Light Bulb in Hydrofluoric Acid (HF)

Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. It is a precursor to almost all fluorine compounds, including pharmaceuticals such as fluoxetine(Prozac), diverse materials such as PTFE (Teflon), and elemental fluorine itself. It is a colourless solution that is highly corrosive, capable of dissolving many materials, especially oxides. Its ability to dissolve glass has been known since the 17th century, even before Carl Wilhelm Scheele prepared it in large quantities in 1771.^[2] Because of its high reactivity toward glass and moderate reactivity toward many metals, hydrofluoric acid is usually stored in plastic containers (although PTFE is slightly permeable to it).^[3]

Hydrogen fluoride gas is an acute poison that may immediately and permanently damage lungs and the corneas of the eyes. Aqueous hydrofluoric acid is a contact-poison with the potential for deep, initially painless burns and ensuing tissue death. By interfering with body calcium metabolism, the concentrated acid may also cause systemic toxicity and eventual cardiac arrest and fatality, after contact with as little as 160 cm² (25 square inches) of skin.

Pouring Mercury into Liquid Nitrogen (slow motion)

Mercury is a chemical element with symbol Hg and atomic number 80. It is commonly known as quicksilver and was formerly named hydrargyrum(/haɪˈdrɑːrdʒərəm/).^[3] A heavy, silvery d-block element, mercury is the only metallic element that is liquid at standard conditions for temperature and pressure; the only other element that is liquid under these conditions is bromine, though metals such as caesium, gallium, and rubidium melt just above room temperature.

Mercury occurs in deposits throughout the world mostly as cinnabar (mercuric sulfide). The red pigment vermilion is obtained by grinding natural cinnabar or synthetic mercuric sulfide.

Liquid nitrogen is nitrogen in a liquid state at an extremely low temperature. It is a colorless clear liquid with a density of 0.807 g/ml at its boiling point (−195.79 °C (77 K; −320 °F)) and a dielectric constant of 1.43.^[1] Nitrogen was first liquefied at the Jagiellonian University on 15 April 1883 by Polish physicists, Zygmunt Wróblewski and Karol Olszewski.^[2] It is produced industrially by fractional distillation of liquid air. Liquid nitrogen is often referred to by the abbreviation, LN₂ or "LIN" or "LN" and has the UN number1977. Liquid nitrogen is a diatomic liquid, which means that the diatomic character of the covalent N bonding in N₂ gas is retained after liquefaction.^[3]

Liquid nitrogen is a cryogenic fluid that can cause rapid freezing on contact with living tissue. When appropriately insulated from ambient heat, liquid nitrogen can be stored and transported, for example in vacuum flasks The temperature is held constant at 77 K by slow boiling of the liquid, resulting in the evolution of nitrogen gas. Depending on the size and design, the holding time of vacuum flasks (Dewars) ranges from a few hours to a few weeks. The development of pressurised super-insulated vacuum vessels has enabled liquefied nitrogen to be stored and transported over longer time periods with losses reduced to 2% per day or less.^[4]

The temperature of liquid nitrogen can readily be reduced to its freezing point 63 K (−210 °C; −346 °F) by placing it in a vacuum chamber pumped by a vacuum pump.^[5] Liquid nitrogen's efficiency as a coolant is limited by the fact that it boils immediately on contact with a warmer object, enveloping the object in insulating nitrogen gas. This effect, known as the Leidenfrost effect, applies to any liquid in contact with an object significantly hotter than its boiling point. Faster cooling may be obtained by plunging an object into a slush of liquid and solid nitrogen rather than liquid nitrogen alone.

Pythagoras Cup (Greedy Cup) filled with Mercury

A Pythagorean cup looks like a normal drinking cup, except that the bowl has a central column in it – giving it a shape like a Bundt pan. The central column of the bowl is positioned directly over the stem of the cup and over the hole at the bottom of the stem. A small open pipe runs from this hole almost to the top of the central column, where there is an open chamber. The chamber is connected by a second pipe to the bottom of the central column, where a hole in the column exposes the pipe to (the contents of) the bowl of the cup.^[1] [2]

When the cup is filled, liquid rises through the second pipe up to the chamber at the top of the central column, following Pascal's principle of communicating vessels. As long as the level of the liquid does not rise beyond the level of the chamber, the cup functions as normal. If the level rises further, however, the liquid spills through the chamber into the first pipe and out the bottom. Gravity then creates a siphon through the central column, causing the entire contents of the cup to be emptied through the hole at the bottom of the stem. Some moderntoilets operate on the same principle: when the water level in the bowl rises high enough, a siphon is created, flushing the toilet.^[3]

