Aluminum is a versatile metal. It’s lightweight, doesn’t rust, conducts electricity, and when mixed with small amounts of other metals, aluminum forms strong, durable “alloys.” But, all this comes with a price: contamination of the air with carbon dioxide. Here’s why:

Aluminum starts out as aluminum oxide. Electricity is used to remove the oxygen atoms, but electricity produced by coal-burning plants spews carbon dioxide into the air. One way to combat this is to generate electricity with hydroelectric plants. An even simpler way is to recycle aluminum — something we can all do! ♻️

🤓 Silicon, the second most common element in the earth’s crust behind oxygen, is a very versatile element.

Silicon alone is the element for semiconductors chips. When bonded to two atoms of oxygen, it forms “silica” — sand 🏖 . When sand is heated, it becomes transparent glass as it cools. Adding boron turns glass into heat-resistant Pyrex, while adding lead improves the way glass reflects light. When silicon is bonded to four atoms of oxygen as “silicate” forms quartz crystals, which become blue amethyst gems when atoms of cobalt are added, green amethyst from atoms of chromium, and amber amethyst from atoms of iron. Adding methane to silicate creates silicone, used in everything from silicone oil to silicone grease, silicone rubber, silicone caulk, and even hair conditioners!

Sulfur bonds with almost every element in the periodic table to produce an enormous variety of products, but none with the breathtaking history of Charles Goodyear’s fanatic search during the 1830’s for a way to prevent sap of the rubber tree from becoming gooey in warm weather. Overcoming abject poverty and the death of multiple children, Goodyear finally, in 1839, mixed sulfur with the rubber sap and accidentally spilled some on a hot stove. What he peeled off the stove would become the start ofthe rubber industry.

Unfortunately, sulfur released from coal-burning power plants creates sulfur dioxide gas which combines with molecules of water in the air to form sulfuric acid, one of the components of the “acid rain” that’s helping to disintegrate marble statues, limestone buildings and monuments, and even steel structures. Commercially produced sulfuric acid, though, is used to make fe lizers, lead batteries, and wide variety of other useful products.

By the way, other sulfur compounds, like hydrogen sulfide, account for the nasty smell emitted by skunks and bacteria in our intestines.

At room temperature, chlorine is a gas – a very poisonous gas. It was used during World War I, but quickly abandoned when unpredictable air currents caused the gas to drift back over the troops that released it. Because of its toxicity, chlorine is still used in bleach to kill bacteria and viruses.

Chlorine’s strength lies in its very strong attraction for electrons, particularly the dense concentration of electrons found in double bonds. Those double bonds help molecules absorb particular light rays and reflect the rest as colors. It is this breaking of those double bonds that allows chlorine bleaches to be used to remove stains in clothing and in woods to make paper bright white.

Bonded to other atoms, chlorine is generally quite safe. Sodium chloride, for example, is table salt, and hydrochloric acid, HCl, is secreted by the stomach to help digest our food. Polyvinylchloride (PVC) tubes and pipes, ammonium chloride in batteries, barium chloride in fireworks, and ammonium perchlorate to power spacecraft into space are just a few examples of useful chloride compounds.

Sir Humphry Davy named potassium in 1807 after isolating it by passing an electrical current through wood ash that had been boiled in large pots – so-called “pot ash” (which was mostly potassium carbonate). Potassium is so reactive that pure potassium does not exist in nature. The list of potassium compounds and their uses ranges from soaps to explosives, ceramics to baking soda, paint removers to insecticides.

Potassium is critical to life. That’s why potassium is one of the key elements in fertilizers along withnitrogen and phosphorus. In animals, potassium is needed for nerves to carry signals and muscles to contract, which is why potassium-rich foods like bananas, peanuts, and raisins are such good pre-game foods.

While calcium is known for building strong bones and teeth and hard shells for crustaceans, calcium has long been a vital building material for civilization. Calcium is mined from the ground as limestone, which is mostly calcium carbonate, mixed with some magnesium carbonate. Limestone is what marble
and chalk are made of, and also makes up the stalactites hanging from caves and dripping onto stalagmites pointing up from a cave floor.

