Biochem...Chapter 2...

This is mildly embarrassing because I feel like this makes me look like I just started chapter 2 yesterday...but I swear I didn't. I know chapter 2. And three! And FIVE! Four, less so...

We can start off by talking about the different types of bonds that we encounter in biochem. And chem. And everywhere. I like to think about covalent and non-covalent like...A covalent bond is like tying someone's shoe-laces together REALLY tight...you can still untie them, but it's going to require more energy to untie shoe-laces that were just interwoven without tying...which I compare to a non-covalent bond. 

In DNA and in Proteins, there's both covalent and non-covalent bonds. Nucleotides in the strands are held in covalent bonds while the 3-D STRUCTURE is stabilized by non-covalent interactions. In Proteins, amino acids are linked (we'll see this a bit later in glyglycine molecule) by covalent peptide bonds (we'll discuss that in chapter 5), while the folding of proteins is managed by non-covalent bonds.

Just some energy comparisons. Obviously non-covalent bonds are much weaker but they wouldn't be if they didn't HAVE to be. Lenny always says "nature doesn't waste energy" and in this case, I guess this would be more like... nature doesn't use less energy on purpose...or something. What I'm trying to say is that it's IMPORTANT for non-covalent bonds to be weaker because that's what lets them break and reform when they need to. This applies to shit like salt bridges and Van der Waal's forces (to be discussed later)... The textbook says that non-covalent bonds are electrostatic, meaning they rely on forces that electrical charges exert on each other

A bad summary of types of non-covalent bonds.

This is kind of all over the place, but this is a summary of non-covalent bonds and examples of those types of interactions. Charge-induced dipole and dipole-induced dipole interactions are shorter in range than permanent dipoles. VDW's are significant at short range. Hydrogen bond lengths are fixed.

We can start off by talking about charge-charge interactions. These are the "long-range" ones. They need less distance between each other to notice each other. These can be cell ions like Na+, K+, Cl-, HPO42-...etc

These interactions, when in a vacuum, depend on Coulomb's Law (above).

When the charges are like, the force between them will be +, indicating they're repelling, and when they're opposite, the force between them will be negative, indicating they're attracted to each other. 

However..in cells, there's no vacuum.

Charges are going to be separated by either water or other dielectric media...this basically means that interactions between charges are reduced, they "sense" each other less than they would if they had only air between them. 

E is the measure for dielectric constant...measures how much the electric field between charges is reduced..I guess you can compare this to walking to a cookie through water or through jello. It's easier to get there in water than in jello. I'm bad at these comparisons.

Anyway. A higher dielectric constant means a weaker interaction (less force) between charged particles. 

In water, E = 80, which is high, since the E of other organic liquids is 1-10. In water, particles interact weakly with each other unless they're really close to each other.

We can now define a newish value, the not E, but the E. Energy of interaction. That's different between E which is the dielectric constant. E is the energy that you need to separate two charged particles from a distance r to infinity. 

When the shit is attracted to each other, Q1, Q2 have opposite charges, energy is going to be negative.

As r becomes large (long, whatever), E --> 0

This basically means that in terms of charge-charge interactions, energy of interaction is inversely proportional to distance and interactions are strong over greater distances. 

This shit is important when you want to isolate proteins from a mix of other cellular crap.

We can now move onto permanent and induced dipoles. This shit depends on orientation..which end is pointed towards which end. Molecules that have no net charge (neutral crap) can become an induced dipole by getting asymmetrical internal charge distribution.

Here, μ is the magnitude of polarity...how polar the molecule is...how much separation of charge there is inside it. When you see δ it means "partial" so, not a full charge. Partial charge. Carbon monoxide has a permanent dipole of μ

Partial charges are separated by a distance x, and the dipole moment is indicated by a vectro pointing towards the partial positive end of the molecule. The magnitude of the dipole is the charge * distance

In water's case, the dipole moment is the vector sum of the two dipole moments that are along each bond between O & H...the polar bonds. Water has a dipole moment because the electrons get pulled FROM the hydrogens TOWARD the oxygen because of a difference in electronegativity. 

Here's a few dipole moments. 

Here's two things with significant dipole moments. The values of their dipole moments are large because they have real charge separation over a large distance (compared to the other molecules that don't have so much distance between separation of charges)

At neutral pH, glycine has both a positive and a negative charge on its ends. These are whole charges, separated by the length of the molecule, giving it a largeish dipole moment. 