Underwater Caesium - Periodic Table of Videos

Caesium metal is one of the most reactive elements and is highly explosive when it comes in contact with water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion. This can occur with other alkali metals, but caesium is so potent that this explosive reaction can even be triggered by cold water.^[10] The autoignition temperature of caesium is also −116 °C, so it is highly pyrophoric, and ignites explosively in air to form caesium hydroxide and various oxides. Caesium hydroxide is a very strong base, and will rapidly corrode glass.^[15]

The isotopes 134 and 137 are present in the biosphere in small amounts from human activities and which differs between locations. Radiocaesium does not accumulate in the body as effectively as many other fission products (such as radioiodine and radiostrontium). About 10% of absorbed radiocaesium washes out of the body relatively quickly in sweat and urine. The remaining 90% has a biological half-life between 50 and 150 days.^[107] Radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.^{[108][109][110]} Plants absorb caesium differently, some do not absorb it much, and some take it large amounts, sometimes displaying great resistance to it. It is also well-documented that mushrooms from contaminated forests accumulate radiocaesium (caesium-137) in their fungal sporocarps.^[111] Accumulation of caesium-137 in lakes has been a high concern after the Chernobyl disaster.^[112][113] Experiments with dogs showed that a single dose of 3.8 millicuries (140 MBq, 4.1 μg of caesium-137) per kilogram is lethal within three weeks;^[114] smaller amounts may cause infertility and cancer.^[115] TheInternational Atomic Energy Agency and other sources have warned that radioactive materials, such as caesium-137, could be used in radiological dispersion devices, or "dirty bombs".^[116]

Underwater Potassium - Periodic Table of Videos

Potassium metal reacts very violently with water producing potassium hydroxide (KOH) and hydrogen gas.

2 K (s) + 2  H₂O (l) → 2 KOH (aq) + H
2↑ (g)

This reaction is exothermic and releases enough heat to ignite the resulting hydrogen. It in turn may explode in the presence of oxygen. Potassium hydroxide is a strong alkalithat causes skin burns. Finely divided potassium will ignite in air at room temperature. The bulk metal will ignite in air if heated. Because its density is 0.89 g/cm³, burning potassium floats in water that exposes it to atmospheric oxygen. Many common fire extinguishing agents, including water, either are ineffective or make a potassium fire worse.Nitrogen, argon, sodium chloride (table salt), sodium carbonate (soda ash), and silicon dioxide (sand) are effective if they are dry. Some Class D dry powder extinguishers designed for metal fires are also effective. These agents deprive the fire of oxygen and cool the potassium metal.^[100]

Potassium reacts violently with halogens and will detonate in the presence of bromine. It also reacts explosively with sulfuric acid. During combustion potassium forms peroxides and superoxides. These peroxides may react violently with organic compounds such as oils. Both peroxides and superoxides may react explosively with metallic potassium.^[101]

Because potassium reacts with water vapor present in the air, it is usually stored under anhydrous mineral oil or kerosene. Unlike lithium and sodium, however, potassium should not be stored under oil for longer than 6 months, unless in an inert (oxygen free) atmosphere, or under vacuum. After prolonged storage in air dangerous shock-sensitive peroxides can form on the metal and under the lid of the container, and can detonate upon opening.^[102]

Because of the highly reactive nature of potassium metal, it must be handled with great care, with full skin and eye protection and preferably an explosion-resistant barrier between the user and the metal. Ingestion of large amounts of potassium compounds can lead to hyperkalemia strongly influencing the cardiovascular system.^[103][104] Potassium chloride is used in the United States for executions via lethal injection.^[103]

Underwater Sodium - Periodic Table of Videos

Aqueous solutions[edit]

Sodium tends to form water-soluble compounds, such as halides, sulfates, nitrates, carboxylates and carbonates. The main aqueous species are the aquo complexes [Na(H₂O)_n]⁺, where n = 4–6.^[24] The high affinity of sodium for oxygen-based ligands is the basis of crown ethers; macrolide antibiotics, which interfere with Na⁺ transport in the infecting organism, are functionally related and more complex.^{[citation needed]}

Direct precipitation of sodium salts from aqueous solutions is rare because sodium salts typically have a high affinity for water; an exception is sodium bismuthate (NaBiO₃).^[25] Because of this, sodium salts are usually isolated as solids by evaporation or by precipitation with an organic solvent, such as ethanol; for example, only 0.35 g/L of sodium chloride will dissolve in ethanol.^[26] Crown ethers, like 15-crown-5, may be used as a phase-transfer catalyst.^[27]

Sodium content in bulk may be determined by treating with a large excess of uranyl zinc acetate; the hexahydrate (UO₂)₂ZnNa(CH₃CO₂)·6H₂O precipitates and can be weighed. Caesium and rubidium do not interfere with this reaction, but potassium and lithium do.^[28] Lower concentrations of sodium may be determined by atomic absorption spectrophotometry^[29] or by potentiometry using ion-selective electrodes.^[30]