The Gaming Den

Fun fact: The human attention span is about 15 minutes. 20, tops. So that's two-and-some hours of lecture that are being squandered each week, and it's really the fault of whatever dumbass instructor planned it like that.

Diets and drugs are probably going to help less and be overall riskier than simply acquiring good study habits. Don't ask me what those are, though: studying is simply not one of those things I've had to do a lot. Maybe someone else has more to say on that, though. Also, if you go to Wikihow and look up how to study, I'm sure it'll have some recommendations to start you off.

The best way to learn something is to teach it. This can take any number of forms. Study groups where you explain concepts back and forth are common and effective.

If you're studying on your own, don't just read the book! Take notes, try to get to the point where you can explain what you're reading in simple terms. Then do that. Out loud, if you have to. Your brain will try to trick you and be all like "oh I totally get this, no problem, we can move on now". When you have to explain it out loud, though, you'll see if you really know it or not.

Figure out systems for rote learning that work for you. Flashcards can be good, whatever. There's a lot of stuff you have to just *know* for chemistry, so quiz yourself.

When you do test yourself, take it seriously. Don't just flip the card over and go "oh yeah, carbon atoms have 4 valence electrons, I remember now". Make notes of the ones you didn't know, and go back and review them.

I wouldn't worry about vitamins and pills and whatnot. Get plenty of sleep, eat healthy, don't study too much at one time (and try to find productive times of day to study). Sensible living habits and good study habits are all you need.

Chemistry is really elegant, but it also makes a lot of assumptions that can be hard to follow. Remember that the discipline came into being before we ever made an electron microscope or had any idea what was actually going on at the tiny scale. So we use the approximations that work because they work, and not because of any special relationship to the small scale interactions that are actually happening.

The math really breaks down to extremely simple algebra and massive fraction reductions. Seriously, in chemistry we make water by combining 602,000,000,000,000,000,000,000 atoms of Oxygen and 1,204,000,000,000,000,000,000,000 atoms of Hydrogen. That's fucked. But the actual ratio is jut 2:1. Chemistry is entirely about dividing both sides of an equation by a very large and identical number so that it's manageable and still true.

The central concept revolves around a completely arbitrary across the board division we made called the "mole." The mole is your friend:


It's a ghastly number. It's about 602,000,000,000,000,000,000,000. And it was chosen, because it's the number of atoms of an atom or molecule that in aggregate happen to mass in at a number of grams that is equal to the relative mass of an individual atom or molecule measured in individual neutrons. Which is in turn an awful number, it's 1.67492729(28)×10−27 kg. But seriously, we just call it "one neutron." Remember we got into this situation because we were doing this for generations before we found out what the numbers actually were. If you divide both side of an equation by X, the equation is still true even if you don't know what X is.

To a first approximation, since an electron plus a proton weigh the same as a neutron, that you can just add up the protons and neutrons in an atom and that will be its molar mass. And then you can add up the molar mass of all atoms in a molecule and get the total mass of a molecule. But if you're doing really advanced chemistry, you have to take into account that the bonds themselves use energy and since energy and mass are interchangeable that the actual measured mass will be less for more stable atoms and molecules than for less stable ones. Furthermore, ions weigh slightly different amounts, because electrons have mass. But not very much mass so we usually ignore it.

If this has been at all helpful, I'll do a thing about bonding and activation energies.

-Username17

FrankTrollma wrote:The central concept revolves around a completely arbitrary across the board division we made called the "mole." The mole is your friend:

Hope the people around you haven't started calling moles 'krteks' even if it is a funny character.

As a tip of my own, I did relatively advanced chemistry in high school and my level of success was quite decent. The book we used for the first year had around a thousand pages, but I managed to find a smaller version of it, I'd call it something of a summary, that had around a hundred pages. The moral of this story is that getting a book that summarises things can help a ton.

And whoever said something about writing notes: He was right. Only the act of writing them helps a bit, but going through them a few times is even better.

All in all, I have always found it easier to read material that is compact. Standard books can be long-winded...

You don't live anywhere near me, but good doctors are always useful. And, honestly, sharing information and opinions is neat.