In glycylglycine, there's a covalent link of two glycines (remember peptide bonds and shit?). The dipole moment will basically be twice as large because you doubled the charge separation (distance). 

Molecules with large dipole moments are very polar. 

This one's blurry. I'm sorry. It's late and I'm lazy.

The presence of an electric field can make molecules that can be polarized (induced to have a dipole moment) polar. You can either do it externally or it can come from a neighboring charge or dipole particle. Benzene has no net charge or dipole moment on its own but a nearby charge can force benzene's electrons to redistribute throughout the rig to produce an induced dipole moment. 

So...aromatic rings are polarizable because electrons can be displaced throughout the rings. 

In induced dipole interactions, an anion/cation can create a dipole moment in a polarizable molecule, and then become attracted to it. I KNOW this shit is sort of redundant but that's what I need. This interaction depends on how polarizable the molecule is. Permanent dipoles, not just charges, can induce dipoles.

In neutral-neutral interactions, we find the most sensitivity to changes in distance. They're strongest when the distance is small enough. Two molecules with no net charges or permanent dipoles can become attracted to each other if they're close enough, and within them, the distribution of charge will fluctuate.

VDW forces, also called dispersion forces, involve fluctuating charges synchronizing up. Two molecules can approach each other, sync up their charge fluctuations, and get a combined attractive force. You see this shit in proteins and nucleic acids. Individually, VDW forces are weak but collectively, they make really significant contributions to stability/structural integrity of shit. 

When we think about why benzenes have a tendency to stack, it's because of the fluctuation in their electron clouds, reacting with each other, producing dispersion forces. 

Now...lets consider the VDW Radius...when molecules with no covalent bonds between them approach each other, their outer electron orbitals will eventually overlap (when they come too close), causing mutual repulsion. At this rate, the repulsion will increase as the distance between the centers of the molecules decreases. That's kind of what this giant next thing is about:

This is a non-covalent energy summary of two particles that approach each other. The energy of repulsion is steep at short distances...who wants to be THAT close to someone else? It can act like a barrier until you get to rv...which is the distance of closest approach. That's the VDW distance, where you get the closest molecular packing.

ro is the distance of minimal energy. It's the most stable distance between the centers of the particles. If you let them do whatever they want, this is how close they will come to each other.

So, by this thing, the total interaction, the total ENERGY of interaction, at a distance of...whatever the distance, is going to be the sum total of Energy of attraction and energy of repulsion. 

As the distance between the particles decreases (going from the right side to the left), the attractive energy and repulsive energy increase, but at very different rates. In longer-range interactions, we have attraction controlling things, but the repulsive energy will go up so quickly that it is going to act as a prevention, which then defines the distance of closest approach and the VDW radii. 

Last thing I want to talk about is hydrogen bonds. They're technically non-covalent but they're also sort of covalent. This is confusing. They're important because they determine structure of ...pretty much fucking everything.

The hydrogen bond is defined as the interaction of a hydrogen that's covalently bonded to a "donor" atom to a pair of non-covalent electrons on an "acceptor" atom.

The ability of something to act like a donor depends on electronegativity. Something that's MORE electronegative is going to take away more negative charge from the H, making the H more and more positiveish, and attracting is strongerly (right, strongerly) to the electron pair on an acceptor atom. Only N & O are strongly electronegative enough to act as strong donors. Hydrogen bonds are going to be the strongest when there's a 180degree angle between the donor and acceptor atoms.

A hydrogen bond can be thought of/compared to a charge-charge interaction between a partially charged H (δ+) and a partially negatively charged pair of electrons (δ-). But there's electron sharing going on between the Hydrogen and the Acceptor atom...kind of covalentish. 

And the lengths of hydrogen bonds reflect this "double character". They're smaller than you would assume based on a table of VDW radii. The energy of hydrogen bonds is higher than that of other non-covalent bonds because it is slightly covalent in character. 

Tomorrow, I'll finish talking about shit in chapter three...so I'll discuss water, ionic equilibria, which I still fucking hate, then molecules that have several ionizing groups, and ionic strength. Then we'll start chapter 3, my least favorite chapter so far...deals with energy and free energy and other things that I don't believe in. Just kidding.