<p>i can help with buffers! i actually get that lol, we were doing them in class today. it’s all the same form of calculation once you get the hang of it. thelemonisplay, one formula is Ptotal = P1 + P2 + P3 etc… but there’s another which is like “a rigid vessel contains 7 atm of gases. there are 3 moles of O2, 6 moles of N2, and 1 mole of H2. what is the partial pressure of O2 gas?”. which one do you need help with?</p>
<p>A buffer is just a solution with a weak acid and its conjugate base. It just makes the pH of the solution harder to change. </p>
<p>…I wouldn’t be surprised if there was a buffer question. They’re usually thrown in Question 1.</p>
<p>Woo! 2 days and 14 hours…
Does anyone know a rough estimate to get a minimum 4?
I really need this and time is not on my side…
Thanks!</p>
<p>Rough estimate? 55%? I believe…</p>
<p>I’m not saying to go and bomb your MC, but just know the curve is great.</p>
<p>can someone clarify a few things for me:</p>
<p>Is HClO3 a strong or weak acid. Some books have different answers.</p>
<p>And can you explain formal charge to me.</p>
<p>Formal charge is the # of bonded electron PAIRS minus the number of unbonded electrons in a lewis structure. If the formal charge of each element is closer to zero, the structure is more likely to occur.</p>
<p>The curve is indeed great! When I took my 2008 practice exam MC + 3 FRQs about 4 weeks ago…with no review (we just ended our last chapter) (we simulated actual AP conditions and took it all in one freaking day on a Saturday; it was crayyyy), I only got 52/70 on the MC But I was two composite points from a 5. So I’m just reviewing my weaknesses, basically reactions! Soooo hoping a lot my score goes up! </p>
<p>Anyone super pumped to just clear out your Chem notes? It’ll be sooon overrr! Except I still have three more APs, but no biggieee.</p>
<p>@diddly The second one. I know that first formula, but it doesn’t help when they give you moles of all the gases and total pressure.
@Trixzoh HClO3 is a strong acid, I’m pretty sure that’s what it’s classified as for AP purposes.</p>
<p>How comparable is the SAT Chemistry to the AP exam? I just took a practice test in my old subject test review book and got a 750, should i get a comparable grade on the MC? Also can someone explain to me the relationship between pH and pKa? In class my teacher just said it was a number to use for the henderson hassleback equation, but in the Princeton Review they seem to be using it to directly talk about titrations and pH</p>
<p>when pH=pKa, the mols of conjug base of weak acid and mols of weak acid are equal. This is also when a buffer will be the most resistant to pH change.</p>
<p>pH=pKa+log(salt/base or acid)
When pKa=pH, it’s an effective buffer</p>
<p>Hm. On sunday, how should I prepare for the exam? Im consistently-ish hitting 50-60% MC.</p>
<p>Oh god I don’t understand IMF and bonding at all. I’ve tried reviewing different sources and stuff but I just don’t get them, however luckily they’re the only type of questions i get consistently incorrect, but there’s probably gonna be an FRQ about it… Can someone help me out? Mainly about dipoles, LDP, etc.</p>
<p>I’m going to be reviewing everything by tonight so that I can only practice tomorrow…
REALLY NERVOUS!</p>
<p>Lemon:</p>
<p>Remember PV=nRT. At constant pressure P and temperature T, you can say that partial pressure P is proportional to the number of molecules, or P alpha n. </p>
<p>Then you have Ptotal = P1 + P2 + P3 + … = (n1)RT/V + (n2)RT/V + (n3)RT/V +… </p>
<p>Under constant volume, the equation changes into Ptotal = (n1 + n2 + n3 + …) * RT/V</p>
<p>So given the total pressure and the total number of molecules in the container, you can find out what RT/V is. Then you use that to find the partial pressure of each one: P1 = (n1)*RT/V etc.</p>
<p>But to simplify that, you can just use mole fraction = partial pressure / total pressure, or n1/(total n)=P1/Ptotal and so on.</p>
<p>Buffers:</p>
<p>For Buffers, just think of this. If you have a weak acid + its conjugate base mixed in solution, they don’t react with each other because they are in equilibrium (recall HA <—> H+ + A- with Ka). Adding a salt of the conjugate base, for example NaA, which disassociates into Na+ and A-. Remember that Na+ does not react with anything in solution because it disassociates completely (it doesn’t form NaOH or anything like that). So basically you are adding more A-, which shifts the equilibrium to the left (recall Le Chatelier’s Principle). This way, you can actually make the buffer at any pH you like (and to calculate the pH, you just use the Henderson-Hasselbalch equation)</p>
<p>An effective buffer is usually such a mixture where [base]/[acid] between 0.1 and 10. Now what’s so special about it? First notice that HA is a weak acid, and A- is a weak base. For our convenience, we assume that in an reaction between a strong acid and a weak base, or strong base and a weak acid, the strong acid/base is always consumed, as long as there is still weak base/acid left in solution (other wise known as not exceeding the buffering capacity). Therefore, we have these equations (with single sided arrows):</p>
