jdurg

Based on the fact that your initial solution had a low pH of 2.5, and the light peach color of it, I would guess that due to all the HNO3 you had added to the solution there were some nitrogen oxides in there (They are a brownish color) in a very low concentration that boiled out of solution when you put it on the hotplate. At least, that's the logical assumption one could make based on the information you've given.

My apologies for misunderstanding you. I should have a) read what you wrote more carefully, b) not jumped to a concussion.

Try not to jump while in a low-ceiling room. That should avoid the concussions.

It was perfectly clear to me what he meant.

 

Calling them "rare earths" when most of them are not rare and none of them is an earth seems unhelpful so calling them the lanthanides or actinides is more correct.

I'm more comfortable calling them Lanthanoids and Actinoids so as to avoid confusion about the negatively charged ions of Lanthanum and Actinium.

You need to certify that you are an industrial, or educational institution. Wholesale chemical supplier will not sell directly to individuals. Too much liability on their part. As a result, your only option is to buy from sellers who purchased their supply from the wholesale suppliers. That results in an increased charge to get the product as those companies need to make money to stay alive. (http://www.unitednuclear.com is an example of this).

I think one would also have to consider the formation of Lithium Nitride when doing these reactions also. Lithium metal reacts with atmospheric nitrogen, so the product you get will have some Lithium Nitride mixed in. (Can't recall off the top of my head how Lithium Nitride reacts with water, even if it does). It's just something to take into account when doing your calculations.

Welcome to a very addicting hobby! If I had some extras, I would be happy to donate to you, but the only "extras" I have are some gold and platinum, and you really can't call those elements "extra".

If you are looking for a place to purchase some samples, there are two that I used extensively in my collecting. One is unitednuclear.com, and the other is http://www.elementsales.com. The latter is run by a very good friend of mine who took our common hobby and turned it into a business. He offers VERY good quality samples and is also able to do custom samples if you need them. Some of the prices are somewhat high, but remember, he needs to make money to keep a business running, and the prices are fully worth it. Dave is a really great guy and knows his stuff quite well. His prices are much better than the prices offered by numerous other vendors out there.

If you truly want to learn more about radiation and nuclear sciences, Physics is really the best way to go. Nuclear chemistry is really quite limited, and while it's been a long, long time since I've taken a chemistry course, I believe that the basics of nuclear chemistry really just focus on the rates of decay, the types of decays, and how that decay affects the way the element/compound reacts.

Feel free to correct me if I'm wrong. I'm jet-lagged, medicated, and not in a full state of mind. (Just got back from Lisbon last night).

Well, how far along have you gotten in your own process? We are here to help people out with their questions, but we will NOT do homework for people. If you are willing to show us what you've done so far, or where you are getting hung up, we will be more than willing to help you out. It's just that the manner in which you've made this post seems more like someone looking for an answer and not looking for help.

Because many, many, many, MANY scientists have done the reactions before and the key part of scientific research is documenting everything you do. Therefore, all they have to do is look it up in a book (if it's not something that they've already memorized) and they have the proper catalyst to use?

Now how did that particular catalyst get discovered in the first place? Experimenting. For millenia scientists have tried using different chemicals to speed up the rates of reactions. Early on, it wasn't exactly safe or wise, but they would just take a random chemical and throw it into a mixture to see if it had any effect. Nowadays, the structure of the chemicals in use, the reaction pathways, and various other characteristics about the reaction itself are used to determine beforehand if a catalyst exists for that reaction and if it would work well.

But what other energy besides kinetic can light have? If light transfers any kinetic energy, it'd no longer travel at c.

 

 

Are most of those examples particle decay?

 

 

For instance, light? Or would it not be considered to be a particle and/nor energetic?

Light always travels at c. The frequency of a lightwave can be calculated using the equation c = frequency x wavelength. The wavelength of a lightwave can be calculated by the equation E (energy) = h(Planck's Constant) x v (frequency).

So because Planck's constant and the speed of light are set constants, when light loses energy (E gets lower) the frequency of the light would have to decrease. (If the total energy of the light was 20, and Planck's constant was set at 2, then the frequency would be equal to 10. If the total energy dropped down to 10, then the frequency would drop to 5 in order to keep things equal).

So if the frequency drops, then the wavelength would have to increase. As a result, the light would go from violet to red if the transition was in the visible range. When light loses energy, the wavelength of the light will increase. Theoretically, one could take UV Radiation and pass it through a medium which would absorb energy from it and make it visible.

Baking SODA needs an acid to generate co2.

 

Baking POWDER is a mixture of Baking soda, cream of tartar and only needs water to activate. It does get weaker with age , so you need fairly fresh. Humidity will degrade it faster.

 

From baking chemistry lemon juice in water is a good cheat for demonstrations, it will activate either.

Ahhh. I had forgotten the composition of baking powder.

Odd; did the baking powder have a "best before" date on it?

