Hello! It turns out when you take an accelerated summer course while working full time and then immediatelly roll into the fall semester, still working full time, starting biology 1 and organic chemistry 1, you are kill. I am becoming a little less kill, as evidenced by posting again for the first time in [looks at calendar] [looks at you] [hides the calendar] a while. That said, I have an organic chemistry exam torrow morning, and while I was able to coast through the first one on the back of Actually Understanding General Chemistry, this one is likely to be a little different as we are getting into new territory for me.

This article is going to be a small collection of topics I feel I need more preparation on, and will be explaining to you, dear reader, as a way to process it for me. Hopefully, everyone wins! We'll be covering:

• Acids and bases (contributers to strength of, a brief summary of categorizations)
• IUPAC naming of branched substituents and compounds other than alkanes (mid-level standard compound naming rules)
  
• Identifying and naming certain types of isomers (compounds that have the same number and types of atoms, but are shaped differently, which makes them different from each other so they get special names!)

Right, well, it doesn't seem Fourthwall has header markup, so I suppose we will go with all-caps to section things out. Allons-y!

ACIDS AND BASES AND BEARS - WAIT HOLD THAT LAST ONE

There are three typical categorization systems for acids and bases. These use different definitions to define what an acid is and what a base is, which would have potential to be very confusing... but the nice part is that each later definition includes all acids and bases from the previous definitions, so you never wind up with "this categorization calls this an acid, but this other one says it's a base!" as a problem. The categorizations, in chronological order, are as follows; keep in mind reading these that H⁺, also known as a hydrogen ion, is often referred to as a proton or bare proton. Its nucleus is usually just one proton, no neutrons, so when it loses its electron as an ion... it is just a proton floating around.

1. Arrhenius: Acids have an H that falls off the compound in solution as H⁺, while bases have an OH group that falls off the compound in solution as OH⁻.
   
2. Brønsted-Lowry: Acids produce protons, and bases accept protons. Note here that this doesn't require H in an acid or OH in a base; if a compound can take an H off of H₂O, that leaves OH⁻ behind and accepts a proton - making that compound a base by this definition! This makes the Brønsted-Lowry definition broader than the Arrhenius definition.
3. Lewis: Acids are called "electron pair acceptors" and bases are called "electron pair donors." This one gets a little funky to wrap your head around, especially if you're used to the first two definitions! This definition was created with the intent of broadening the concept of acids and bases even further so that it wasn't so water-centric (focused on H⁺ and OH⁻ falling off of things and creating H₂O). It helps to think about what is happening to the electrons in a compound when an acid dissociates. Let's look at H₂O. This is a neutral compound, with two electrons shared between each H and the central O, and four additional electrons on the O itself. The shared electrons mean each H has 2 electrons filling its valence shell. If one H falls off as H⁺, it isn't taking any electrons with it. The pair of electrons are accepted by the rest of the compound, giving us OH⁻, the conjugate base of water - the other hydrogen still shares 2 electrons with the O, the O still has its original 4 electrons to itself, but now it has the previously shared 2 electrons also to itself (giving OH⁻ that negative charge; a neutral O here would have to have 1 electron itself, so the charge is -1). Next, let's look at OH⁻ as the conjugate base to H₂O. It can pick up that H⁺ out of solution to become H₂O again, because it has that extra pair of electrons that we just said were left behind on the O. That pair of electrons are donated to the incoming H⁺ to fill its valence shell, and we get a neutral compound again. Note that H₂O is amphoteric, in that it can act as an acid or a base - the O has those two extra electron pairs, after all, so it can donate one of those to a proton to create H₃O⁺, in this case acting as a base!
   