<p>A- + H+ —> HA
HA + OH- —> A- + H2O</p>
<p>To calculate the pH after adding acid or base to a buffer, just calculate how much [HA] and [A-] is left, and then use Ka and the Henderson-Hasselbalch equation.</p>
<p>The difference is that in a buffer solution of 100mL, if you add say 3 or 4 drops of HCl, the pH barely changes. But if you add 3 or 4 drops of HCl into 100mL of pure water, the pH changes drastically. </p>
<p>Acids:</p>
<p>Strong acids: HCl, HBr, HI (notice the pattern?), HNO3 (nitric acid), HClO3, HClO4, There is one diprotic acid: H2SO4. I’d say just remember it, it’s not hard. </p>
<p>Strong bases: Any hydroxides with an alkaline metal (NaOH, KOH), and any hydroxides with heavier alkaline earth metals (Sr(OH)2). Note that Mg(OH)2 actually isn’t very soluble, but still a very strong electrolyte. Therefore, it is considered a strong base as long as there isn’t too many of it (so much that it precipitates out).</p>
<p>Notice that HClO is actually a weak acid, and this can be explained using the properties of oxyacids. Increasing the number of O on the central atom (which is Cl in this case) helps stabilize the conjugate base by spreading out the negative charge, and pulls electron density away from OH bond, both of which make it more likely to disassociate.</p>
<p>-----This is only my understanding of these topics, if I’m wrong please please please correct me! </p>
<p>Source: Chemistry, the Central Science.
-s4</p>
<p>Intermolecular forces:</p>
<p>If I recall correctly, there is only 3 types: London Dispersion Forces, Hydrogen Bonds, Dipole-Dipole interaction. Other than that, I do recall that there is something about how easy molecules stack on top of each other or something… </p>
<p>Dipole-Dipole interaction:
Some molecules are polar, some are not. H2O is polar, and CH3CH2CH2CH3 (butane) is not. Polar molecules tend to dissolve in polar solvents (i.e. H2O). How do you know if a molecule is polar? You have to know the structure (Lewis structure) and molecular geometry (VSEPR), and then add all the dipole moments (which is basically the difference in electronegativity, with that little arrow (also known as vector) pointing to the more electronegative). Remember that vectors have direction and magnitude? If two vectors have same magnitude and opposite directions, they cancel. Hence, CO2 is nonpolar. CH4 is also nonpolar because it has point symmetry. Why is H2O polar? Because it is bent. Polar molecules attract other polar molecules by orienting themselves so that the negative side of their dipole moment is pointing at the positive side of another. </p>
<p>London dispersion forces: </p>
<p>Basically an addition to dipole-dipole interactions. They say that the electrons can be anywhere in their electron clouds - so it can be on one side, or the other, or spread out etc. At random times, a nonpolar molecule may gain temporary polarity because all its electrons decide to shift to one side. If there is two molecules that happen to gain temporary polarity, they attract each other. The general rule of thumb is that larger molecules have more electrons and so they have more randomness. So the larger the number of total electrons a molecule has, the more London dispersion force. </p>
<p>Hydrogen bonding:</p>
<p>Nitrogen, Oxygen, Fluorine have very high electronegativities. If bonded to Hydrogen, they basically make the hydrogen very positive. They are now known as hydrogen bond donors. Then this hydrogen is attracted to the lone pairs of N, O, F molecules (because of their high electronegativity). They are known as hydrogen bond acceptors. (To be an acceptor, it does not have to bond with hydrogen.) Just remember N, O, F, so that as soon as you see HF, or OH or NH in a molecule, you know there’s hydrogen bonding going on.</p>
<p>Let’s take H2O as an example. Why exactly is it in liquid phase at room temperature despite having such low molecular mass (18 g/mol)?
- Well there’s definitely not a lot of London Dispersion Forces, because it has like 10 electrons in total.
- We talked before that it is polar, so it attracts other H2O to itself.
- The main reason is that H2O is both a hydrogen bond acceptor and a hydrogen bond donor. In fact, because this is the case, in ice, H2O orient themselves in such a way (sort of a hexagon) that their hydrogen bonding is the most. This is why ice is less dense than water. Very interesting stuff right?</p>
<p>Source: Chemistry, the Central Science
-s4</p>
<p>I’m trying to get a 3 or 4
i have taken couple practice tests, i’ve been getting 35-40 MC and 5-7’s on my FRQ’s
What should i do to improve my scores?
My weakest link in bonding and molecular geometry</p>
<p>Well hi guys, I’m trying to get a 4 (which is going to kill me before I even get to my AP Lang exam). My question is bond order. Like which molecules would contains a bond order of 1.5?</p>
<p>^Ones with resonance structures. This is really helpful for bond order: <a href=“http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/General_Principles/Bond_Order_and_Lengths[/url]”>http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/General_Principles/Bond_Order_and_Lengths</a></p>