The whole purpose of the stuff is to fizz (slowly) when added to water-containing foods.

Incidentally, baking soda will fizz in very hot water but that might not help you much.

To see any appreciable fizz from baking powder or soda, the water must be acidic. Otherwise, what is the reaction that will lead to Na2CO3, or NaHCO3 forming CO2 gas?

You might tell your friend to search for

"ginger beer recipe"

It's likely to make a better product.

Thanks John. You just made me really thirsty for some homemade ginger ale, but it takes a good week to make it properly.

if an atom decays in a molecule, it will remain bound. the daughter product however will be an ion. this will change the propertied of the molecule and likely cause it to break appart somewhat.

I don't think that's an absolute statement that is true for every instance. For example, in the example given earlier, if T2O decays, I believe that the Helium becomes freed from the chemical bond (as the decay of T provides enough energy to break the bond) and the result is a couple free radicals which make it just as nasty.

Don't forget Bismuth. It was once thought that Bismuth was the last element with a stable isotope, but they have found that the isotope that was once thought to be stable is actually radioactive with a half-life that is in the order of BILLIONS upon BILLIONS of years.

Chlorine is an interesting element. I have, in my element collection, both gaseous and liquid chlorine. The price for puchasing a small sample of liquid chlorine is definitely worth it, as the only safe way to really obtain liquid chlorine is just too tedious and costly in it of itself to make it worth it. My liquid chlorine sample is approximately two milliliters of it sealed in a quartz ampoule further sealed in an acrylic resin casting. Liquid chlorine is truly a yellow liquid with a slight green tint to it.

For the gas, the method I used to obtain it was to take an earlenmeyer flask with a two holed stopper in it. One hole had an inlet for concentrated hydrochloric acid, and the other hole had a glass tube that led to another flask with another two-hole stopper in it. This flask was filled halfway with water. The inlet tube was under the surface of the water, and the outlet tube was above the surface. This outlet led to a test-tube with ANOTHER two-hole stopper in it. The inlet was under the surface of a bunch of sodium bicarbonate. The outlet was again above, and this was led into a large test tube with a one-hole stopper. I had the first flask filled quite a bit with a bunch of calcium hypochlorite. I poured the concentrated HCl into the first flask, then covered the top to prevent gas from escaping. Cl2 was generated and it moved through the outlet tube, into the water where some dissolved and all of the HCl dissolved, then the chlorine moved through the NaHCO3 where any water was absorbed, and the Cl2 that came out from that was pretty pure. It all collected in the large test tube where I let it collect until the tube obtained a green-yellow color. At that point, the stopper was taken off and the tube was quickly melted shut at the top. My friend then went and encased the tube in an acrylic resin and that is where it still sits, showing it's green-yellow color. Little to no gas escaped. This was all done outside and at no point did I notice any Cl2 odor.

While you do have to be safe when working with Cl2, it's not nearly as bad as working with H2S or HCN. My only bout of chemical stupidity did involve Cl2, however. I foolishly took some household bleach and put it in a sizeable bottle. At that point I poured some concentrated HCl into it and saw some vigorous fizzing and bubble formation. (Bubbles of Cl2 gas). The gas wasn't green like I thought it would be, however. So I went and put the container near my nose and inhaled. OUCH!!!!!!!!! I then learned that low concentrations of Cl2 aren't colored at all. It was painful, but no long term damage. That's when I realized why the act of smelling a gas involves wafting towards the nose and not direct inhalation.

REPLY: But still , a couple follow up questions : How can something as thin as that dead layer of skin stop what I thought was at least a fairly fast moving particle such as an alpha particle emitted during radio active decay ? Also, how can this thin layer stop such a particle if it is not hitting any thing as you said it`s electrostatic forces allowed it to do? Also,would not this alpha particle very rapidly grab from any available source [ the atoms that make up what ever tissues it passes through ] THOSE TWO ELECTRONS required to fill that first electron shell . Stop and think about it a moment. You have 2 bare naked protons fused together requiring 2 electrons to balance those 2 protons to make it an electrically balanced atom, and also the octet rule and that force also at work here, except in this unique atom it would be called the doublet rule. the force that compels atoms to fill their outer electron shells. Does an alpha particle have the strongest possible electronegativity value? It`s power of attraction for electrons. So if nothing else, it is going to be setting in motion, two sets of free radical chain reactions for every alpha particle that passes through any tissues. I do not see how these alpha particles can be as relatively harmless as you suggest they are. Your Friend, Dr.Syntax

In addition, the atoms which are left over after Pu has alpha decayed end up being two electrons rich. While the alpha particle may strip electrons off of other atoms, those electron rich atoms will donate the two electrons fairly quickly as well.

Still, I would not ingest plutonium ever. I really wouldn't be too worried about the alpha radiation (though the lining of your digestive tract would suffer some pretty nasty damage and colon cancer would pretty much become a certainty), but the gamma radiation that is also emitted would cause some nasty damage to your body.

