hermanntrude

whoever taught you that didn't teach you much in the way of details. Certainly wikipedia can be unreliable. So can textbooks. So can teachers. It's important to realise wikipedia isn't perfect. However, anything you read (particularly in the science sections) in wikipedia is most likely going to be very useful in your search for the truth. For instance, if I wanted to find out what exactly the relationship was between a sample's kinetic energy and the speed at which the molecules in that sample are moving, I'd be prepared to bet that wikipedia would tell me the EXACT truth, perhaps in rather more detail than I needed. I wouldn't ask wikipedia to tell me whether OJ Simpson was really guilty or not, though. Seriously about the avocado armadillo thing... think of a number involved in chemistry. Think of the names of some scientists chemistry.

there are methods for capturing and measuring gases. I don't want to give you too many answers, since part of your project is undoubtedly supposed to be research. Try looking in your textbook in the kinetics chapter. I really don't know what the catalyst is. There is a link between the temperature of a sample and the kinetic energy of the atoms in a sample, and there is a link between the speed of the atoms and the kinetic energy they have. Find those links and the link between collision theory and temperature will become clear. Once again, your textbook would probably be quite helpful. Also, always remember that google and wikipedia are your friends.

seems like a good idea... the only trouble will be making sure that each sample of magnesium is EXACTLY the same weight. The rate of a reaction can be measured in more than one way... perhaps another way you might want to consider (and you certainly don't have to... your idea will work too) is to measure the products appearing rather than the reactants disappearing. Gases are particularly easy to measure...

Just a note before the thread is closed. The cinnamon wasn't necessary for the synthesis, BUT it is the reason it glowed. The oil of cinnamon, combined with white phosphorus is a strong phosphorescent mixture. The guy who did this first must have had trouble believing his own eyes. Holy crap, this stuff glows in the dark! I've made the sun!

don't involve tests like taste and smell and touch. Real chemistry students would be ejected from the lab if they were caught tasting or touching chemicals. Sometimes chemists will cautiously smell something if they're sure it's not toxic by inhalation. there is a technique for that. I'd guess if you want to go for a coloured powder, tell the students that it's inorganic. that way you could probably do a flame test and perhaps one or two more tests to be sure what it is. the flame test would tell you which metal was involved and then perhaps one other test would tell you the other part (sulfate, chloride, etc etc). you could make the test for the other part simpler by telling the students it is one of several choices. or by choosing something insoluble in water (there are less of those and they tend to be more distinctive. My choice would probably be copper (II) hydroxide. The tests could go as follows: flame test: gives a greenish (slightly blue) flame. This indicates the salt is copper(II) and that the other half is not a halide. solubility: forms a gelatinous mixture but doesn't dissolve. A good student might know that copper(II) hydroxide makes a weird gelatinous mixture with water. Any student ought to know that copper (II) hydroxide, and all hydroxides except those of the alkali metals and some of the alkali earths are insoluble. heating a small sample: this would result in decomposition to copper (II) oxide, which is black. This would be pretty much conclusive that the salt is copper (II) hydroxide.

