Jor - 1-2-2008 at 08:55
Something very weird happened today. I wanted to dissolve some iron in hydrochloric acid, so I could test some sodium diethyldithiocarbamate on it
(with copper I got an intensely brown precitipate, about 100mg, wich made the solution very strong brown, even after diluting to approx 100mL)
I diluted 32% hydrochloric acid (Fisher certified reagent) 4 times (approx 2,5M). I then placed about 200mg iron fillings (lab-grade) in 4mL of the
diluted acid. I expected quite some bubbling , but instead, it bubbles very slowly (temperature is 10 degrees Celcius). After 20 mins, only 3/4 of the
iron has dissolved. After 30 minutes everything dissolved.
I remember we dissolved approx same amount of iron in class, in 2M HCl, and this was done in like 2 minutes. Temperature in the school is 20 degrees.
Why does my iron dissolve so slowly? Is it the temperature? Is it another factor? Its quite weird, because at school the dissolving went like 10 times
as fast.
12AX7 - 1-2-2008 at 09:27
Why not go with strong HCl and dilute it after reaction?
Things definitely go slow in a cold lab.
Tim
BromicAcid - 1-2-2008 at 15:19
Occasionally iron filings are coated with oils. Though in your case with lab grade material it is unlikely, it's still a simple matter to give some
filings a wash with some acetone, blow them dry, and give the HCl another try.
chloric1 - 2-2-2008 at 06:35
If I am not mistaken, for every increase in temperature of 10°C the reaction rate doubles? Anyway, I would do as bromic says and degrease your
filings. In fact, all the technical literature such as patents and journal articles that perform chemical proceedures on metallic surfaces etc the
first step is degrease in trichlorethane, acetone, methylene chloride etc.
For kicks , make up a table and lable 10 divisions of temperature(10°C,20°C,30°C.....100°) and the other parameters can be the minutes to
dissolve say 5 grams of metal. This would be a great excersise in reaction rate determination.
woelen - 2-2-2008 at 12:27
I recognize the experience of Jor. Many metals dissolve in acid at a slow rate, and especially if they are very pure. With this kind of experiments, I
just am patient and let the stuff sit overnight. The same slow reactions I observed for titanium (giving deep purple/blue solutions), iron, chromium
and tin. Especially tin is extremely slow, even in conc. HCl.
Impure metals tend to dissolve faster. E.g. steel-wool dissolves quite fast, while my very fine iron powder (reagent grade) is very slow. My tin (from
www.emovendo.net, very pure material) dissolves very slowly, while some technical piece of tin dissolves more rapidly
I do not have a theoretical explanation for this effect, maybe someone else could comment on that?
microcosmicus - 2-2-2008 at 12:52
Depending on the type of impurities, they might form a battery, thus dissolving
the metal faster electrolytically. For instance, suppose that the impurity consists
of some other metal which is not evenly dissolved. Then, once one puts the
metal in acid, a region with higher concentration of impurity will be at a different
potential than a region with lower concentration of impurity, so a current will
flow which should help corrode the metal faster.
The_Davster - 2-2-2008 at 16:56
Most metals have distinct crystaline forms, introduction of impurities distorts the crystal lattice, making packing within the lattice less
efficient/less dense. Whether by fitting into holes in the lattice(followed by expansion of the unit cell by repulsion) or by replacing the local
metal atoms. So on an atomic scale, more surface exists for reaction.