Alkaline earth metals

The alkaline earth metals are the six elements in Group 2 of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These metals form the second column of the periodic table, sitting just to the right of the highly reactive alkali metals of Group 1. Like the alkalis, each alkaline earth metal atom has two electrons in its outer shell and readily loses those electrons to form divalent cations (M²⁺) in compounds. This means their compounds are usually ionic salts in which the metal is present as a 2+ charged ion. For example, calcium forms Ca²⁺ in calcium chloride (CaCl₂) and magnesium forms Mg²⁺ in magnesium oxide (MgO).

The name “alkaline earth” comes from an old classification: centuries ago, chemists called certain oxide minerals “earths” if they were nonmetallic, water-insoluble, and stable to heat. The oxides of Group 2 metals fit that description and produced basic (alkaline) solutions, so they were termed alkaline earths. When scientists later isolated the metals from these oxides, the metals themselves became known as the alkaline earth metals. All the Group 2 metal oxides form alkaline solutions in water (for instance, calcium oxide produces calcium hydroxide, a strong base). An interesting exception is beryllium oxide, which unlike its heavier congeners, is amphoteric – it can behave as either a base or an acid. In general, as you go down the group from beryllium to barium, the metals become more reactive and more strongly metallic (electropositive). Beryllium is relatively uncommon and has some non-typical behaviors, whereas barium is quite reactive (it must be stored under oil like sodium) and forms very strong bases such as barium hydroxide. Radium, at the bottom of the group, is highly radioactive, which dominates its chemistry and reactivity.

Physically, the alkaline earths are shiny, silvery metals, somewhat harder and denser than the alkali metals in the adjacent column, but still soft and lightweight compared to many transition metals. They are never found in pure form in nature because of their reactivity – instead, they occur in minerals and ores throughout the Earth’s crust. In fact, magnesium and calcium are among the most abundant elements on Earth (calcium is the 5th most abundant element in Earth’s crust, and magnesium the 8th). Calcium and magnesium compounds are ubiquitous: calcium carbonate makes up limestone and sea shells, magnesium salts permeate the oceans (imparting a bitter taste to seawater), and both elements are essential to all living organisms. Strontium and barium are less common but still moderately abundant (strontium is about the 15th most abundant element in the crust), found in minerals like celestite (SrSO₄) and barite (BaSO₄). Beryllium is rarer, found in certain gemstones and ores (like beryl and bertrandite), and radium is exceedingly rare – it exists only as a trace radioactive byproduct in uranium ores.

All these metals are reactive, tending to form compounds by giving up their two valence electrons. When heated, they react with oxygen to form oxides, and with halogens to form halide salts. Most also react with water, at least when the water is warm or the metal is powdered: for instance, calcium will slowly produce bubbles of hydrogen in cold water, and barium reacts readily even with room-temperature water to yield barium hydroxide and hydrogen gas. (Beryllium is the exception – it does not react with water due to a protective oxide layer on its surface.) Because they tend to form colorless ionic compounds, many alkaline earth salts are white solids or give colorless solutions. However, one spectacular signature is their flame colors: calcium and strontium compounds impart brilliant red colors to flames, while barium yields a green flame. This makes strontium and barium especially useful in fireworks.

From a human perspective, the alkaline earth metals play diverse roles in science, industry, and everyday life. They literally form the infrastructure of the modern world: calcium in concrete, plaster, and our very bones; magnesium in aluminum alloys, aircraft frames, and even the chlorophyll in green plants; strontium and barium in fireworks, ceramics, and medical imaging; beryllium in high-tech aerospace and electronics applications; and radium, historically in luminous paints and early cancer treatments. In the sections below, we’ll take a brief tour of each of these six elements, highlighting their key characteristics, interesting facts, and how they touch our lives.

Summary of Group 2 Elements: Before diving in, the table below summarizes the six alkaline earth metals with their symbols, atomic numbers, and one notable use or property of each:

Element	Symbol	Atomic Number	Notable Use or Property
Beryllium	Be	4	Very lightweight, stiff metal used in aerospace structures and X-ray windows.
Magnesium	Mg	12	Lightest structural metal; used in alloys (e.g. aircraft parts) and bright-burning flares.
Calcium	Ca	20	Essential for bones and teeth; major component of limestone, concrete, and seashells.
Strontium	Sr	38	Gives brilliant red color in fireworks and flares; used in glow-in-the-dark paints and old TV tubes.
Barium	Ba	56	Dense, opaque barium sulfate is used in X-ray contrast “barium meals”; also gives green fireworks color.
Radium	Ra	88	Highly radioactive; glows in the dark and was once used in luminous watch dials (now handled only with strict precautions).