bear in mind that the octet "rule" is more of a guideline. There are three classes of atoms which can disobey the octet rule: 1) the atoms Beryllium and boron (and also technically lithium helium and hydrogen, but no-one uses lithium much in Lewis diagrams for molecules and He and H are usually dealt with by mentioning that they like tohave only two within the original statement of the octet rule). These atoms like to have 4 and 6 electrons respectively. The usual stated reason for this is simply that they are quite small. I suspect there's more to it than that, but don't freak about it. There are really only six molecules to look out for in a question: [ce]BeH2[/ce], [ce]BeF2[/ce], [ce]BeCl2[/ce], [ce]BH3[/ce], [ce]BF3[/ce], [ce]BCl3[/ce]. Others do occur but don't show up very often in exams. Basically, if you see B or Be, expect them to have incomplete octets 2) radicals. these have odd numbers of electrons and so they also fit into category 1. The most common example of a stable radical is NO (nitrogen monoxide). If your electron count comes to an odd number, first check to see if you forgot to add or deduct electrons for the charge on the ion, and if you didnt, then you're probably seeing an example of category 3. The unpaired electron usually goes on the central atom. 3) atoms which can have expanded valence shells. These are found in the 3rd period and lower. NEVER draw a Lewis diagram with more than 8 electrons on an atom in the second period. The reason these atoms can do this is something you may or may not have learned yet: They have access to their empty 3d subshells, which can participate in hybridization to form either sp3d or sp3d2 hybrid orbitals. the sulfate ion is a good example of this kind of species. Sulfur is in the 3rd period so it can safely break the octet rule. You can tell from the Lewis diagram in which all the atoms have complete octets that the formal charges aren't happy. The structure with four single bonds has four negative formal charges, one on each oxygen and a 2+ formal charge on the central sulfur atom. For a good Lewis structure, we want the formal charges to be as small as possible. In an atom from the second period we would simply have to suck it up and obey the octet rule, but since sulfur is in the third period, it can disobey the octet rule and get some double bonds. Notice also that sulfur is on the list of atoms which tend to form double bonds (N, O, C, P and S... there are others but these are the most common). What you didn't notice when you wrote about the two structures you drew was that neither of them obeys the octet rule. A double bond contains 4 electrons, not two. the first structure you drew has complete octets for the oxygen atoms, but the sulfur has 12 electrons. the second structure you drew also has complete octets for the oxygen atoms but the sulfur now has 16 electrons. How to decide between the two? well let's look again at the formal charges. remember that the formal charge is calculated by taking the number of valence electrons on the atom itself, then subtracting one for each electron in a lone pair and one for each single bond. The first diagram you drew has two oxygens with zero formal charges, two with negative formal charges, and the sulfur has a zero formal charge too. The second one has all of the oxygen atoms with zero formal charges and a 2- formal charge on the sulfur. the reason this is no good is that if there has to be a negative formal charge (and there does have to be, since the total charge is -2), it prefers to be on the most electronegative atom, in this case, oxygen. It's also true that it's very rare for any atom to have more than 6 bonds on it. The only example I can think of are XeO4, which has 8 bonds (four doubles). So in this case the best compromise between formal charges and the octet rule is to have the first structure you drew, which has sulfur with 12 electrons, which fit nicely into the 6 sp3d2 hybrid orbitals. In some cases (particularly the oxyanions of the halogens... things like [ce]ClO4-[/ce], [ce]BrO3-[/ce], etc), it becomes difficult to tell exactly how much the central atom ought to break the octet rule by, and there are some nit-picky rules to follow for that decision. However, most courses just skim over that and never ask you to draw a lewis diagram for those ions. However, a useful rule I like to follow is that for oxyanions (polyatomic ions which contain a central atom and several oxygen atoms), there are as many singly-bonded oxygens as the charge on the overall ion. Here is a website with loads of examples of lewis diagrams, nicely presented to help you learn the methods. Note that it does include some of the nit-picky rules I mentioned so if you've learned another method, just use that. I hope that helps you out :0)

timing what, exactly? the safety risks should be fairly clear. The best thing to do in this type of situation is look at all the substances involved and be sure you know the hazards related to them, then look at the actual reaction and the hazards related to that (will it get hot? will it spit? might it explode? will there be any gases given off? etc)

what is your intention for this storyline? is the substance toxic? or is it just a side-line? something they have to do in class but with no consequence? what happens next? is it important to the storyline what conclusion the student comes to? At high-school level, the students would have to be given more information. Either they could be told that the sample is inorganic and (for instance) contains the sulfate ion (or something like that), or they could be told that the sample is one of a number of given substances... for instance, unknown compound A is either copper(II)sulfate, copper (I) sulfate, lithium carbonate, nickel (II) chloride or charcoal.