Okay. Now we have defined what our acids and bases are. We've even alluded to the idea in the Lewis acid and base discussion that acids and bases can "turn into" each other - H₃O⁺ is the conjugate acid of H₂O, which is the conjugate acid of OH⁻, and OH⁻ is the conjugate base of H₂O, which is the conjugate base of H₃O⁺. The strength of acids and bases is related to the strength of their conjugates - if you think about it, if an acid and its conjugate base were equally strong, they'd just be playing "pass the proton" all day and none would be left to linger in solution. This is actually related to why we consider a neutral pH to be pH 7. pH looks at the concentration of H⁺ present, and there is a related concept of pOH that measures the concentration of OH⁻. Because H₂O is amphoteric, it and its conjugates tend to play "pass the proton", making its pH 7 and pOH also 7 - equal. You can imagine, looking at the first two acid/base definitions, how these two are likely related as well. As pH goes up, pOH goes down, and vice versa. The "p" part of pH and pOH indicates that we're taking the reverse logarithm of the actual concentration, so that we can look at just the exponent (and make it nice and positive most of the time, since these are usually very tiny numbers). A pH of 7 means therefore that your concentration is around 10⁻⁷. Going "down" as pH means your exponent is actually getting closer to 0, meaning your concentration is less tiny! This is why pH going up and pOH going down mean a solution is more basic (often confusing when new to acids and bases). 

So... if conjugates aren't always the same strength, then one has to be the stronger one, right? And it then follows, since we have H₂O where they do seem to be of similar strength, that different compounds must be stronger or weaker acids? How do we tell? I will be skipping over the first question for now (the short version is we measure a special solubility constant, called Kₐ, which is larger for stronger acids and lower for weaker acids). Instead, I need to worry about comparing acid strengths for this exam, so I'm sure you'll forgive me if I jump straight to that.

In our organic chemistry class, we are using the acronym ARIO (atom, resonance, induction, orbital) to evaluate acids by structure to determine relative strength. The electronic (pertaining to electrons and charge) structure of compounds has a lot to do with the reactivity and stability of compounds; you can imagine that since like charges repel and opposite charges attract, the presence of a lot of protons and electrons in a compound means a lot of repelling and attracting! These forces can either stabilize or destabilize a compound. When looking at acids, the more stable the acid itself is, the less likely it's going to be to want to lose its proton... and the more stable its conjugate base is, the less likely it's going to be to pick any more protons up. Since an unstable acid losing a proton becomes more stable, and vice versa, this means there's a relationship here between the stability of an acid and its strength.

• A (Atom): Two factors here: electronegativity (tendency to attract electrons) and size (how big is the atomic radius? How far out are its valence electrons from its nucleus?). Note that both of these only care about the atom directly connected to the H atom.
  • Higher electronegativity on an atom that is directly connected to an H atom makes for a much more stable conjugate base if the H is dropped as a proton - the atom is happy to just yoink those electrons and sit with them. Electronegativity increases up a group and to the right of a period (ignoring noble gases) - Fluorine (F) is the most electronegative atom on the periodic table, and Cesium (Cs) is the least.
  • A bigger atom connected to an H atom has more spread out electrons, with a lot more room to stay away from each other. If a very big atom drops its H as a proton , it's not really in a rush to grab it back - the electrons can avoid each other okay, making the conjugate base more stable. A very small atom, however, like F, has a very dense negative electrical charge. If it's left with an extra electron due to its H falling off as a proton, its electrons are very close to each other and not very happy about it. HF is therefore a much weaker acid than HBr, for example, despite being more electronegative. If looking at atoms in the same group, size is more important than electronegativity.
    