Beryllium (Be)

Main article: Periodic table/Alkaline earth metals/Beryllium

Beryllium is a steel-gray metal, very light in weight but with a remarkable stiffness (its modulus of elasticity is about one-third greater than steel’s). This rare metal is quite brittle at room temperature and toxic to handle, but its unique properties make it invaluable in specialized applications. Beryllium has a high melting point for such a lightweight metal, and it resists corrosion by forming a thin, tough oxide layer on its surface. As a result, it remains stable in air and at normal temperatures will not react or tarnish quickly. Early chemists actually noted that some beryllium compounds had a sweet taste – the element was first called glucinium (from Greek glykys, “sweet”) – but we now know beryllium is highly toxic, especially if dust or fumes are inhaled, so one should never taste or ingest it. In fact, beryllium and its salts are classified as carcinogenic and must be handled with special precautions in industry.

Despite the hazards, beryllium’s advantages are significant. It is about 1.85 times as dense as water (lighter than aluminum) but exceptionally rigid, and it has excellent thermal conductivity. These traits, plus the ability to form a strong, passivating oxide, make beryllium useful as a structural material in high-performance aerospace and military applications. For example, beryllium is used in satellite and spacecraft components, high-speed aircraft, precision instruments, and even the mirrors of advanced telescopes, where low weight and dimensional stability are crucial. Beryllium is also transparent to X-rays – it transmits X-ray beams much better than heavier metals (about 17 times better than aluminum) – so thin beryllium foil or windows are used in X-ray tubes and scientific instruments to allow X-rays to pass through.

Much of the beryllium in commerce is not pure metal but in alloys. Notably, adding about 2% beryllium to copper produces beryllium-copper alloy, which is six times stronger than pure copper and has the valuable property of being non-sparking. Beryllium-copper tools (wrenches, hammers, etc.) are used in oil rigs, grain silos, and munitions factories where a stray spark could trigger an explosion. Small additions of beryllium to other metals can also improve their performance; for example, a bit of beryllium in magnesium alloys helps form a protective oxide skin that reduces flammability.

In nature, beryllium does not exist as free metal. It is found in certain minerals, most famously beryl (a beryllium aluminum silicate). Gem-quality beryl gives us emerald and aquamarine – the green and blue varieties of this mineral. In fact, emerald was known and treasured by ancient Egyptians, but it wasn’t realized until the late 18th century that emerald’s composition was similar to ordinary beryl. Beryllium was first isolated in 1828 by Friedrich Wöhler and Antoine Bussy (after its oxide was discovered by Louis Nicolas Vauquelin in 1798). Today, the United States, China, and a few other countries produce beryllium, extracting it mainly from minerals like bertrandite and beryl.

One historical scientific footnote: beryllium played a key role in the discovery of the neutron. In 1932, James Chadwick bombarded a beryllium target with alpha particles (from a radium source) and observed a new kind of penetrating radiation – neutrons – being emitted. This experiment was crucial for nuclear physics, and it was made possible by the unusual nuclear properties of beryllium when hit with alpha rays.

In summary, beryllium is a niche metal – you won’t encounter it in daily life except perhaps as a component inside your cell phone or in aerospace hardware – but it enables technologies that demand a combination of light weight, rigidity, and thermal stability. Just remember: its old nickname “sweet” metal is a misnomer, and tasting or breathing beryllium is definitely off-limits for health reasons!

Magnesium (Mg)

Main article: Periodic table/Alkaline earth metals/Magnesium

Magnesium is a shiny, silvery-white metal that is both extremely common on Earth and extraordinarily useful. It makes up about 2% of the Earth’s crust by weight, occurring in many minerals such as dolomite (calcium magnesium carbonate), magnesite (magnesium carbonate, shown in the image below), and various silicates. Yet you will never find magnesium as a free metal nugget in nature – it’s always locked in compounds. Sir Humphry Davy first isolated pure magnesium in 1808 by electrolyzing molten magnesium oxide, and French chemist Antoine Bussy later produced it in quantity. The name “magnesium” comes from Magnesia, a region in Greece where magnesium-rich minerals were known since antiquity (for example, “milk of magnesia,” a suspension of magnesium hydroxide, got its name from the same root).

