levine, PhD