Sacrificial anodes are used to protect another metal from corroding. A metal with a more negative standard reduction potential is attached to the desired metal to protect it. This creates a galvanic cell were the desired metal is the cathode. The sacrificial anode will corrode leaving the cathode intact.
To show this, two cells were set up using iron nails and either zinc or nickel. The electrodes were connected by copper wire and placed in ~500mL of tap water.
Standard Reduction Potential (V)
Since Zinc has a more negative standard reduction potential than Iron, it corroded. Iron has a more negative standard reduction potential than Nickel so Iron corroded in that cell.
Figure 1 Iron and Zinc cell. Zinc corroded although it is hard to tell since zinc oxide is white
We showed this reaction using iron nails in water. Most nails come with a galvanized coating to prevent rusting. On half the nails this was removed by soaking in 12M HCl. The nails were washed and sanded lightly to remove any coating of rust. One nail with anti-rust coating and one without were placed in beakers containing ~400mL distilled water, tap water, boiled tap water, 1M NaCl, or 1M HCl.
The nails in 1M NaCl rusted the most, followed by tap water, then boiled tap water. The nails in the distilled water rusted the least. The nails in 1M HCl didn’t form Fe2O3, they caused the solution to turn green. This showed that removing impurities and oxygen from the water slowed the rate of oxidation.
The tops of liquids are at equilibrium between gas phase and liquid phase molecules. This means that in a closed container, there is a layer of vapor above the liquid. This causes vapor pressure. If possible, this gas will diffuse until it is evenly distributed throughout the space it is contained in.
To demonstrate this fact, we connected 9 sealed 500mL Erlenmeyer flasks together with glass and rubber tubing. The flask on the left side was filled with ~300mL of concentrated Hydrochloric Acid. The flask on the right was filled with ~300mL of concentrated NH3. The middle seven flasks were filled with ~300mL of water. Thymol Blue pH indicator was added to each of the flasks so we can see the pH changes in each of the flasks caused by the gas. At first the HCl solution was pink, the NH3 solution was blue, and the water was yellow. The flasks were all connected using glass and rubber tubing.
The vapor built up in the bottom flasks, it began to travel up the tubing into the second flasks. The gas dissolves in the water and slowly changes the pH. The pH of the water on the acid side goes down and the pH of the water on the basic side goes up. Once the pH on the acid side reaches 2.8-1.2, the color changes from yellow to pink. Once the pH on the basic side reaches 8.0-9.6, it changes from yellow to blue. Overtime, through the changing colors of the flasks, you can see how the gas is diffusing through the system.
The first time this experiment was set up, only the first flask on the acid side and basic side changed pH enough to change the indicator color. The experiment was set up again using ~500mL of each NH3 and HCl in 500mL side arm flasks. The seven 250mL side arm flasks in the middle were filled with ~100mL of deionized water. Thymol blue was added to each of the flasks again. The flasks were connected with glass and rubber tubing with a wider inner diameter than the first attempt.
The indicator color still only changed in the first flask on the acidic side and the first 3 flasks on the basic side. Testing with pH paper showed that the HCl and NH3 were diffusing up the flasks, but the concentration wasn’t changing enough to change the indicator color.
2.5 L of HCl can be purchased for $87.40. 500mL of HCl would cost $17.48. 2.5 L of NH3 costs $87.30, so 500mL cost $17.46. It takes 0.1g of thymol blue solid to make 250mL of thymol blue solution. We used ~1mL of this solution. Thymol blue costs $30.20 per 5 g. This means we used less than $0.01 worth of indicator.
Total cost of chemicals in this experiment was $34.94
Hmm, here’s a quick little puzzle that I’m in the middle of solving. If anyone knows the answer, feel free to let me know.
Attempting to make the acid chloride of 2-picolinic acid by SOCl2. In dichloromethane, combined ~5g of 2-picolinic acid with a slight excess of SOCl2 (~1.2:1), refluxed for ~20 minutes, removed solvent, set it aside. The initial solid was tan-brown. After a few days, it was dark. Massively dark. And green-blue, as far as the eye can tell. After a week, it’s bone dry and still massively dark. A rough yield of around 6g. So it picked up some mass…
Solubility testing… Very soluble in water, forms a deep purple solution. Soluble in methanol, solution is a bit more blue. Very slightly soluble in chloroform, purple solution. Very slightly soluble in acetonitrile, greenish leaning toward blue solution.
NMR (in D2O)… Four clean aromatic signals, consistent with the 4 protons of an ortho-substituted pyridine ring, nothing else.
UV-Vis (in water)… Peak at ~540nm, shoulder at ~650nm. Strong peak at ~260nm.
This has not been exhaustively lit searched yet, but the intensity of the color of the product is intriguing AND a ~20% mass increase makes it a fun little puzzle. We will be exploring more…
Recrystallization is a process that takes an impure substance and purifies it, removing the unwanted substances from the wanted solid. This method involves heating the solid to dissolve it in a solvent , cooling it to reform the solid, then vacuuming and drying the newly formed pure crystals. First, you have to choose a proper solvent, something that the desired solid is insoluble in at room temperature, but soluble in when heated. If the solvent is too strong, the solid won’t fully reform, resulting in a poor yield of final product. If it is too weak, the start material will not fully dissolve and the final product will still contain impurities. You then dissolve the solid in the solvent using heat, when the solid is dissolved, the impurities are also dissolved and escape into the solvent. Then the solution is removed from heat and cooled to room temperature, allowing the crystals to reform, leaving the impurities in the solvent. The recrystallization process can be helped by placing the solution in an ice bath to cool it down to 0o C. Once the crystals have formed you filter out the solvent and impurities and dry the purified solid. Once you have obtained pure, dry crystals, you can take the mass of them to discover your mass yield, and test the melting point of the solid to determine the purity of the solid. For this experiment we had crystals that contained Lead bromide (PbBr2), Lead iodide (PbI2), and Lead chloride (PbCl2). We chose water as the solvent for all of our solids. We heated all of the solutions on hot plates while they were being stirred. Once all of the solid had dissolved, we cooled the solution at room temperature, continuing to stir them. Crystals as the solutions cooled, crystals began to form, see below for pictures. Since these recrystallizations were done for the purpose of show, we did not filter and dry the solids.