Ionic and Covalent Bonds, Hydrogen Bonds, van der Waals - 4 types of Chemical Bonds in Biology

There are four types of chemical bonds essential for life to exist: Ionic Bonds, Covalent Bonds, Hydrogen Bonds, and van der Waals interactions. All of these bonds represent an emergent property of atoms. Individually, we can study the makeup of atoms - how many protons, neutrons, and electrons they have.

But when atoms come in to close proximity with each other, a new characteristic becomes apparent. If an atom doesn’t have a filled valence shell of electrons, it can either share electrons with another atom, or it can completely transfer electrons. This way, an atom can gain enough electrons to fill its outermost shell, or shed enough electrons to empty its outermost shell, leaving a complete valence shell underneath at the next lower Energy level.

The interactions between atoms that take place as a result are called chemical bonds. So this is an example of an emergent property. You don’t know anything about an atom’s chemical behavior until it comes in close contact with another atom - then the chemical behavior EMERGES.

One thing that determines what kind of bond an atom will be involved in is its Electronegativity. Think of electronegativity as the degree to which an atom pulls on an electron and holds it close. If two atoms come together and they have fairly similar electronegativities, they will share an electron between the two of them.

This especially makes sense when the two atoms are identical. Like - a hydrogen and a hydrogen share an electron and become an H2 molecule. Or, let’s look at two oxygen atoms.

Neither one has a complete valence shell of electrons. The outermost valence shell in oxygen has 6 electrons. So it needs two electrons to fill up that valence shell and become more stable.

When two oxygen atoms come into close proximity, they can share two electrons and so form a double bond, becoming an O2 molecule. They have the same electronegativity, since they’re identical, so the electrons are equally shared between the atoms. We call this a NONPOLAR COVALENT BOND.

Now, if there IS a difference in the electronegativities, the electrons will spend a greater amount of time next to the atom with a higher Electronegativity. Remember, electrons are always in motion, moving around in kind of cloudlike regions called orbitals. For instance, in a water molecule, H2O,Oxygen has a higher electronegativity than Hydrogen.

So in both of the bonds in this molecule, the electrons will spend a greater part of their time closer to the oxygen atom versus next to the hydrogen atoms. This results in a water molecule having a partial negative charge around the oxygen atom, and a partial positive charge around the hydrogen atoms. We call this kind of chemical interaction a POLAR COVALENT BOND.

Now if the electronegativities are different enough, one atom will actually donate one or more electrons to the more electronegative atom. This results in two charged species - ions. The atom that donates one or more electrons is called a cation.

It becomes positively charged. The atom that takes on electrons becomes negatively charged. This is called an anion.

As a result of becoming oppositely charged, these ions are now attracted to each other through electrostatic attraction. That attraction holding the two ions together is called an IONIC BOND. Now here’s a funny thing about ionic bonds in biology.

When you learn about ionic bonds in chemistry, like say, the bonds holding the salt NaCl together, we think of it as quite a strong bond. But in biology/ biochemistry, you must remember that everything - all these chemical interactions - are taking place in the context of water. In water, ionic bonds quickly dissociate.

So for this reason, in biochemistry, we consider ionic bonds weaker than most covalent bonds. In chemistry, we always discuss the range of strengths of the bonds. You can learn more in a video we have called ionic bonds vs covalent bonds in our chemistry video series.

This may surprise you, but it’s really important that we have the capability to make bonds of varying strengths. There are some cases where we want very strong bonds. Say, when we’re building structures.

But there are other cases where it makes more sense to have weak bonds. Bonds that can be used for REVERSIBLE interactions. For example.

Let’s say we have a receptor, and a signalling molecule that binds it - like a hormone or a neurotransmitter. We don’t want that signalling molecule to bind to the receptor once, and get stuck there, blocking the receptor. We want it to bind REVERSIBLY - to bind, send the signal, and then fall off.

In fact, there are some poisons that work this way - they bind permanently to a receptor and don’t fall off, so they mess up signalling in your body. The next weaker bond is the Hydrogen bond. People often get confused about hydrogen bonds because the one example you’ll always hear of a hydrogen bond is in water.

We use that example because it’s incredibly important for why water is essential to life. . .

and that’s a topic for another video. But meanwhile, the confusion comes because we just said there are polar covalent bonds in water. Yes.

In INDIVIDUAL water molecules, those bonds are polar covalent bonds. There are hydrogen bonds that hold different water molecules together. This is what makes water so cohesive.

For example, why water forms drops that kind of bunch up instead of lying flat. The hydrogen bond is the very weak attraction between one hydrogen atom that is already covalently bonded to something else (so it has a partial positive charge) and something else that is partially negative. This is usually an oxygen, nitrogen, or fluorine atom that is bonded to something else.

So in the case of water, you have the hydrogen from one water molecule attracted to the oxygen in a DIFFERENT water molecule. Notice that these bonds are written differently from covalent bonds. Rather than a solid line joining the two atoms, it’s a dotted line.

That’s to remind you that this is a very weak interaction. These bonds sort of blink on and off. Now, let’s talk about the very weakest kind of bond, the van der Waals interactions.

To understand these, you must keep in mind that idea that electrons are constantly in motion. They occupy an orbital, which a cloud, or a region, around their atom’s nucleus. You can’t pinpoint exactly where it is going to be at any one moment, but you have an idea of this region where you can find it with high probability.

So now imagine two typical bulky molecules coming into contact. That receptor and signalling molecule we mentioned earlier. On the surface, the very very surface, there’s a cloud of electrons moving around.

Some of the time, there will be a temporary uneven distribution of electrons, which results in partial positive and partial negative regions. The opposite regions of the receptor and binding partner will line up, and briefly be attached to each other. And then in the next instant, they are released.

Again,sort of like the hydrogen bonds, imagine these bonds twinkling on and off like christmas lights, except now even faster. Here’s a challenge for you. For each one of these kinds of bonds, give us an example of it playing an important role in biology.

Write your answers in the comments! This is a good strategy to adopt as you go forward in biochemistry. If you can keep an example molecule in mind, that can help keep the more abstract concepts clear.

So - this is what a hydrogen bond looks like. This is what a fatty acid looks like. You won’t get so easily confused later on.