For students taking Chemistry 101 at the University of Victoria in the Fall 2013 term.
Why is CH3CN more polar than CH3CHO?Oxygen has higher electronegativity than Nitrogen, so shouldnt the O-H or O-C bond be more polar than the NH bond or NC bond?
Is the reason why ice floats on water directly because of hydrogen bonding, or is it just because the solid form happens to form large spaces?By this logic, NH3 or other molecules capable of hydrogen bonding should form solid structures that float on there respective structures. Can solid NH3 float on liquid NH3?
Why does bmim+PF6- have a higher melting point than Bmim+BF4-?BF4- is a smaller molecule with the same charge. Therefore, the interaction with Bmim+ should better form a solid, making the melting point higher than that of bmim+PF6-
639: You can think of C and N sharing 6 electrons, so N is able to attract more charge towards itself than in a C-O or O-H bond (sharing only 2 electrons).655: Both. Large spaces wouldn't form if not for H bonding. No, solid ammonia doesn't float, because it can't form the same sort of open network as ice - there are 3 N-H bonds, but only one N lone pair.737: More efficient packing makes up for increased cation-anion distance.
Recall that like substances tend to be drawn towards eachother, like the lipid bilayer in cell membranes.If we take 2 C2H6 molecules, is there strongest force of attraction due to london dispersion forces? The molecules are not polar...Then if we take 2 O2H2 molecules, which are polar, will the forces holding together O2H2 be stronger than C2H6?The molecules are relatively the same size and length...Are london dispersion forces the only thing holding hydrocarbons together?When we say "like dissolves like," for lipids, are the forces that dissolve the lipids purely london dispersion forces, or is there some other force at work?
Pretty much "yes" to all your questions. Non-polar molecules are held together by dispersion forces.
One of the mastering chemistry questions asks you to arrange Hg, Fe and F2 in order of increasing boiling point. Why does it not go F2, Fe, Hg because of increasing mass?
So then why doesn't a hydrocarbon and a polar molecular tend to be attracted to eachother?In the polar molecules london dispersion forces are still there. And the textbook says london dispersion forces usually account for more attraction than polar groups.
912: See your notes about metallic bonding. Hg has lots of bonding electrons... but just as many antibonding.921: well, they are, but polar molecules are much more strongly attracted to each other than they are to the hydrocarbon. So they will separate into different liquid phases (e.g. oil and water).
I am in the A02 class. I wanted to do the quiz, but realised I didn't know if we had finished Chapter 5 yet. Are we done it as of today? Or is there more tomorrow?
Hello. Could you please answer the questions posted on the "Revision for labs" blog entry? Thank you.
1154: We have finished Chapter 5, yes.929: Yes, done. Basically, I'm not the person you should be contacting about lab work - your TA and the senior lab instructor are the people that can help you.
ist the following substances in order of decreasing boiling point. CH4, CCl4, CBr4.I can't do this purely using my intuition; It's something I'd have to look up. Are we expected to be able to just list these, or do we need to look it up? Is there a trend?The correct sorting is CBr4,CCl4,CH4, but i really don't get how bigger atoms form stronger covalent bonds. Aren't the electrons in a huge atom held loosely compared to a small atom?
"Hydrogen bonding occurs when a hydrogen atom is covalently bonded to an N, O, or F atom."Hydrogen has a large electronegativity difference with N,O,F. The bond would be polar covalent. Can we assume polar covalent is the same thing as covalent? In my books there is a huge difference.Also, do atoms with relatively the same electronegativities tend to form stronger bonds than those with different electronegativities?
It seems like as we move down a column, the melting point of metals gets lower. Ex. Sodium metal has a higher melting point than potassium.Doesn't more electrons/orbitals mean more molecular orbitals, which means the energies become blending more, and the overall metallic bonding becomes stronger, as they share more electrons?Why does melting point decrease down a column?