Let's talk about bonding, which is the process by which atoms stick together to form bigger things. There are lots of different flavors of it, which include ionic bonding, covalent bonding, hydrogen bonding and vdw forces. But basically it all comes down to two facts:

• positive charges attract negative charges and same charges repulse.
• electrons dance around nuclei in distinct shells and they are much more stable when those shells are full than when they not.

So the interaction between those two facts can get fairly complex sounding. But really it is just those two things. Basically if an atom has a number of protons that is not equal to a number of electrons that would fill a shell, then it would like to gain or lose electrons until the outer shell is full or empty (which is just like having a full outer shell, because the outer one stops existing and the next shell is the outermost and is by default full). But oh snap! Now the atom has a net positive or negative charge and is attracted to another atom (or several other atoms) that have the opposite charge. And now things stick together, and we go from This to This.

Now let's consider the different ways this can happen:

• Ionic Bonds: This is the simplest case. Sodium has just one electron in its outer shell, so if it ditches that electron its new outer shell i full. Chlorine is missing just one electron from its outer shell, so it likes to pick one up. But then Sodium is positively charged so it sticks to negatively charged things, and chloride (yes, we rename it when it becomes ionized, just go with it) is negatively charged and it sticks to positive things. So they stack together in a regular cubic pattern like magnetized oranges and the result is a cubic matrix of table salt that is of unlimited size. Seriously, as long as you keep throwing in layers of sodium and chloride you can just keep growing the crystal limitlessly.
  
  Covalent Bonds: Covalent bonds are more complicated, but they are the most important ones from the standpoint of carbon based life. The concept is that if you have two (or more) atoms that both have a relative lack of electrons that they can "share" these electrons to make hybrid orbitals that fill the shells of both parties. Now because the atoms are sharing electrons there is no net electron imbalance and no magnetic tug, but since the electrons are being shared by both atoms they still can't move away.
  
  Hydrogen Bonds: Remember how I said that two atoms sharing electrons had no net electric charge? Well that's not exactly true. Because at any given time the shared electrons are actually at a specific place. And while the shared electrons are chilling over at atom A, atom B has a relative positive charge, right? So these covalent bonds actually slosh charge back and forth along them and generate momentary static attractions and repulsions. And if the affinity for electrons is different on atom A and atom B, we can expect these relative positive and negative charges to happen a lot. We talk a lot about the Hydrogen-Oxygen bond, because we are bloated sacks of mostly water. Oxygen has a very high affinity for electrons and hydrogen would have a full outer shell if it gained or lost an electron, so in its covalent bond the electrons spend most of their time around the Oxygen. And that's why when you cool water down, it looks like This.
  
  van der Waals Attraction: Even among purely neutral covalent bonds, on a moment to moment basis electrons are sloshing around creating instaneous positive and negative charges that attract each other. so matter kind of sticks a bit at the small scale to a greater degree than gravity can account for. This was described by a Dutch dude named van der Waal and now we are stuck with the name.

The most important atoms are Hydrogen, Carbon, and Oxygen. They all have super powers that we use a lot in biochem and help us live.

• Hydrogen: Hydrogen is really just a single proton. It's so small that the only electron shell it has available is full at 2. That means it can gain or lose 1 electron to achieve maximum stability. It also means that it is incredibly small and easy to push around. Almost all matter in the universe is Hydrogen, so you can basically bank on it being an important part of pretty much anything.
  
  Carbon: Carbon is another atom that goes both ways, and is also very common. It outer shell is a hybrid SP orbit that contains 4 electrons and would be full at zero or eight. Nothing in nature is particularly happy being an undefended +4 or -4 ion, so Carbon is pretty much required by law to make covalent bonds, and it needs to make a lot of them to attain stability. And the end result is that we get something that can make these cool tetrahedrons with bonds out at every angle. We often draw it as a plus sign with the bonds out in compass directions, but they really exist in three-space and look like d4s.
  
  Oxygen: Oxygen is one of the most common elements in the galaxy and it is incredibly electron grabby. Only Fluorine is more adamant about taking electron from things, but while Fluorine is happy enough taking just one electron (and thus serves no meaningful structural purpose), Oxygen needs two. Which is enough to attach one thing to another thing with some of the most powerful bonds in the universe.