• R (Resonance): This refers to compounds with resonance structures. These are structures where there's more than one possible bond configuration for the same overall atomic structure of the compound - so all configurations are the same compound, and we've observed that in resonant atoms they seem to either be swapping the bonds around OR have a a structure that is partially all resonant structures at once! I am not going to go into depth on this topic here; all you need to know on this point is that resonance "delocalizes" electrons (they're not locked into one specific compound structure, they're able to "move around the compound" so to speak). The more resonance structures that a compound has, the more stable its conjugate base is going to be. The conjugate base is able to "spread out" the extra negative charge more than a non-resonant compound would be able to, so it's again not in a rush to grab extra protons to attain a more neutral charge. More resonance structures = more ability to move that charge around = less stressful on the compound to keep the negative charge = stronger conjugate base = weaker acid.
• I (Induction): Induction is about electronegative atoms in a compound that aren't directly connected to an H atom attracting electrons and pulling them away from the part of the compound where that H atom is bonded, weakening the bond between it and the rest of the compound. This electron-withdrawing group (withdraws electrons from other atoms in the compound like it's at an ATM) creates, or induces, a negative dipole moment (a small partial negative charge in its area of the compound). The closer this electron-withdrawing group is to the H atom without being directly connected to it, the stronger its effect; the more electronegative that electron-withdrawing group is, the stronger its effect. The bond-weakening effect on the H atom's bond to the compound destabilizes the acid, making it easier for the H atom to drop off as a proton and leave its electrons behind (they've been partially stolen anyways, so that makes sense, right?).
  • Some other groups can induce a dipole moment by giving up electrons, the only group we have talked about is methyl groups (CH₃); this weakens an acid by strengthening the bond between H and the compound. Withdrawing electrons = negative dipole moment = weakens bond = easier for H to fall off, positive dipole moment = harder for H to fall off. They're just opposites. Think of carbon as a bit of a wuss that gives its lunch money (electrons) up if anyone in a higher group on the periodic table so much as looks at it wrong.
• O (Orbital): Okay. I don't really have time to talk about orbital hybridization right now. This article is already getting WAY longer than it should be, and I have two more topics to discuss. It is a very cool topic I'd love to circle back to another time, but the quick and dirty version is: orbitals are how atoms hold their electrons, can hold up to 2 electrons each, and there are different kinds; "s" & "p" orbitals are primarily used for bonding to make compounds, and atoms that aren't H or He have 1 s orbital and 3 p orbitals. "Orbital hybridization" suggests that to bond with other atoms, an atom has to merge its s orbital and at least one p orbital so that it can stick together with the other atom (which, unless it's H or He, likewise hybridizes). How the atom hybridizes depends on how many atoms it wants to stick to. 1 s + 3 p = 4, so you can hybridize by combining all your s and p orbitals (sp³ hybridization) to connect to 4 other atoms. To connect to 3 other atoms, you need 2 p orbitals (1 s + 2 p = 3) (sp² hybridization). If you only need to connect to two (like the O in our old friend H₂O), you can get away with only using 1 of your p orbitals (1 s + 1 p = 2) (sp hybridization). Okay, so what does that mean for us here?  You'll notice the s orbital is ALWAYS used in these hybrid orbitals. The more hybrid orbitals you make, the more p orbitals you use, and the more the hybrid orbitals "stretch out" to avoid each other - think about it, if you stick 2 atoms on, they can be on either side of the nucleus pretty comfortably. 4 atoms, though, are getting a little too close for comfort and need to be a little further away! This means more hybridization = the H atom is also further away = its bond is longer. Longer chemical bonds are generally weaker chemical bonds and easier to break, so the more atoms that the H atom's buddy is bonded to, the easier it is for the H atom to pop off on its own as a proton... and the less likely the compound will miss it very much, now that the other bonded atoms can get a little cozier without a fourth in the room. This means that compound with an sp³ orbital bonded to the H is a stronger acid than one using an sp² orbital, and even stronger than one using an sp orbital. Remember, you can figure out which type of orbital an atom is using by counting its bonds! More bonds besides the H bond = stronger acid. 

IUPAC NAMING OF BRANCHED SUBSTITUENTS 

So. This is maybe a bit of a misnomer, as I am 90% sure the naming I'm about to review is actually common names for the branched substituents, not IUPAC standard naming. If you're unfamiliar with this, IUPAC is the International Union of Pure and Applied Chemistry, and they set up a bunch of rules to standardize how we name compounds. These rules are all based on structure, so in theory you can figure out the exact structure of a compound someone's just discovered just by picking apart the name. Handy! Also complicated. Alkanes are unbranched carbon chains - however many carbons bonded together in a single row, and all their other bonds are with hydrogens. They're the simplest hydrocarbons, and all the rules are based on them; you take the number of carbons, use that to slap a prefix on, and end it in -ane for alkane. Methane for 1, ethane for 2, propane for 3, butane for 4, and thankfully after that we get into normal Greek prefixes with pentane (5), hexane (6), heptane (7), octane (8), and so on. The first four are unfortunately grandfathered in, blame the earliest alchemists and chemists for that. 