A masteringchemistry questions says to rank CCl4, Si, and Ar from lowest to highest boiling point. The correct order is Ar, CCl4, and Si.But...this only holds true if Si forms a covalent network like in diamond. How can we be sure Si will form a covalent network like this? If instead of Si, the element were something like Al, would Al form covalent networks similar to diamond/Si? Are the group 14 elements the only ones that can form these continous, covalent networks?Should the melting points of the group 14 elements be the highest out of all elements in the periodic table, then?
Dipole dipole bonding, London dispersion forces, and hydrogen forces must be overcome to boil/melt CH3CH2OH.But arent Hydrogen bonds a specific type of Dipole dipole bonding? Why do we need to overcome both h-bonding and dipole dipole bonding, when only hydrogen bonding is occuring.
True or FalseFor molecules with similar molecular weights, the dispersion forces become stronger as the molecules become more polarizable?Arent dispersion forces strictly london dispersion forces and have nothing to do with dipole moments?Why is the answer true?
Let's look at HF and HCl. HF is more polar than HCl due to electronegativity differences. Does this mean the intermolecular forces in HF are greater than in HCl?A question in MC asks why the boiling point of HF>HCl and one of the answers is that because HF is more polar. Another is that hydrogen bonds occur in HF. Its obviously hydrogen bonds, but wouldnt the HF being more polar work too? Does the quality of dipole interactions depend on polarity?E.G the greater the difference in positive & negative charge for 2 magnets means they will be attracted more strongly to eachother?
In benzene, you were able to calculate bond order for a carbon carbon bond by simply saying, "look, theres 1 and a half bonds so the bond order is 1.5." This works for the H2 molecule, which has a bond order of 1, and so there is a single sigma bond.But what about for [He2]+? The bond order is 0.5. Is there only half a sigma bond in He2+?
509: All of these molecules are non-polar by a geometric argument, so you need only consider London dispersion forces, and these get greater the more electrons you have. You're not considering covalent bonds; you're considering intermolecular forces.515: Hydrogen bonds are a special type of especially strong dipole-dipole force; so much stronger than other dipole-dipole forces that they get a special descriptor. Polar covalent means a bond between atoms with different EN; a rule of thumb is that atoms with EN differences of 0.5-2 will have polar covalent bonds. And, no, bond strength is not so easily predicted, which is why you are given tables of bond strengths.535: it doesn't always do so! Consider group 6. It's to do with the relative extent of overlap of orbitals, and the factors that determine that are beyond the scope of this course.540: Ar and CCl4 exhibit only London dispersion forces, and they can be discriminated on the basis of greater number of electrons = higher boiling point. It really doesn't matter whether Si was a metal or a covalent network, you'd expect the bonding to be stronger than this! Al is in fact a metal. The m.p.s of the group 14 elements are indeed mostly fairly high.615: you'd expect dipole-dipole forces between C-O bonds in different molecules, and London dispersion forces occur in all materials.642: "polar" and "polarizable" mean different things. Polar implies a permanent dipole; polarizable implies that an instantaneous (or induced) dipole is easy to set up, because there is a large and easily distorted cloud of electrons.646: Yes. A hydrogen bond is just an especially strong dipole-dipole interaction. The strength of dipole-dipole interactions depends on the polarity, yes.919: There is just half a sigma bond in [He2]+, yes.
Hello, this is a question from mastering chemistry which I don't understand:In which mixture do you expect to find ion-dipole forces?Ca(NO3)2 in waterCH3OH in waterCould you please explain this? Thank you.
You expect to find ion-dipole forces where you have both ions and polar molecules present. Water is polar, but ethanol is a polar molecule (so no ion-dipole forces) and calcium nitrate is made up of Ca2+ and NO3- ions (so ion-dipole forces are present).
For the first quiz question doesn't SO3 2- have resonance structures making it delocalized, meaning the answer should include all three molecules? thanks
Just piggybacking off the previous question, shouldn't there be no resonance in SO2 and SO3, since it is possibly to have a formal charge of 0 for each atom?