Magnesite – crystallized magnesium carbonate – is one of the common minerals that contain magnesium. Pure magnesium metal is not found free in nature, but magnesium is abundant in compounds like magnesite and dolomite.

One of magnesium’s claims to fame is that it’s the lightest structural metal – among the metals that are used for building things, magnesium has the lowest density (about one-quarter the density of steel). In fact, solid magnesium can even float on very salty water. This low weight, combined with decent strength (especially when alloyed), has made magnesium a go-to component in lightweight alloys. Magnesium is often mixed with aluminum, zinc, or manganese to create alloys that are used in automobile parts, aircraft bodies, spacecraft components, and portable tools. For example, some car engine blocks and aircraft gearboxes are made from magnesium-aluminum alloys to save weight. Older examples include the Volkswagen Beetle’s engine block, which was largely magnesium, and military aircraft parts during World War II. Magnesium alloys are easy to machine and fabricate, which is another plus for manufacturers.

Magnesium metal has another dramatic property: it burns with a brilliant white flame. Finely divided magnesium ignites easily and produces an intense white light. This led to its use in early photographic flash powders and flashbulbs in the 19th–20th centuries. Even today, magnesium is used in flares, fireworks, and emergency signal fires, where its bright light (and hot flame) is advantageous. Military incendiary bombs in WWII were packed with magnesium as well, to produce high-temperature fires. A simple classroom demonstration is burning a strip of magnesium ribbon – it will flare with an eye-searing white glow and leave a flaky white ash of magnesium oxide.

Chemically, magnesium is reasonably reactive but not as explosively so as the alkali metals. A piece of magnesium metal will tarnish slowly in air as a thin oxide coat forms, but it won’t react violently unless ignited. It reacts with acids (like hydrochloric acid) to produce hydrogen gas, and it will react with water very slowly at room temperature. In hot water or steam, however, magnesium reacts more rapidly, yielding magnesium hydroxide and hydrogen. In comparison to its Group 2 neighbors, magnesium is less reactive than calcium or barium, partly because the protective oxide layer sticks well to magnesium’s surface. Magnesium’s compounds tend to be white solids (for instance, milk of magnesia is a white suspension). A fun fact: dissolved magnesium salts in seawater contribute to the ocean’s bitter taste – seawater contains about 0.13% magnesium (mostly as magnesium chloride), which is a key factor in “sea salt” flavor.

One of the most important aspects of magnesium is its biological role. Magnesium is absolutely essential to life. It is a central component of chlorophyll, the green pigment in plant leaves that drives photosynthesis: at the heart of every chlorophyll molecule sits a magnesium ion, coordinating a complex ring structure that allows plants to capture sunlight. Without magnesium, plants couldn’t produce glucose from carbon dioxide and water, and the food chain would collapse. In the human body, magnesium is the fourth most abundant mineral. It’s critical for hundreds of enzymes, muscle and nerve function, and bone health. Humans get magnesium from foods (like leafy greens, nuts, and grains), and a magnesium deficiency can cause muscle cramps and other health issues. On the flip side, magnesium compounds have medical uses: Epsom salt is magnesium sulfate, used to soothe muscle aches, and milk of magnesia is magnesium hydroxide, a common antacid and laxative.

From a broader perspective, magnesium even has an extraterrestrial significance: it is forged in the hearts of stars from the fusion of lighter elements, and it’s the eighth-most abundant element in the universe. Here on Earth, it’s the third most abundant element dissolved in seawater (after sodium and chloride ions). So the next time you see a bright firework or take a dose of antacid, remember magnesium – a metal that not only lights up the sky but also powers the biology of life itself.

Calcium (Ca)

Main article: Periodic table/Alkaline earth metals/Calcium

If one element can claim to be the building block of both civilization and life, it’s calcium. Calcium is all around us: in limestone cliffs, concrete buildings, chalk, marble statues, bones and teeth, seashells, and even the milk we drink. It’s no surprise that calcium is extremely abundant. In fact, calcium seems to come in fifth place wherever we look: it’s the 5th most abundant element in Earth’s crust (after oxygen, silicon, aluminum, and iron), the 5th most abundant dissolved ion in seawater (after sodium, chloride, magnesium, and sulfate), and the 5th most abundant element in the human body by mass (after oxygen, carbon, hydrogen, and nitrogen). It is the most abundant metal in our bodies – about 2 pounds (1 kg) of calcium is in an average adult, almost all of it locked up in our bones and teeth as calcium phosphate.