Now, we're here to talk about branching. A "substituent" is just a fancy term for "a thingie that's attached to the main other thingie we're looking at"; a carbon attached to four hydrogens has four hydrogen substitutents. An alkane with one branch has the branch as its one substituent. If the branch is also in the form of an alkane (just appended on to a carbon on the main alkane, making for at least one carbon in the compound that has fewer than 2 hydrogen bonds) then it's named the same way as an alkane but with -yl instead of -ane. If you take butane (a string of four carbons) and slap another carbon onto it in the middle where a hydrogen would be, that's a methyl group (assuming it has the usual three hydrogen bonds). If it's a two carbon chain you're adding on, that's an ethyl group, and so on. You number the main chain's carbons, starting with 1, and label any non-hydrogen substituents in the chemical name with the number of the carbon it's on; counting is oriented on the compound so that you minimize the numbers. This means no matter which middle carbon we use in our butane examples (because there's two), it'll be named 2-methylbutane or 2-ethylbutane, since that's the lowest number we can get it to. (If we added it to the end, we'd just have a normal pentane, so it's not 1 or 4, and if it's our only substituent then there's no good reason to use 3 instead of 2.)

Okay. So... can substituents branch? Yep, starting with propyl and butyl since they have enough carbons to get funky with it! That is the problem I personally have here, because you can name them the normal way, or there's a handful of common names for the different configurations of propyl and butyl. (Please note that I'm using "alkane" here to describe how these get attached to compounds, but they're still called this regardless of if it's an alkane or another organic compound they're attached to.)

• Propyl has 3 carbons, making for a maximum of 2 configurations.
  • Normal propyl, which is a simple little chain with one end attached to an alkane, is just propyl-. Apparently its common name is n-propyl, with an n for normal. Pretty sure we all agree this is the normal configuration just because it gives everyone the least headaches.
  • The Other Propyl. 3 carbons, with the middle carbon attached to the alkane, making for a kind of V shape in the propyl group. The IUPAC name for this would be 1-methylethyl, since if you squint this can also be thought of as a 2 carbon chain (ethyl) attached to the alkane, with a 1 carbon (methyl) attached to the ethyl. It unfortunately has a common name, however, because I guess a lot of chemists don't like just using the normal rules. This is apparently isopropyl-, because it's the same atoms in a different configuration from "normal." "Iso" tends to be a very important prefix in chemistry.
• Butyl has 4 carbons, and more headaches.
  • Normal butyl is a simple 4 carbon chain with one end attached to the alkane: butyl- or n-butyl.
  • A simple butyl chain bonded by one of its two middle carbons to the alkane is either:
    • 1-methylpropyl (note the similarity to isopropyl) or sec-butyl (because chemists hate me personally), which looks like a string of 3 carbons and an extra buddy off the side right next to where it's attached
      • Okay actually after typing up tert-butyl down there I have realized this one actually makes sense - sec likely stands for secondary, and this group is considered "secondary" due to the attaching carbon having two carbon bonds in its group. So this is the "secondary" butyl form.
    • 2-methylpropyl or isobutyl (because chemists hate me personally), which looks like a string of 3 carbons and an extra buddy off the middle of that string
      • Later note in the vein of above - I guess they call isobutyl this because they don't have a better option, so I guess we can remember this one as the most unloved butyl child. Poor isobutyl.
    • Remember that there isn't a third option here despite this being a "string of 3", since if we add that fourth carbon as a buddy on carbon 3, we just get a normal string of four - n-butyl.
  • Butyl has one more configuration that's kinda like an elevated methyl group - a central carbon (this is the one attached to the alkane) connected to three carbons, which then each have three hydrogens to fill out their bonds. This is formally called 1,1-dimethylethyl, because you can look at this as 2 methyls on a chain of 2 (ethyl), or tert-butyl, which makes some sense - tert is similar to "tertiary," and this is considered a "tertiary" group since its attaching carbon is bonded to three carbons in its group. 
    

So we have n- for normal chains, we have sec- for secondary groups to distinguish from primary and tertiary groups, we have tert- for tertiary groups, and we have iso for the unloved children who are either not secondary or tertiary but also not normal (or for secondary groups that don't have a tertiary group to distiguish from... so only isopropyl). Pentyl has all of butyl's options for prefixes, but also neopentyl, which is like a butyl attached to another carbon which THEN connects to the alkane (formal name 2,2-dimethylpropyl). Clearly, the chemists were running out of options and scrambling to find something new (haha, get it? Neo? New? I'll see myself out).