I too have had bad experiences with chlorine. At the very least it can make you very sore for a day or two. at the worst it can kill you in a very unpleasant way.

Very true. Even as an experienced chemist, I have had a not-so-nice experience with chlorine gas. I was trying to produce chlorine gas for my element collection and performed the reaction in an Erlenmeyer flask. Not having ever seen Cl2 in person before, I had expected the gas to exhibit the characteristic green color. I knew from reading plenty of literature that it really only shows up green in a fairly strong concentration, but for some reason that escaped me. I put my nose to the top of the flask with the bubbles in it and immediately found out that yes, I did produce chlorine gas. The copious amounts of mucus in my nose, thanks to a sinus infection I was battling, probably kept me from inhaling it all into my lungs. My nose and sinuses were immediately cleared out though, and the sinus infection went away very quickly thereafter. Yes, it helped get rid of a bug, but I am never doing that again. My chlorine gas is now a sample that is quite concentrated, and even a liquified sample, securely sealed away never to be opened again.

Yep, it depends on the definition.

 

I always learned that a salt is a compound that has an ionic bond (i.e. is composed of ions).

If SnCl4 has covalent bonds, then that means it's no salt, even though it is composed of a metal and a halogen.

The interesting bit is what happens to SnCl4 when you dissolve it in water. I haven't looked it up, but my gut feeling says it will break up into ions.

 

However, I can imagine that people put this one in the category of salts, which does make the definition of "salt" a bit more vague.

When I was being taught chemistry, my chemistry teachers always emphasized that there really isn't a difference between ionic and covalent bonds. It's a continuous spectrum related to where the electrons in the bond are spending most of their time. In a salt like NaCl, the electrons are spending nearly all the time around the Cl atoms in the crystal. In a typical covalent bond H2, the electrons spend a fairly equal amount of time around each of the H atoms. Again, this is all pretty simplified but it's a good description. Try to think of bonds like temperature gradients. While we call things "hot" and "cold", it's really just a matter of which object has more heat energy in it.

If you take a look at HCl, it is by basically all defintions a covalent bond, but when you dissolve it in water it dissociates into H+ and Cl- ions.

I found this interesting little .pdf file with a google search regarding White P contamination in Red P; chemical_behaviour_of_red_phosphorus_in_water.pdf

To pretty much remove all White P in the Red P, you can go and wash your Red P with Carbon Disulfide, or some other organic solvent as White P is highly soluble in organic solvents while Red P is pretty much insoluble.

Substances which aren't phosphorus. Silicon, carbon, etc. etc. White Phosphorus is just not stable enough when exposed to the atmosphere to exist at all in the situation you describe. There may be a lot of other contaminants in the P that you have. (God that sounds horribly icky when I read it back out loud. ) I don't really have a good enough knowledge in how Red Phosphorus is produced on an industrial scale to know what else may be in there.

Unless your phosphorus was stored under water, any white phosphorus in there would have rapidly burned up in the atmospheric oxygen. The faint garlic odor is characteristic of some phosphine gas as John Cuthber pointed out, but like so many odiferous gases, the nose picks up on the odor in INCREDIBLY minimal quantities. The percentage of white phosphorus that naturally formed in your can of red phosphorus is incredibly minimal. Nowhere close to 10%.

Strong bases may also harm the glass as well, so I would avoid them.

Also, when thalidomide came out, clinical trials were a LOT different than they are now. They were conducted in the age before computers were available for the use of clinical trials. Data obtained from the clinical trials was hand collated and the tables and listings that are used for the final analysis and reporting were VERY prone to human error. Plus, since there was no computer based "checking" of the data, there was no way to validate the data to ensure it was correct.

Due to the limitations of the data collection and analysis, clinical trials at the time generally collected FAAAAAAAAAAAAAR fewer datapoints than they do now. Thanks to computers and complex databases, you can collect a large amount of data, validate it on a daily basis to ensure it is correct, and analyze and report it in a very short time.

I work in the pharmaceutical industry on the data management of clinical trials. I oversee the setup of the data collection instruments, the data validation procedures, the clarification of questionable data, the documentation of all processes used, and much, much more. The clinical trials conducted today are so much more detailed and safer than what was around when Thalidomide came to market. It's almost frightening seeing how pharmaceuticals got to market back then.

I will just re-state the obvious here, but the difference between liter of 100% ethanol and 95% ethanol is 50 mL of H2O. Ethanol is also the alcohol of intoxication, so when you purchase concentrated ethanol it typically contains a small amount of methanol and some various other compounds to make the drinking of it impossible. (Concentrated ethanol suitable for drinking is subject to MANY taxes. The type that is commonly bought for chemistry labs is denatured, yet suitable for whatever it would be used for).

If you just want something that burns and is also able to be used to make cold solutions, I agree with the acetone suggestion. Acetone is VERY cheap, VERY flammable, and when mixed with dry ice creates a solution that is incredibly cold.