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  • 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 theyre 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 were 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 youll 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.

  • Thanks for watching Socratica

There are four types of chemical bonds essential for life to exist: Ionic Bonds, Covalent Bonds,

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