IFL IMFs! So I hope you’ll stick around for a chat on how InterMolecular Forces determine whether molecules stick around, whether a protein flows straight through a chromatography column, or gets bound. But that’s just one example – IMFs are all around – no matter where you look, these forces can be found!
Molecules (things such as individual H2Os and protein chains) are groups of atoms that are held together really tightly – think of your fingers connected to your hands – they don’t just fall off. Similarly, the atoms in molecules don’t just “fall off” because they’re attached through a strong kind of bond called a covalent bond. But just like your hand can interact with other things, can touch and untouch, grab & let go, molecules can interact with other molecules through interactions that *can* “fall apart” – they’re directed by INTERMOLECULAR FORCES (IMFs) – different variations of these IMFs have different names, but at their heart they involve the same key concepts.
There are a few key things to keep in mind – if you can remember these you just need to “think small!” – really small – and you’ll begin to understand it all! (seriously though, I don’t know why people always laud “big thinking” when the super small is so fascinating!) So, anyways, the key points
- opposite charges attract, like charges repel
- it “doesn’t matter” where the charge comes from
- these charges don’t have to be “full charges”
- these charges don’t have to be “permanent”
- this applies both at the “large scale” (whole molecules interacting with one another through intermolecular forces (IMFs)) and the “small scale” (electrons within atoms within molecules)
- the bigger the opposite-chargeness, the stronger the attraction
- the further apart the opposite charges, the weaker the attraction
- charge-based attractions are fickle friends – they’re “weak” and “easily” reversible
- every bit counts – lots of small attractions can add up to pretty powerful stickiness (just ask the Geiko gecko’s feet!)
You’ve likely heard the “opposites attract” thing – but some of those other terms may be unfamiliar. To fully appreciate why molecules stick or don’t stick you need to understand what a molecule is and to understand that you need to know what an atom is and to kind of understand that (I’m not pretending to be an expert here!) you need to understand what subatomic particles are. But don’t worry – we’ll take it step by step and it’s not as bad as it sounds!
I often talk about amino acids as the “building blocks” of proteins and nucleotides as the “building blocks” of nucleic acids (RNA and DNA). But these building blocks are themselves “preformed pieces” made up of even smaller building blocks called atoms. When you look at a chemical structure or see a formula, each of the letters represents an atom – and the different letters stand for atoms of different elements (e.g. H for hydrogen, C for carbon, O for oxygen, N for nitrogen). Different elements are defined by having a distinct, unique, number of a subatomic particle called a proton (e.g. H has 1, C has 6, N 7 and O 8).
The proton is one of 3 key subatomic particles – protons are positively-charged and they’re held in the center of the atom (the atomic nucleus) in a dense cluster, helped kept glued together by neutral subatomic particles called neutrons. And whizzing around them are negatively-charged subatomic particles called electrons (e⁻).
So, an atom is a “unit” of an element that’s made up of subatomic particles called protons (+-charged), neutrons (neutral), and electrons (- charged), with the # of protons defining the element (e.g. hydrogen or carbon)
The electrons are constantly moving around in an “electron cloud” and you never know exactly where they’re going to be, but the places they spend the most time are called “orbitals” and the orbitals of different atoms can merge together to share electrons, and this electron-sharing leads to a strong bond we call a covalent bond. It is this sort of bond that holds together atoms WITHIN a molecule.
So, a molecule is a group of atoms held together through strong covalent bonds that involve electron-sharing.
Protons & electrons have equal but opposite charges. The # of protons is “set” for a given element – if you change the # of protons you’ve changed the element. But the # of electrons & # of neutrons is negotiable.
If you change the # of neutrons, the charge doesn’t change (since neutrons are neutral) – you just get a different “isotope” of the element which may be more or less stable (if you skew the proton/neutron balance too far you can get a really unstable nucleus that can radioactively decay)
If you change the # of electrons, the nucleus doesn’t really care – but the overall charge of the molecule changes – if you have more electrons than protons, you have a net negative charge (anion) & if you have more protons than electrons you have a net positive charge (cation).
Anions & cations can be single atoms (like the Na+ & Cl- you get when you dissolve table salt (sodium chloride)) or bigger molecules – like whole proteins.