O...M....G! This is the most random thing i have ever seen: LINK theodore gray's article

I' not sure if there is a catalyst for this reaction, but in general it doesn't need one. My advice is to think of a standard experiment, one which uses medium concentration HCl, medium surface area of Mg, medium temperature, and then try to think of a way of measuring the rate of the reaction quantitatively (using numbers). Then once you've performed the standard experiment a few times and found out how repeatable it is (does it always give the same results?), then try altering one thing at a time. When you alter the temperature, keep the concentration and the surface area the same, and so on... the most important part is designing your experiment so you can change it easily and safely, and so that it always gives the same results when you use the same conditions.

yes, it just seems like a strange hypothetical question to ask if you have no responsibility in chemistry education. A blue powder could be a LOT of things. Sometimes context will give you a clue... for instance if you found it in a chemistry set for kids, it'd probably be a copper salt, most likely copper sulfate or copper chloride, and you could test that using a flame test. However if you have no idea where it came from you'd have to do at least a few tests. Test its properties... things like solubility, melting point, etc. The most likely thing is that it's inorganic, and if that's the case it's most likely a copper containing salt... however, you'd need to be sure of both of those assumptions before even trying to find out which copper salt. Seriously... do you have a blue powder somewhere and you want to find out what it is?

thanks for that, UC. And just for an indication of the level of risk involved... i've been considering this demonstration for a year now, and i'm only just finding the confidence to try it.

I've been toying with the idea of using thermite for a demonstration again, but I only have 8 mesh aluminum. will that be fine enough or should i go and buy something finer?

follow this link to the final page. I suspect you can use the original recipe and just add the colourants used in the recipes on the final page. I wondered whether it'd be possible to use a metal salt to make the flame one colour and a colourant to make the smoke another colour. Perhaps you could make red white and blue using red as the flame, blue as the smoke from one half and white as the smoke from the other half? (glue two recipes together with water)

if you are not a chemist, why are you being asked to develop a chemistry lab activity? A non-chemist in a chemistry lab is a dangerous thing indeed

I recently tried this. I used a 3:2 ratio of KNO3 and sugar by volume. It worked very nicely and burned quite fiercely with a fair amount of smoke (although not as much as I had hoped). I plant to repeat the experiment with some metal salts involved to change the colour of the flame. The smoke isn't very important to me anyway since i just like the fact that it burns so fast. There's always a small amount of yellow residue left over which seems to suggest the presence of KNO2... I discovered today that KNO2 is used as a food additive (preservative) so it must be fairly non-toxic, which is good to know.

I still think that if it's at all possible (and yes, it is, even if you don't have a number pad) to write µ, you should do so. Use the character map if you have to.

chemistry is physics that chemists understand. Physics is chemistry that's been scaled up a bit so physicists can understand it.

my apologies. Did you check those links in the second post? the original poster seemed satisfied with them.

Jene, this thread was started in 2005. Check the date before posting

hands up who still uses a typewriter? anyone? no? ah... ok.

esterification requires either an acid or a base catalyst, OR you can use the acid chloride instead of the acid, OR you can use a coupling agent like DCC/DCU (I may have those acronyms wrong and i don't remember what they stand for either). My guess would be that trimethyl citrate (notice it's triMETHYL, not triETHYL, as you said) would be hard to make by simple methods, since there'd be a lot of steric hindrance on the citric acid (it's crowded).

you mentioned you could smell the NO2, which suggests that either you weren't using a fume hood or your fume hood isn't good enough, either way, not smart. Don't mess with that stuff, it can easily lead to pulmonary oedema.

in general, the least electronegative element goes first. However, sometimes the order is changed in order to give some information about the structure of the compound involved. chemical nomenclature depends on the class of chemical you're naming and the context within which you are naming it.

i have no advice for you, then. it's hard to do chemistry if you don't know what your starting materials are