Simply heating limestone drives off the carbon dioxide in calcium carbonate, leaving behind calcium oxide, called “quicklime.” Water added to quicklime turns it into “slaked lime” – calcium hydroxide. Slaked lime is used as a smooth paste to coat walls. Instead of just heating limestone, adding water and sand turns quicklime into mortar to bind rocks and other building materials together. Adding, in addition, iron and clay (a mixture of silica and aluminum), then heating the mixture into hard rocks and grinding it all into a powder creates “cement.” Then by adding water and rocks to cement, a slow chemical reaction begins that binds them all together into concrete.

Plaster of Paris is a close relative of limestone made up of calcium sulfate

Titanium is lightweight but very strong. It is commonly used in aircraft, spacecraft, missles, ships, crutches, golf clubs, and high-end bicycles.

Titanium is also used in artificial joints because it bonds so well with bone without being rejected by the body’s immune system.

Oceangoing ships built with titanium hulls are much better at resisting rust and corrosion, because oxygen in the air reacts with titanium to form a barrier of titanium oxide.

Painters use titanium dioxide because it refracts (bends) sunlight in so many directions that sunlight reflects back as bright white. Titanium dioxide’s ability to reflect light like this is also why it’s used in sunblockers to help prevent sunburn.

Vanadium is a lightweight metal mostly used to make very strong alloys of steel without adding a great deal of weight, useful for things like piston rods, crankshafts, and axes. One of its key features is the large number of orbitals around the nucleus for electrons to occupy and move about in. Not only does this make vanadium ideal for large industrial batteries, but with so many orbitals for electrons to occupy, vanadium can manifest many beautiful colors. That’s why vanadium is named after the Swedish goddess of beauty, Vanadis. Vanadium’s many electron states may help explain why shallow-water animals like ascidians and sea squirts absorb vanadium from the sea, but vanadium’s actual role in those animals is still unclear. Also unclear is why Amanita mushrooms — those pretty but poisonous mushrooms with the red caps dotted with white spots — use vanadium to make a complex molecule called “amavadin.”

Manganese and iron are often found together in the earth as manganese and iron oxides. Not surprisingly, these two oxides have been identified as the first pigments used in prehistoric cave paintings. Another early use of manganese was to remove the greenish tint from glass and make it clear.

By bonding to stray atoms of sulfur in iron, manganese is able to strengthen iron into steel alloys strong enough for rifle barrels, train tracks, and prison bars.

Manganese’s ability to give and receive electrons makes it a useful component of enzymes. For example, all the oxygen we breathe is a byproduct of photosynthesis produced by an enzyme that splits molecules of water into hydrogen and oxygen. That enzyme, called the “water-splitting enzyme,” depends on its four atoms of manganese.

Fungi feed off of dead wood by digesting the lignin in wood with its manganese-containing enzymes. We depend on our manganese-containing enzymes to protect us from dangerously reactive atoms of oxygen “radicals,” and to ensure the growth of healthy bones. Too much manganese, though, can damage the brain and lead to Parkinson’s disease. In fact, the early studies of L-DOPA treatment for Parkinson’s disease were carried out on manganese miners.

High concentrations of manganese can be found in nodules lying on the ocean floor, just waiting to be mined.

Iron is the metal for strength. It may rust easily and it may not be lightweight, but it is strong and can also be hammered, tempered, annealed, and forged into practically any shape. Iron can be mixed with carbon to make different varieties of carbon steel, which can then be alloyed with manganese, nickel, tungsten, titanium, copper, vanadium, or chromium to make a wide variety of strong, corrosion-resistant steel.

Even hemoglobin uses atoms of iron to carry oxygen around our bodies. The iron in hemoglobin is why our blood is red. Iron is so essential to our diet that it’s even incorporated into breakfast cereals – check it out.

