Water’s really wondrous – from its crazy high molarity (~55.5M!) to its ability to form hydrogen bonds, to its ability to drive together “hydrophobic” molecules. So don’t forget about it and/or blow it off as “just solvent”! It affects *everything* in biochemistry!
I don’t have time to write much new stuff, but wanted to share some thoughts about water in biochemistry – thoughts that touch on many things I talk about in other posts. So I’m going to link to those, as well as link to some other helpful resources I have found, and here’s a bit of text revamped from past posts.
Water molecules are really “sticky” towards each other because they’re highly polar – basically atoms (like the 2 hydrogens and the oxygen in H₂O) have smaller parts called subatomic particles – positive-charged protons & neutral neutrons in a dense central nucleus with a cloud of negatively-charged electrons whizzing around. Atoms can form bonds by sharing electrons, but they don’t always share fair. Oxygen is much more electronegative (electron-hogging) than hydrogen, so it pulls their shared electrons closer to itself, making the O partly negative (δ-) and the Hs partly positive (δ+). And opposite charges attract, so the H’s of one water molecule can hang out with the O of another. Each water molecule can form up to 4 “hydrogen bonds” (H-bonds) with other molecules. ⠀
Unlike the strong covalent bonds holding the Hs to the O in each individual water molecule, these inter-molecular bonds are weaker, so they can stick and unstick. As long as water molecules have sufficient energy to temporarily break free of the bonds, the water molecules can move around & explore, breaking and forming interactions with other water molecules as they travel. These water molecules can occupy many different “states” and the term we use to describe this is high entropy (aka “disorder” or “randomness”)⠀
But they can’t interact readily with hydrophobic molecules, which are characterized by being nonpolar (electrons are evenly distributed so there aren’t partly or fully charged regions) and thus don’t offer tantalizing charge opportunities. So each water molecule that has to be next to part of a hydrophobe has part of its stickiness “hidden” and is limited in its binding opportunities – it can occupy fewer “states” and thus has lower entropy (is less disordered)⠀
We use a term called free energy, G, to describe how “comfy” a molecule is – it takes into account entropy (S) (that disorder) as well as something called enthalpy (H), which has to do with bond energy. Molecules interact spontaneously in ways that make them comfier (reduce the free energy) and we describe this using the equation⠀
ΔG = ΔH – TΔS⠀
This says that the change in (abbreviated delta, Δ) free energy equals the change in enthalpy (do the new interactions have more or less energy than the old ones) minus temperature (in Kelvin) times the change in entropy (do the molecules have more freedom now?) Negative G is “good” (means a reaction is favorable) – and you can get to it if the new bonds are much less energetic (easier to hold together) and/or the molecules gain freedom of movement.⠀
Say you have a sea of water molecules and you toss in a hydrophobe. Some of the water molecules will have to hang out with it – there’s no getting around that – and because there’s only so much space around a water molecule, it’ll have to break up some of its water-water bonds to do this. And this requires putting in energy (without getting a better bond in return) so you have a + ΔH⠀
Now imagine you keep dropping in hydrophobes. Each time a water molecule swaps an interaction with water for an interaction with the hydrophobe, it has to “spend energy,” so you keep racking up “enthalpic penalties.” and it loses binding opportunities – so you have “entropic penalties” as well. But, if those hydrophobes all cluster together, fewer water molecules will have to give up the “better” opportunities offered by water. ⠀
But the hydrophobes have no intrinsic desire to clump themselves together – instead it is the water molecules around them that kind of shepherd them together – by “reaching out” to other water molecules in their network (remember each water molecule can form up to 4 H-bonds), they draw together (kind like how surface tension can lead to drops of water staying spherical – when hydrophobes combine, water molecules get released from the clathrate cage, leading to an increase in entropy (+ ΔS). And this offsets the energy you have to put in (ΔH) to break up the individual cages when you merge them. So the hydrophobic effect is ENTROPY-driven. And powerful – it’s the driving force of protein folding! more here: http://bit.ly/hydrophobiceffectPSA
And it can provide powerful “binding energy” that enzymes often take advantage of to get reactions going by allowing you to get the activation energy needed to get over an energy barrier. Water also comes into play with enzymes in terms of enzymes “desolvating” reactants – basically replacing some water-reactant bonds to enzyme-reactant ones and kinda excluding water so reactants can find one another without water hiding them. more here: http://bit.ly/enzymecatalysis
And that molarity/molality thing?
The molar mass tells us how many grams are in 1 mole (mol) of a chemical. A mole is 6×10²³ (Avogadro’s number) and it’s like the biochemist’s “dozen” – it is just a set number of things – anything – but if you ordered a mole of bagels you’d get 6.02 x 10²³ of them… (that’s 602 and and then 21 0’s…) We can use that to figure out the molaRity (mol/L). But, the volume of a liquid can change with changing temperature… at a higher temp, molecules can move around more, so they take up more space. To account for this, we can use molaLity instead of molaRity. molaLity is moles per kg solvent & we represent it with a lowercase m. Helpfully, at room temperature, the density of water is 1.00 kg/L so the molarity & molality for water-based (aqueous) solutions at room temp are basically the same. more here: http://bit.ly/solutionconcentrations
some helpful resources:
- Hydrogen Bonds, Anthony Day, 1996: http://www.cryst.bbk.ac.uk/PPS2/projects/day/TDayDiss/HBonds.html
- Introduction to Chemical Biology 128. Lecture 05. Non-Covalent Interactions, DNA; Gregory Weiss: https://youtu.be/mUc7RCkM66I?t=1304
- Chem 201. Organic Reaction Mechanisms I. Lecture 01. Arrow Pushing. Part 2; David Van Vranken: https://youtu.be/3TsRbCj2Awk?t=2249
- Chem 201. Organic Reaction Mechanisms I. Lecture 09. Addition to Sigma Star (σ*); David Van Vranken: https://youtu.be/kR1PAIRhezI?t=897