-Username17

I know a fair bit about chemistry. A few expansions on what Frank said:

Charge density is an important concept; this is just the charge something has (protons - electrons, in chemistry), divided by its volume. I'll come back to it repeatedly, but, in general, higher charge densities are more reactive if free, and form stronger bonds when not.

Hydrogen is indecisive about losing its electron. If it gains an electron, it's perfectly happy with its two electrons, but if it loses its electron it becomes just a bare proton, which is actually extremely reactive because of its charge density (most atoms are on the order of 10^-10 meters; protons are on the order of 10^-15 meters). This is called an acid. Usually, stronger acids are made with things that have a stronger pull on the hydrogen's electron (Chloride, Bromide, and Iodide make strong acids), but Fluoride pulls so hard that the hydrogen sticks with it. Hydrofluoric acid is extremely reactive in other reactions, too, making it more corrosive than its acidity would indicate.

You can also make an ion out of multiple covalently bonded atoms. Typically this will be a negative ion made out of one X atom and some oxygen atoms, with enough electrons to fill up all of their shells. X can be anything, Sulfur (SO3--, SO4--), Chlorine (ClO4-), and so on. An exception to both of those rules is Ammonium (NH4+).

There's also Metallic bonding. Here, basically, you have a whole mess of metal atoms that all want to lose electrons, so they do, creating a bunch of free electrons. The ions arrange themselves into a crystal structure (similar to ionic bonding, but without - ions), held together by all of the free electrons between them. Because the electrons are free to move, heat and electricity are conducted well by metals. Biochemistry doesn't care much about this kind.

Ionic, Covalent, and Metallic bonding are much stronger than Hydrogen and Van der Waals bonding. The former three are lumped together as Primary Bonding, and the latter two as Secondary bonding.

Hydrogen bonding is substantially stronger than Van der Waals bonding.

An important concept Frank touched on is Polarity. Any bond between different elements is probably polar; exactly how polar depends on the electronegativity difference. Fluorine has the highest electronegativity, followed by Chlorine, then Oxygen and Bromine; high-period alkali metals like Francium and Cesium have low electronegativities.

When you have a polar bond, the electrons tend to clump on one side of the bond more than the other, creating what's called a Dipole Moment, which you can symbolize with an arrow (vector) pointing from the lower electronegativity to the higher, with a length proportional on the difference. Adding up all of the dipole moments in a molecule (remember that direction is important. Add arrows by moving them to put the tail of the second on the head of the first, then drawing a new one from the tail of the first to the head of the second) gives the molecular dipole moment. Highly symmetric molecules (like linear CO2 and tetrahedral methane) tend to be non-polar overall, since symmetric dipole moments cancel out.

I imagine that in biochemistry, when you have really big molecules, the dipole moment for small parts of the molecule (single amino acids in a protein, for instance) is more important than the moment for the whole molecule, but I'm not a biochemist.

Polar molecules tend to stick to other polar molecules well, and to ions, by matching unlike poles; non-polar molecules don't stick together too well. This is why Methanol (CH3OH) is liquid at room temperature, while Methane (CH4) is not; the oxygen-hydrogen bond is very polar, while the dipole moments in methane are small and all cancel eachother out. In general, displacing hydrogen to add a hydroxyl group (OH) to something makes it more polar.

This is what Frank was talking about with water. It's also why so many ionic compounds dissolve in water; the water molecules can match their poles to opposite ions, pull them out of their lattice, and then surround the ion so that it becomes a lump of water molecules with more diffuse charge (lower charge density), which keeps them loose in water.

Larger molecules have stronger Van der Waals forces, which is why the first few hydrocarbons (Methane, Ethane, Propane, Butane) and so on are gasses at STP (0 C, 1 atmosphere pressure), but Pentane and up are liquids (Gasoline is a mixture of heptane, octane, and a bunch of other things)

Are there any specific concepts you need explained?