Also building on the 12:41 comment after looking online everything says that all three molecules have resonance structures so I'm a bit confused.
1241, 142, 242: I've passed this on to Dr Briggs. Stay tuned.
From the molecular forces, will all molecules have London-dispersion forces? And what does a molecule that it "momentarily polar" mean?What are induced and instantaneous dipoles? And what is polarizability?
Yes. "Momentarily polar" means that there is a finite probability that more electrons are on one side of an atom than another, and for an *instant*, the atom will be polar. This will *induce* an adjacent atom to also become polar. Polarizability relates to the ability of an atom to become polarized in this way; it basically scales pretty closely with number of electrons.
If we have 2 carbon chains of similar shape and length, but they vary by either a CN bond or CO bond, can we infer that one of these will have stronger attractions and thus higher melting point?Although the CO bond has a larger electronegativity difference, the CN bond is a double bond, making CN exhibit greater polar characteristics.But if we take a CO bond vs a CI bond, and assum theyre both singular sigma bonds, is it fair to assume that the CO bond will exhibit greater polar characteristics, making the overall attractional forces for CO greater than CI?
These are fairly subtle effects, but yes, overall you're basically right.
Why are ionic liquids replacements for organic solvents?Organic solvents aren't supposed to have net charges on them, and so they dissolve similar molecules.But...ionic liquids have this seperation of charge, so shouldn't it act similar to water?What exactly is an organic solvent?
I have no idea how to approach Question 25 on the 2012 midterm.How are we supposed to know Benzene has a higher melting point than Toluene? Toluene has that extra methyl group, which makes it more polarizable.I realize that it may be due to irregular packing of the molecules, but then how can we be so sure that the extra attractions of polar groups are not negated by the fact that there is inefficient packing?
545:Ionic liquids do have charges that are separated, but they tend to also have hydrophobic side chains or sections of the cation that is organic and non-polar. This presence of both charged and nonpolar sections make ionic liquids versatile solvents. Depending on the ionic liquid and its hydrophobic properties it could act similarly to an organic solvent in what it dissolves, or it could be similar to water in what it dissolves. An organic solvent is a solvent (something that does the dissolving of the solute) that is organic (it is carbon based).
Do all liquid crystals have aromatic hydrocarbons to them?
755: Yeah, this was a tricky one. This exact example is straight out of the (later) notes, though. The b.p. of toluene is higher than benzene because of greater London dispersion forces, but the m.p of benzene is higher due to very efficient packing of benzene. 931: many do, but no, not all
When asked to rank the substances Ni, S, and Pb in order of increasing boiling point, how are we supposed to best tackle this problem? ( the answer reads Ni> Pb> S )
Because of the increasing atomic radius and therefore size of a noble gas as it moves down a column, can one say that the dispersion forces and the boiling points will increase?
932: by looking up the answer. Not a question you're likely to encounter on an exam.934: Yes, you add more electrons and they are increasingly polarizable.
For long chains of hydrocarbons with 1 OH group, the dominant force is still london dispersion forces. What is the cutoff here?like...at what chain length will london forces>polar forces.So for something like methanol, polar forces are more important than london, because its only 1 carbon. But if we take like a 10 carbon chain with 1 hydroxyl, group, london dispersion forces will be the dominant force
Have a look at the boiling points of methanol vs. the different chain length hydrocarbons; methanol has a similar b.p. to hexane. So the London dispersion forces associated with -CH2CH2CH2CH3 are roughly equivalent to one H-bond (this assumes that CH3OH and CH3CH2- will have similar London dispersion forces).
sample excercise 5.3 of the text asks to predict which molecules will exhibit liquid crystalline properties.For (iii), there is a benzene ring attached to CH2 attached to carbon with a double bond oxygen attached to a single bond oxygen with negative charge. This entire thing is attached to a Na with positive charge.But..the explanation for WHY this isnt a liquid crystal is that its ionic, and that ionic substances usually have high MPs. But...it shouldnt matter what MP something has; all that is required is that the substance exhibits liquid crystal properties above its MP.So why can't ionic substances be liquid crystals?