Calcium is a soft, silvery metal in pure form, but you’ll almost never see pure metallic calcium in everyday life. It’s far too reactive to exist freely. If freshly cut, calcium has a shiny silvery-white appearance, but exposed to air it quickly turns dull gray by forming calcium oxide and calcium nitride on the surface. Pure calcium metal will even react (slowly) with the moisture in air. It was first isolated in 1808 by Humphry Davy (the same pioneer who isolated sodium, potassium, and magnesium via electrolysis). Davy produced calcium by electrolyzing a molten mixture of lime (CaO, derived from limestone) and mercuric oxide. The name “calcium” comes from calx, the Latin word for lime (calcium oxide), which was known and used by ancient civilizations. Indeed, compounds of calcium were in use long before calcium metal was purified: the Romans built structures with lime and volcanic ash, and medieval manuscripts describe the use of plaster of Paris (calcium sulfate) to set broken bones as early as the 10th century.

While pure calcium metal has few direct uses (it tarnishes quickly and is a bit too reactive, though it’s sometimes used as a reducing agent to extract other metals from ores), calcium compounds are enormously important. Calcium carbonate (CaCO₃) is one of the most common substances on Earth – it forms limestone rock, marble, chalk, and coral reefs. We burn limestone to create lime (CaO) for cement and mortar; when water is added, lime becomes calcium hydroxide, which hardens by reacting with CO₂ from the air, returning to carbonate in concrete. The Great Pyramids of Giza and the Roman Colosseum are enduring testaments to calcium-based building materials. Calcium carbonate is also used in agriculture to neutralize acidic soils (as lime) and is the main ingredient in antacid tablets (usually as chalk or powdered limestone). Gypsum (calcium sulfate dihydrate) is another crucial calcium compound – when heated it becomes plaster of Paris, used for casts and drywall. Calcium fluoride (CaF₂), or fluorite, is used to make specialized lenses and windows for microscopes and telescopes because it’s transparent to ultraviolet and infrared light.

In water, calcium compounds are responsible for what we call “hard water.” As groundwater flows through limestone or gypsum deposits, it dissolves calcium ions. Hard water isn’t harmful – in fact, it can contribute to dietary calcium – but it leaves limescale deposits in kettles and pipes (that white crust is largely calcium carbonate). Interestingly, brewers pay attention to calcium levels in water, as it can affect the taste of beer and the brewing process.

Biologically, calcium is indispensable. It is classified as a “macromineral” for nutrition because we need it in substantial amounts. Calcium ions (Ca²⁺) in our bodies are what make our bones and teeth hard – they combine with phosphate to form hydroxyapatite, the mineral matrix of bones and enamel. But calcium’s role goes beyond structural; dissolved calcium ions in our cells and blood serve as vital signals for physiological processes. For example, a sudden influx of Ca²⁺ into a muscle cell is the trigger for contraction, and Ca²⁺ in neurons helps transmit impulses. Our blood won’t clot without calcium (it’s a co-factor in many enzymes). That’s why maintaining calcium levels (through diet or supplements) is important, especially for children, pregnant women, and the elderly. Vitamin D and hormones help regulate calcium by ensuring it gets absorbed from food and deposited in bones.

Calcium in everyday life often comes up in the context of diet (“drink your milk for strong bones”) and health, but it’s also literally under our feet and in the walls around us. The minerals of calcium have built human infrastructure for millennia – limestone for construction, gypsum for plaster, cement for modern cities – and they continue to do so. It’s poetic that the same element forming the skeleton of a skyscraper is also forming the skeleton in your body. When you admire a stalactite in a cave, a seashell on the beach, or even the concrete skyline of a city, you’re seeing calcium’s handiwork. This versatile element truly earns its title as “nature’s most renowned structural material”, integral to both the living world and the built environment.