Koumei, could you maybe mention some specific topics you'd want some clarification on? I assumed we were talking about organic or molecular chemistry, but most of the stuff Frank talked about I remember from the intro classes at college, so if we knew what things specifically you'd like help with it'd better focus the attention.

I did great in High school Chem, and learned everything I needed to beat college chem. Which meant the class was so boring, I nearly flunked the tests.

Follow frank's advice and remember it's all just blind puzzles, and you'll be fine.

One tip: Try to get sample or last-year's tests from the instructor, on the assumption you want to know what he's asking so you're studying he right stuff. That's a good way to make sure you're not prepping wrong.

-Crissa

Chemical Engineer Major here. Some more general "study tips" that I personally found helpful for chemistry:


-Become practiced at unit analysis. If you were having problems with the math of titration, this may be one of the problems - if you know how to do unit analysis everything becomes so much easier, and you can even easily write formulas you may have totally forgotten based on it alone.

-Don't just stare at the books. It's better to do example problems for 1 hour than try and memorize text for 3 hours. This goes for pretty much any science/math course. You really do learn stuff better by continually writing it down and using it.

-Keep a sheet of any formulas you have to memorize. Write them out both in symbols, and in words. I've found associating the meaning of the words to them helps, a lot.

-Learn to think of things in ranked order, remembering the why for it. This should not be difficult for a denner. Chemistry has a lot of stuff where there's a whole list of things that may or may not apply, and generally you only care about whatever's there and on top of the list. Bond strength, melting temperatures, electronegativities, boiling points, acidity - all things you'll want to remember in ranked order, and with a general idea of why it happens.

-Are you dealing with geometry of molecules, angles, and shapes? If so, try actually constructing some out of paper - it's not hard, and you maybe able to find some online. I remember the methane "molecules" we made out of paper in GenChem 1 and thought it kinda silly, but I can remember the exact shape and bond angle (109.5) just fine. (Edit: As Frank pointed out, some of these are actually the shape of dice that a normal person wouldn't be familiar with, but you are.)

-You might not be taking Organic Chemistry now. You might not ever take it (if you want to be a pharmacist, you probably will). If you do take it, your professor may suggest making reaction flash cards. You'll probably find this a bit inane and childish. Listen to him - make those goddamn cards. You can tell the OChem students in the library as they're the ones carrying around a box of index cards, and for good reason.

I'll second the advice for specific topics. Even just a general stuff of what you feel like you're missing from year 11, or doing now.

Larger molecules have stronger Van der Waals forces, which is why the first few hydrocarbons (Methane, Ethane, Propane, Butane) and so on are gasses at STP (0 C, 1 atmosphere pressure), but Pentane and up are liquids (Gasoline is a mixture of heptane, octane, and a bunch of other things)

It's worth noting the why here - Van der Waals forces involve two molecules kinda "touching" each other. This is based off of surface area - molecules with larger surface area have stronger Van der Waals forces. For molecules that are the same "size" but with different surface area, this is important. Although it may be going into too much detail then what you're interested in now.[/i][/b]

Koumei wrote:H2O + SO2 -> ???

H2SO3 has it happens. Which is itself an acid. so expect it to subsequently go to SO3H- and H+. And then to SO3-2 and 2H+.

1. Balancing an equation. I get that if you have 3H and 7O and 2N on the left side, then you need that on the right side. And that in some cases you might have to multiply part of the things on one side (and thus, also on the other side) to get a nice balanced amount that actually forms something.

Right, the atoms on the left side must equal the atoms on the right side. In general, the simplest whole number ratios are the correct ones. If you get a ratio of 2:1 you're probably looking at 2 and 1. If you get 1.5:1 you're probably looking at 3:2. And so on.

There are exceptions, especially in organic and biochem when you deal with polymers. For example, a Carbohydrate has the ratio of 1C:1O:2H (it's name is literally "Carbon and Water"), but they can be of literally any length. Glucose is C6O6H12, and Glycogen may well be C504000O504000H1008000. You know, whatever. But as long as you have reduced whole number ratios your stuff will work out from a lot of perspectives. For example, it doesn't really matter how long any particular molecule is for purposes of energetics, because you get a roughly constant amount of energy out of the Carbohydrate + Oxygen -> CO2 + Water reaction. So what matters is how many moles of Carbon you are throwing around, not specifically how many moles of actual sugar molecules you have.