Iron is also great for making magnets, which is good news for us because the central core of the earth is mostly iron, much of it in liquid form, and as it flows, it creates a magnetic field around the earth that protects us from the harmful electromagnetic rays of the sun.

Nickel is a transition metal that is incredibly resistant to oxidation and corrosion. Nickel can be mixed in with other metals to make strong corrosion-resistant alloys like stainless steel. The airplane and aerospace industries use nickel to make strong aluminum alloys. Nickel’s heat resistance produces alloys used in toasters and ovens. And don’t forget the nickel-copper alloy that makes up nickel coins! Nickel, like chromium, can also be plated onto other metals to produce a less shiny coat than chromium.

Nickel is also one of three metals that’s permanently magnetic, iron and cobalt being the other two, hence its use in rechargeable batteries.

Nickel and iron are often found together, especially in meteorites. One particularly large such meteorite crashed into Canada and around 1900, the English chemist Ludwig Mond discovered a way to extract the nickel. Mond found that by passing carbon monoxide gas though the mixture of iron and nickel to form nickel carbonyl, he could then gently heat nickel carbonyl into a gaseous vapor, which readily decomposed back into nickel and carbon monoxide. Great care must be taken handling nickel carbonyl, because once inhaled, both the carbon monoxide and the nickel are highly toxic and often fatal.

Copper has been mined and used for many thousands of years, because its relative softness allows it to be shaped into many different tools and utensils. Its real prominence began about 3300 B.C., when it was discovered that adding tin greatly strengthened copper into bronze.

The Stone Age quickly shifted to the Bronze Age, allowing civilization to begin building useful tools, especially agricultural tools. The Bronze age lasted until about 1200 B.C. when civilizations around the Mediterranean Sea began to collapse, and iron began to out-compete bronze.

Copper is still widely used, not just in our coins, but in many aspects of life. Its low cost and high conductivity of heat make it ideal for pots and pans. Its high conductivity of electricity makes it an important component of power plants, and its ductility allows copper to be pulled into thin electrical wires and shaped into delicate printed circuits.

Oddly enough, we need small amounts of copper in our diet to replenish key enzymes in our body, but more than that and copper turns poisonous. Various copper-containing insecticides, bactericidal agents, and antiviral agents use copper to poison insects, bacteria and viruses.

In 1869, before Gallium was discovered, Russian chemist Dmitri Mendeleev, the creater of the periodic table of elements, predicted gallium’s properties and its location below aluminum in the periodic table. Six years later, in 1875, gallium was discovered by a French investigator, Paul-Emile Locoq de Boisbaudran, who named it after his country’s Latin name, “Gallia”.

As a pure metal, gallium is soft and silvery, much like aluminum. Despite a rather low melting point of around 30 degrees Celsius, 86 degrees Fahrenheit — which suggests weak intermolecular bonding between gallium atoms — gallium has a very high boiling point – over 2400 degrees Celsius, almost 4400 degrees Fahrenheit. Its low melting point and high boiling point mean that gallium is liquid over a very wide range of temperatures.

Gallium’s usefulness is mostly confined to compounds of gallium, not pure gallium. Gallium arsenide and gallium nitride are both used to make semiconductors, much like silicon, only faster. They’re also used in solar cells to convert light into electricity, and in light-emitting diodes (LEDs) to emit light on receiving an electrical current.

Selenium is essential for the proper function of certain key enzymes in humans, but more than a few hundred micrograms a day in our diet is toxic.

Added to clear glass, selenium adds a reddish tint, but at the same time it’s able to neutralize the greenish tint seen in glass contaminated with iron atoms.
When teamed up with sulfur, right above it on the periodic table, selenium lends a particularly foul odor to skunk spray.

What’s most interesting about selenium, though, is the discovery around 1870 that shining a light on selenium causes it to generate an electrical current.

It took Albert Einstein to explain in 1905 that light rays carry packets of energy we now call “photons.” The discovery of the electron a few years earlier allowed Einstein to propose that light’s photons were providing selenium’s electrons with enough energy to escape from their atoms and flow through wires as an electrical current.