I'm guessing you're imagining some sort of ionic liquid/liquid crystal? It's because the electrostatic interactions between cations and anions dominate over the weaker forces (London dispersion, dipole-dipole) that align the molecules along the same axis.
Any update on this comment?"For the first quiz question doesn't SO3 2- have resonance structures making it delocalized, meaning the answer should include all three molecules? thanks"
Yes. Dr Briggs has given the whole class full marks for this question.
how many pi bonds/double does benzene have?Since these bonds are delocalized do we say there are 6, or 3?In class today, Dr.Fraser said that a an no2 group can't be attached between a 2 benzene structure, because the carbon would have 5 bonds. Would it be more correct to say it would have 4.5 bonds, or 5 bonds? If 5 bonds is correct, then we assume all the bonds in benzene are double bonds
3 pi bonds, 3 double bonds - but they're delocalized over 6 positions. Regardless of where the pi bond is located or if it is delocalized, that carbon would still have 5 bonds: 4 sigma and one pi. But the only way you will be able to convince yourself of that is if you draw all the possible resonance structures.
If we look at methanol, do we assume dipole dipole, london, and hydrogen bond forces exist or just hydrogen bond and london forces?The textbook seems to say that hydrogen bonding and dipole forces play a role but hydrogen bonds are a type of dipole, so wouldnt this be incorrect?
There are still C-O dipoles in addition to the unusually strong O-H dipoles we call hydrogen bonds.About the only cases I can think of where you would have hydrogen bonding but not any other type of dipole-dipole interaction would be H2O, NH3, HF.
If asked to determine which of c6h5cl or c6h5br on a test how would you do this?A cl-c bond has a greater electronegativity difference making dipole moments stronger but br is more polarizable.How should we approach this on a test?
^ determine which has higher melting point
Yeah, that would be difficult because the two effects oppose one another. You're not likely to get a question like that one.
How can we determine the temperatures of the different liquid crystal states (nematic, sematic, & cholesteric) ?
Why does CF4 only have london dispersion forces acting upon it?
I am very confused with how to find the boiling points for a single molecule or for a whole compound. In one of the MC questions, it displayed a compound that has H bonded to N with the highest boiling point, but in the other question, the options were Kr, Cu, & HF in increasing boiling point. The correct answer was: Kr, HF, Cu. What are the guidelines for finding the bp's or mp's for compounds or single molecules?
In a question given in the SOS practice midterm, it is asked which intermolecular forces must be overcome when dissolving KBr in water. I assumed that the reaction we are discussing is the dissociation of KBr, so I thought that the dipole-dipole forces and dispersion forces of KBr itself only would have to be overcome. The answer however is ion-dipole and hydrogen bonding. I feel as if those are the forces to be overcome if BOILING a KBr solution, not simply to dissolve the KBR in the first place. Kindly lessen my confusion?
937: Because all the bond dipoles sum to 0941: because metallic bonding is stronger than dipole-dipole forces which are stronger than London dispersion forces953: what is the SOS practice midterm?
In question 23 in the practice midterm, why do the boiling points from lowest to highest follow...CF4 < CH3CH2CH2CH3 < CH3CH2OH < HOCH2CH2OH...I understand CF4 has the lowest BP because it only has london dispersion forces which are very weak.
0 hydrogen bonds < 1 hydrogen bond < 2 hydrogen bonds
The SOS practice midterm was made by students that ran the Chem 101 Midterm Help Session on Wednesday!
1) Can you please explain why the molecule CH3CH2-O-CH2CH3 has a higher melting point than CH3CH2OH ?2) Why does SeBr2 have more polar bonds than PF3?
I can't, no, because neither are true.
Post a Comment