2. The whole volumetric analysis thing. Okay, so I can find the titration point of H2SO4 + H2O + H2O2 using KMnO. Great. And I know the concentration, volume and molarity of the Sulfuric Acid there, and the volume of the H2O2. Apparently this can tell me the concentration and molarity of it. How? I just don't understand it.

OK, Sulfuric Acid is a "strong acid" which means that it deprotonates essentially completely. That means every mole of Sulfuric Acid produces 2 protons. These protons are individual acid sources. Each Acid will detonate with a Hydroxide to form water and cease to exist when that is an available option. So when you know the moles per liter of starting protons and the number of liters you know an actual number of protons you have in a pile. When you cancel them back to neutrality (your titration point), that means that the number of protons and hydroxides were equal.

So you at that point know an actual number of hydroxides you put in. So you also know that the moles per liter times the number of liters equals that number. So you can pretty easily calculate one from the other.

-Username17

But how am I supposed to tell what will form what? For instance, H2O and SO2 are mixed together.

H2O + SO2 -> ???

How much of this is required for you to learn at this point, and generally, what kind? The specifics of how you're supposed to tell depend on those, and can range from very easy to mind-numbingly complex, depending. The only reactions I ever remembered needing to learn how to write the products for myself at that point were Acid-Base reactions (which are generally very easy), metals + acid on the activity table (which is a bitch to memorize), and combustion reactions (which, again, are very easy).

Most of the times I've seen, the full reactions are usually given except for the above - but you may be doing something differently than I did.

And I know the concentration, volume and molarity of the Sulfuric Acid there, and the volume of the H2O2. Apparently this can tell me the concentration and molarity of it. How? I just don't understand it.

It's worth noting, "Molarity" IS a concentration value. There are others, too, such as Normality, Molality, %mass, %volume, and so on, and are all *generally* easy to calculate one from another. (There are a few exceptions when you get into non-ideal solutions, but that's generally covered in a thermodynamics course).

Frank's explanation is really good, but I think it would be helpful to list the steps on their own, and give examples. I'm not familiar with the exact reaction you used, so apologies if I get something off with respect to that.

Generally, you titrate to a point where everything has reacted. First, you need to take the concentration and volume you used and know to get the number of moles. (concentration x volume) gives you an amount. (mole/L * L = mole) is the way to think of it, unit analysis wise. If you have an acid with a concentration of 6M (mole/L) and use 0.5L, then you used (6 mole/L * 0.5L) = 3 moles of acid.

You then need to figure out how many of those moles are needed to react. This is found in the coefficients balanced equation. To give an example, where "A" and "B" are not defined...

H2A + 2 BOH -> 2 H2O + AB2

If you're titrating with the acid (species H2A) you know you need 1 mole for every 2 moles of base that will react. So if you know from the previous example that you put in 3 moles of acid, then you know that you reacted with 6 moles of base. (3 moles acid * 2 moles base/1 mole acid) = 6 moles base.

Now, you should hopefully know the volume of the base. You find the concentration in exactly the reverse of the first step - (6 moles base / 0.100 L base) = 0.6 mole/L, or 6M.

Koumei wrote:But how am I supposed to tell what will form what? For instance, H2O and SO2 are mixed together.

H2O + SO2 -> ???

Like DragonChild mentioned, this can be somewhat complicated depending on what you're dealing with. Strictly speaking, it isn't really part of balancing equation. Balancing the equation means stuff in = stuff out. If they tell you in the problem what you're getting out and ask you how much, you balance the equation and tell them. But if they don't tell you what you're getting out, it's a whole different story. There's helpful rules and tables, but there's kind of a lot of them, which I'm hoping someone who's taken chem more recently than me will be more able to recall.

You're supposed to know that H2 will form more readily than HO, and that SO2 is unstable like HO is; so you end up with 2 H2 and SO4 based upon the shapes of the little diagrams you're supposed to memorize. But even that varies by temperature, and they won't ask you those edges in the level you're in.

-Crissa