Because selenium created an electrical current on exposure to light, selenium was commonly used by photographers in hand-held light meters to measure levels of light. Selenium’s real success, though, was in the development of the xerox machine in the late 1930’s and 1940’s. The key structure in the early xerox machines was a metal plate coated with selenium. The selenium layer was made electrically positive by a high-voltage instrument called a “Corotron.” When a bright light was shined onto the page to be copied, photons from the white areas of the page were reflected onto the positively-charged selenium surface, releasing selenium’s electrons which then neutralized the positive charges. Since the black letters on the page being copied didn’t reflect any photons, those areas on the selenium surface remained electrically positive.

The next step was to dust the selenium surface with tiny, black, negatively-charged toner particles which adhered to the positive charges on the selenium layer. Since the areas of positive charge represented the black letters on the page being copied, simply laying a sheet of white paper onto the selenium- coated plate and heating the paper fused the toner into the paper for a permanent copy. 

Bromine is a halogen – along with fluorine, chlorine, and iodine. There is no known function for bromine in the human body, and in fact bromine is quite toxic when the brownish-red gas is inhaled.

At one time, bromine was considered helpful for extinguishing fires and preventing fires when incorporated into flammable products like cushions, bedding, and drapes. In other compounds, bromine was the key element for killing insects, nematodes, bacteria, fungi, and even rats. Bromine also had a role in a variety of beneficial pharmaceutical agents. Bromine’s vast protective role was sharply curtailed once bromine was discovered to be destroying the ozone layer at a much faster rate than even chlorine.

Our curtailment of bromine use hasn’t prevented the oceans from incorporating bromine into a variety of compounds that are still harvested from highly concentrated sea water, like the Dead Sea and deep wells holding natural brine.

Rubidium sits right below potassium and sodium in Group 1 of the periodic table. Like potassium and sodium, rubidium has a single valence electron that, because it sits in a high energy orbit far from the nucleus, is easily removed. This makes rubidium highly reactive. For example, rubidium, like the other Group 1 elements, reacts violently with water by replacing one of water’s hydrogen atoms to form an alkaline solution of rubidium hydroxide (RbOH). The released hydrogen atom bonds to another released hydrogen atom to form hydrogen gas.

When rubidium is heated, the excited valence electron gives off two frequencies of red color that are readily detected with a spectroscope. The red color is where rubidium got its name  – “rubidus”. The name is Latin for deep red, but its name has nothing to do with rubies, which are red because of chromium atoms.

Rubidium has very few commercial uses and no apparent role in plant or animal metabolism. However, because rubidium resembles potassium, which cells accumulate intracellularly, radioactive rubidium injected intravenously has been used to locate brain tumors. And that makes it pretty useful!

Strontium sits right below calcium in the periodic table and behaves in many ways like calcium. That’s why if strontium is ingested, it can be incorporated into our bones, especially the growing bones of children. This becomes a problem when the radioactive isotope strontium-90 is released into the atmosphere after a nuclear explosion or a nuclear accident like the ones that occurred at Chernobyl and Fukushima. The nuclear fallout enters the food chain by settling onto plants that are then consumed by farm animals. The danger of radioactive strontium-90 is that it can lead to bone cancer and damage the bone marrow’s ability to make blood.

The radioactivity emitted by strontium-90 consists of electrons that carry heat and conduct electricity, making strontium compounds a useful source of heat and electricity in isolated places like space vehicles.

Strontium-89 is also radioactive. When given to patients suffering bone cancer, the strontiuim-89 is taken up by the rapidly-growing cancer cells, which are then killed by the high dose of radiation.

Heating strontium atoms in a flame elevates its electrons to higher orbits. When the excited electrons return to a lower energy orbit, they give off a bright red color – hence the use of strontium compounds in fireworks and road flares. Certain strontium compounds glow in the dark because their excited electrons take their time to return to a lower energy orbit.