ISOMERS: IDENTIFYING AND NAMING

I'll be real with you, this section might wind up wildly confusing unless I can come back later and add images. For now, you're stuck with just text, because I am in a hotel room frantically studying before a wedding I am flying home immediately afterwards because I have an exam at 9:10am tomorrow morning. This is, however, our last section... though it's possible it'll end up the longest. You'll find out when I do!

So. Isomers. A compound cannot be considered an isomer on its own; being an isomer means you have the same molecular formula as at least one other compound, but your atoms are oriented differently in space, giving you different chemical properties and therefore making you a distinct compound. This means that studying isomers is essentially studying structure and spacial orientation of atoms and bonds. Again, sorry if you have to read this without any images, I'm in a rush. 

Certain orientations in space are not considered different compounds; if two atoms are connected by only a single bond, they can spin like a top - meaning if they have additional bonds, spinning around that first bond makes their other bonds rotate around within the compound too. However, since they can just keep spinning and get back to the original shape, they aren't considered different compounds - just different confirmations of a single compound. This changes in larger molecules, though, if you get a chain of single bonds that connects to itself, called a cycloalkane for alkane chains. A cyclohexane, for example, is a ring of six carbons connected to each other in a circle (like they're about to start singing "Ring Around the Rosie" holding hands). This ring structure limits the amound of rotation that can happen, since you don't want to smack your neighbors on one side of the chain into the other side of the chain with a particularly unfortunate flip; ring structures therefore have a top and bottom side and limited conformational changes. Any atoms that have a double or triple bond are also limited in their rotation, as double and triple bonds lock their two atoms in one specific orientation towards each other. Those atoms can continue to rotate around any other single bonds they might have (other compound geometry features permitting), the double or triple bond simply will not rotate. 

• Aside: compounds will typically favor certain conformations over others. This isn't because they have different reactivity or chemical properties, but because the stress of certain conformations (if substituents get too close together, bonds get stretched out or their angles are affected, etc) makes the conformation itself have higher potential energy and want to get to a lower energy state by "righting" itself. In simple carbon chains, the lowest energy conformations are staggered (if you look at the other three bonds on one carbon, they form a Y, and the other carbon forms an upside down Y) with the angle of torsion (twisting/rotation around the bond) between its largest substituents maximized - anti refers to 180° between two groups, which is the largest possible torsion angle; gauche refers to 60°, which is the smallest possible torsion angle >0°; eclipsed, like a solar or lunar eclipse, means the two atoms are in front of each other, with a 0° torsion angle. Eclipsed is used to refer to chain conformations where the two Y shapes are both upright, and are the higher energy conformations - the highest is always an eclipsed conformation with the largest substitutents eclipsing each other. Cycloalkanes/rings are a little different since they cannot fully rotate and have to bend and twist a bit to change conformations; all side bonds can be viewed as either vertical or horizontal, with each atom on the ring having at most one vertical or axial (like the Earth's axis - if you're into geology or astronomy shut up I know, it's close enough) and one horizontal or equatorial (like the Earth's equator). Axial bonds on the same ring will be parallel to each other, alternating up and down on the ring, equatorial bonds on the same ring will be parallel to each other and facing out, away from the ring; orienting substituents to make the biggest ones axial will make them closer to knocking against other axial substituents, so conformations that make more of the compound's large substiutents equatorial are more stable. 

So. When deciding if something is an isomer or not, you look at its molecular formula - if it's not the same, it's not an isomer, it's just a different compound. Then, if it is the same, you look at its structure. Can it rotate around to become the same structure? If yes, not an isomer, because it's an identical compound. If no, you have an isomer! There are different kinds of isomers, but we are focused on constitutional and stereoisomers in class. Constitutional isomers have the same molecular formula, but the bonds are different; the atoms aren't all connected to the same things. Looking back at our branching alkane examples, tert-butyl being substituted out for an isobutyl group on an alkane would create a consitutional isomer. Same atom types and counts, but they're not bonded the same way - tert-butyl has branching that isobutyl doesn't.

