LanthanumK's Blog

What temperature can be reached with a three-inch magnifying glass in direct sunlight?

1. Indium melts when directly focused on it. Indium has a melting point of 156.6 degrees Celsius. At least this temperature can be reached by focusing the light onto a shiny piece of metal.

To get a higher temperature, I needed to use a dark surface. I charred the surface of a piece of wood with the magnifying glass and placed my metal sample on the dark surface. Upon focusing the light, the metal melted.

2. Tin, bismuth, and lead easily melted. Their respective melting points are 232, 271.5, and 327.5 degrees Celsius. All of these temperature are easily reached by using this method.

3. Zinc melts with difficulty, requiring an extended period of time to warm. Zinc has a melting point of 419.4 degrees Celsius.

4. Magnesium burns with difficulty. This will not work with a large piece of magnesium; to get the magnesium to burn I cut a thin strip of the foil and focused the light on the super-thin and fragile end. Magnesium autoignites at 473 degrees Celsius. Hotter temperatures may be reachable but not very easily by such a primitive method.

5. Titanium does not burn using a dark charred wood base. The wood turns white due to ash formation and disappears from under the titanium, never allowing it to reach the maximum temperature. Titanium autoignites at 1200 degrees Celsius. A different method will need to be tried to determine if this temperature can be attainable.

I have compiled a list of sources for the elements that are available to the amateur chemist. Neodymium will be discussed here.   

Neodymium is one of the more reactive rare earth metals. It forms one oxidation state, +3, and has the unique property of color-changing salts. A solution of neodymium chloride, pink in sunlight, may appear light grayish-yellow in fluorescent light. Neodymium oxide ranges from blue to purple to pink.

In element form: Neodymium magnets contain about 14% neodymium. Mischmetal has a small amount of neodymium in it.

In compound form: "Daylight" incandescent light bulbs often use neodymium oxide to make them bluish (fluorescent) or purplish (incandescent or sunlight) colored. The color change indicates the presence of neodymium.

Here is my sample of neodymium. It is a neodymium magnet.

Magnesium powder or shavings are useful in many experiments in amateur chemistry as a reducing agent. It is also interesting to burn, although repeated burning is discouraged due to the production of harmful UV rays. There are several ways to obtain magnesium in a more finely divided form.

Buy it: You can buy magnesium turnings for a relatively low price ($4.50 for 50 grams at AlphaChemicals). However, magnesium powder has some shipping restrictions on it.

Knife: When a magnesium bar is scraped, large shavings and flecks of magnesium metal come off. These flecks are useful for reductions and sparkles in pyrotechnic compositions but are less than ideal for thermites or other reactions that need a finely divided form of magnesium.

File: I have found a file to be the best way to create small amounts of finely powdered magnesium. Rubbing the magnesium against the file over a piece of paper allows the magnesium to be collected on the paper. It can later be mixed in with any oxidizer and used to create a variety of energetic mixtures.

Steel grinder: Electric grinders, though much faster than hand, tend to create air currents that blow most of the superfine magnesium powder away. Powder can be created this way but with a large amount of waste.

My first thermite reaction burnt relatively quietly as it used magnesium shavings. The last one, which used a similar oxidizer with magnesium powder, burnt much more rapidly, noisily, and quickly, showing the benefit gained by using powder (generated by a file) instead of shavings.

Trivalent iron has a significant amount of oxidative power to it. Here I will list some of the reactions of this common yet interesting substance.

Thermites: Iron(III) oxide (or red iron oxide) is one of the most common ingredients in thermites. Iron thermites tend to burn extremely hot, yet without much violence, allowing a prolonged heating of the supporting materials. See Magnesium - iron(III) oxide thermite for my experience with an iron(III) oxide thermite.

Most of the other experiments involve trivalent iron in aqueous solution as a chloro complex. To make this, iron(III) oxide is dissolved in hydrochloric acid and the resulting brown and highly acidic solution is the starting point.

Iron: Iron comproportionates with iron(III) chloride, forming a greenish solution of iron(II) chloride. This reaction goes to completion, as evidenced by the complete whiteness when the solution is neutralized with sodium carbonate. See Comproportionation Reaction for other comproportionation reactions including the iron one.

Nickel: Nickel reacts with iron(III) chloride, producing a green solution of iron(II) and nickel(II) ions. While nickel dissolves very slowly with hydrogen peroxide/hydrochloric acid and dissolution is almost nonexistent with pure HCl, nickel dissolves quite rapidly in trivalent iron solution.

Copper: Copper reacts with iron(III) chloride, forming copper(II) and iron(II) ions. However, copper(II) reacts further with copper, producing white and insoluble copper(I) chloride. This then further reacts with more copper(II) chloride, producing a brown solution. This breaks up when diluted with water to white copper(I) chloride and green copper(II) chloride again.

Bismuth: Bismuth dissolves extremely slowly and incompletely in iron(III) chloride solution. When the resulting solution is diluted, some bismuth hydrolysis is noticed.

Tin: Tin reacts with iron(III) chloride solution, forming tin(IV) ions and iron metal. Since tin and its divalent ion are both relatively strong reducing agents, they reduce iron(III) all the way to iron metal.

Lead: Lead reacts with iron(III) chloride, forming insoluble lead(II) chloride and iron(II) ions. However, the reaction does not run to completion because the lead(II) chloride forms an impermeable layer.

As you can see, iron(III) is quite a reactive substance, reacting with and reducing most metals. That is why it is used as an etchant.

I have compiled a list of sources for the elements that are available to the amateur chemist. Barium will be discussed here. 

Barium is a soft, silvery gray metal of the alkaline earth metals. It is the heaviest member of the group that is not radioactive, and also the most reactive. Barium metal reacts vigorously with water, producing soluble barium hydroxide. Barium easily absorbs oxygen and nitrogen from the air when heated. Barium is quite toxic in its ionic form, making it useful as a rat poison.

In element form: Barium metal is used in vacuum tubes and CRTs to absorb any leaking gases. It is visible as a reflective crown on the inside of a tube, but probably not visible in a CRT.

In compound form: Barium sulfate is used to map the digestive tract because it is opaque to X-rays. Barium ferrites are used in credit card magnetic strips. Electrodes in fluorescent lamps are coated with barium oxide to increase the emission of electrons. The screen of a CRT has barium oxide in it to help absorb the radiation.

Here is my sample of barium. It is a vacuum tube obtained from a friend who has a stock of them. The barium is visible on the right side.