In a neutral molecule, the # of protons = # of electrons. But since the electrons are whizzing all around – and they can spend more time hanging out in certain areas – the charge isn’t always evenly balanced everywhere in the molecule, and you often have a situation where electrons (with their negative charge) like to hang out more in one part of the molecule, making that part partly negative (δ-), leaving the other part of the molecule they fled from partly positive (δ+). We call this gradient of imbalance a dipole. δ is pronounced delta and it means partly. (Don’t confuse it for uppercase delta, Δ which means change in and we often use to avoid calculus…)
There are different ways to end up w/such partially-charged regions
Sometimes the charge difference is permanent – we call this a “permanent dipole” and it occurs because some atoms are more electron-hogging (electronegative) than others so if you put them together, the more electronegative atom wins the electron tug-of-war (a common example of this is the dipole in water molecules caused by O pulling e away from H). Since the bonds hooking the H’s up to the O in the molecule of are of that strong covalent kind, they’re not coming apart easily – and the O is “always” gonna want the electrons more – so you’re always going to have that charge imbalance -the molecule is “permanently polar”
But a warning – a molecule can have polar bonds (uneven electron sharing between atoms) but still be “nonpolar” overall – because, as the whole “pole” thing suggests (🐧 live opposite 🎅) polarity is directional. If at least some of the pulling pulls in the same direction, they add together, like in water where you have angled bonds and the O taking from both, so you get a net pole ↗︎ + ↖︎ = ↑. But the bond dipoles can cancel each other out, like in CO2, where the oxygens pull electron density equally but oppositely from the carbon between them. ← →
So, whether a BOND is polar depends on the electronegativity of attached atoms BUT whether a MOLECULE is polar *also* depends on angles at which they’re attached. To get a polar molecule you need to have NONSYMMETRIC POLAR BONDS.
Those dipoles based on the molecular makeup (which doesn’t change) are “permanent,” but dipoles can also be temporary – happening by chance since the electrons are moving around randomly and might just happen to cluster more at a certain part at a certain point in time (spontaneous dipole)- or happening due to being near another external charge (which can come from things like other molecules or electric fields) – the external charge can distort the e⁻ cloud. For example, If ➕ charge comes nearby, e⁻ will go towards it, making that side δ⁻ but leaving other side w/less e⁻ density so it becomes δ⁺. And then that can serve as an “external charge” that can distort the cloud of another molecule, setting off a “chain reaction” of cloud syncing.
As a result, even “nonpolar” molecules can be made polar (at least for a while) – and they can then make other nonpolar molecules polar
Why have I spent all this time talking about poles? Is it just that I’m excited about Santa? Nope! My enthusiasm for dipoles is year-round! Because these are key to why molecules do or don’t stick around! Because where you have charge separation you have charge differences and where you have different charges you have attraction! And when you have attraction you can have molecular interaction! So let’s take a look at INTERMOLECULAR FORCES (IMFs) in action! It’s all about syncing to the cloud – the electron cloud!
The fundamental core of IMFs is that opposite charges attract, like charges repel. Then you just have terminology to distinguish between sources of charge unevenness. The source of the external charge can be…
🔹a fully charged particle (ION) 👉 ION-INDUCED DIPOLE
🔹a partially charged part of a polar molecule 👉 DIPOLE-INDUCED DIPOLE
🔹a partially charged part of a temporarily polar molecule 👉 INDUCED DIPOLE-INDUCED DIPOLE
What if there are no “permanently-charged” things around? You can still get those INSTANTANEOUS DIPOLES I mentioned above thanks to the random movement of e⁻. If you flip a coin hundreds of times, you’ll get some “hot streaks” of heads after heads even though the coin is fair. Similarly, sometimes there will just happen to be more e⁻ on one side of a molecule & this can be enough to induce a dipole in a neighboring molecule IF the charge is strong enough, that 2nd molecule is close enough, & that 2nd molecule’s e⁻ are willing to budge.
Just as it’s easier to sync your data to the cloud when you’re closer to a data tower (where signal’s strongest), it’s easier to sync e⁻ clouds when they’re closer to charge source – IMFs decrease w/distance. So I can’t get my protein to stick to the beads unless I actually put them together – which is probably a good thing or else my columns would be full of proteins my colleagues are purifying at the next bench over!
But you can be right next to the data tower & still be unable to sync if you’ve used up all your storage. Similarly, if all the e⁻are tightly held, they’re already “being used” so they don’t budge. This is the case w/ atoms w/few e⁻ that are held close to nucleus (where p⁺ live) (you can find them in upper right of the periodic table) ↗︎
Another phrase you might come across – we call INSTANTANEOUS DIPOLE-INDUCED DIPOLE interactions DISPERSION FORCES or LONDON FORCES & they’re the weakest type of IMFs but together they add up! And, speaking of adding up, the combination of all these IMFs can be called van der Waals forces – and they literally let geckos walk up walls!
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