Stereoisomers (stereo- used here to refer to spacial/3D qualities) have the same molecular formula, the same bonds, BUT the atoms are still oriented differently in space somehow (and the bonds can't rotate around to make the two compounds match). These are going to be compounds that wind up with the same IUPAC name, and will need some special extra prefixes to differentiate them from each other since that difference in spacial orientation gives them different chemical properties.

Some types of stereoisomers:

• cis/trans isomers have the same duo of different substituents on two carbons in the compound; usually, this is either in a ring, or two carbons connected by a double bond. A cis isomer has the substituents positioned like mirror images; if you look up cis and trans isomers, most examples use an ethene (two carbon chain doubly bonded) and the cis example is reminiscent of a butterfly's wing patterns. The top two wing sections look the same, and the bottom two wing sections look the same, but the top and bottom wing sections don't look the same. trans isomers have one side flipped - if you have substituents A and B on your two carbons and position it like the cis butterfly example,  you have A on top and B on the bottom on one side, and B on top and A on the bottom on the other side. For ring structures, this top vs bottom distinction is how you differentiate cis and trans as well (remember when I mentioned earlier that forming a ring gives you a top and bottom side?) - cis would have both A groups on top and both B groups on the bottom, and trans would have an A and a B on both sides of the ring. To name cis/trans isomers, you take the normal compound name and add either cis- or trans- to the front. Example: cis-2,3-dichloro-alkene vs trans-2,3-dichloro-alkene.
• Spinny isomers. That's clearly (not) a technical term, (don't) write that down. This type of isomer involves certain molecules having a kind of handedness, like left and right hands; hands and gloves being matched because you can't fit your left hand's thumb into a right handed glove very well is used as a common example of this type of isomer. This is chirality (cheir is Greek for hand) and is technically a simple concept, but has more vocabulary than cis/trans isomers do. These isomers can very easily be on alkane chains, and usually have at least one of what's called a chiral center, a carbon with four different substituents. You can have chiral molecules without a chiral center, but they're more complicated to identify (and usually more complicated in general); most organic molecules used by organisms are chiral. The main three terms specific to chiral isomers are enantiomers (isomers with opposite chirality on each chiral center, making them mirror-images-but-not-the-same like human hands), diastereomers (2+ chiral centers and not all of them are opposites, so they're not enantiomers but not identical), and meso compounds (2+ chiral centers, compounds which have other chiral forms but two of their chiral center conformations actually make identical compounds so they are achiral molecules - not chiral themselves). Identifying the these involves labeling chiral centers as R or S (right or sinister, left): "rank" the substituents on a chiral center (at their first point of difference, order by the highest atomic number; for any that are the same, then go to their substiutents in rank order and compare first point of difference, and so on), "rotate" the atom in your mind so that the lowest ranked substituent is behind it and out of sight, and then if you "trace" around the atom in a circle from the highest ranked substituent to the next highest, note if you are going clockwise (R) or counter-clockwise (S). These are, of course, opposites, so if there's only one chiral center, then the two isomers are definitionally enantiomers. If there's 2+, of the possible stereoisomers, the enantiomer pairs have opposite directions from each other on every chiral center - if any have the same direction, that's a diastereomer, unless they're ALL the same, in which case you've found a meso compound form and that's actually only one isomer. When you're naming these compounds, you have to append the R and S designations to the front of the name (before even cis and trans prefixes). If it's just one chiral center, (R)- or (S)- is all you need. If there is more than one, you have to add them all in a single set of parentheses, and label them with the number of carbon it's on.
  • Example: 2,3-butanediol has an -OH group on 2 and on 3; the carbons have the same set of groups, so it has a meso form, making a total of 3 isomers. These are named (2R,3R)-2,3-butanediol, (2S,3S)-2,3-butanediol, and meso-2,3-butanediol (because 2S,3R and 2R,3S happen to be identical)

This is about as far as I am going to get on this one. Hope you enjoyed and were able to learn some things, let me know if you have questions or corrections! (Unless it's about the angle of Earth's axis.)