Strontium (Sr)

Main article: Periodic table/Alkaline earth metals/Strontium

Strontium is a soft, silvery metal that doesn’t often make headlines, but it does make fireworks sparkle red and glow-in-the-dark paint shine. In pure form strontium is highly reactive – even more reactive than calcium – so like its neighbors, it isn’t found free in nature. Instead, strontium occurs in minerals such as celestite (strontium sulfate) and strontianite (strontium carbonate). The element itself was named after Strontian, a village in Scotland, where strontianite ore was first identified in the late 18th century. Pure strontium metal, isolated by Sir Humphry Davy in 1808 (using electrolysis, as he did for other alkali and alkaline earth metals), is so reactive that it must be stored under kerosene to keep it from oxidizing. Fresh strontium quickly turns yellow in air as it forms oxide, and if heated it burns with a bright crimson flame. When strontium metal or strontium salts are added to a flame, they emit a brilliant red color – that’s why strontium compounds (like strontium nitrate or strontium carbonate) are used to produce red colors in fireworks and signal flares.

Sky-blue celestite crystal, a primary mineral source of strontium. Strontium derived from celestite is responsible for the brilliant red hues in fireworks and flares. Natural strontium is stable and harmless, but radioactive strontium-90 from nuclear fallout is hazardous.

One of strontium’s notable uses in the past was in television cathode ray tubes (CRTs). The glass screens of old color TVs and computer monitors contained strontium and lead compounds. Strontium (as strontium oxide) was added to the glass to block X-ray emissions from the electron gun and to improve the optical properties of the glass. With the decline of CRTs in favor of flat panels, this use has faded, but it was significant in the mid-20th century.

Strontium today finds niche uses in materials and medicine. Certain glow-in-the-dark phosphorescent paints use strontium compounds (like strontium aluminate) as the phosphor; these are the coatings that absorb light and then slowly release it, glowing in the dark. If you have seen those glow-in-the-dark stars for bedroom ceilings or emergency exit signs that luminesce after lights go out, you’ve seen strontium at work. Another use is in ceramics and glass for ferrite magnets and in refining zinc. Strontium titanate, a man-made gemstone, has an extremely high refractive index and was once sold as a diamond simulant (though it’s much softer than diamond).

A particularly interesting application of strontium is in science and anthropology: strontium isotope analysis. Different regions of the world have slightly different ratios of strontium isotopes in the soil and water (due to geological differences). These isotopes end up in plants and animals. By measuring the strontium isotope ratios in human or animal bones and teeth, scientists can often infer where that individual grew up or migrated. Archaeologists use this technique to track ancient human migrations or to identify the birthplace of skeletal remains – your teeth, which form in childhood, carry a strontium “signature” of the local geology where you lived as a child.

It’s worth noting that natural strontium, comprised of stable isotopes, is non-toxic and not radioactive – it behaves a bit like calcium in the body, and about 99% of the strontium in our bodies resides in our bones (harmlessly replacing a small fraction of calcium). However, strontium-90, a radioactive isotope produced in nuclear fission (like in nuclear bomb fallout or reactor accidents), is a serious health hazard. Strontium-90 has a half-life of around 29 years and chemically mimics calcium, so if ingested, the body can deposit it in bones and teeth. There it irradiates tissues and can cause bone cancer or leukemia. During the Cold War era, strontium-90 from atmospheric nuclear tests became a global concern; it was one of the reasons above-ground testing was banned. Fortunately, strontium-90 is not found in nature, only as a byproduct of human nuclear activities.

In summary, strontium is a bit of a quiet workhorse among the alkaline earths. You likely encounter it in the form of a brilliant red firework explosion or the persistent glow of a novelty sticker. And if you’ve ever had a toothpaste for sensitive teeth containing strontium chloride (an older formulation), you’ve even put strontium in your mouth – some sensitive-teeth toothpastes used strontium compounds to help block tooth tubules and reduce pain. Rest assured, the strontium in consumer products is the stable kind. So, while strontium doesn’t have the marquee name of calcium or magnesium, it adds color to our celebrations and insight to our science, proving its worth as a member of the Group 2 family.

Barium (Ba)

Main article: Periodic table/Alkaline earth metals/Barium

Barium is the heavyweight of the group in more ways than one. Its name comes from the Greek barys, meaning “heavy,” because barium compounds were noted for their high density. Barium’s most common ore, barite (barium sulfate), is a dense mineral – so heavy that it was used in old drilling muds to weight them down. Pure barium metal itself, isolated by Humphry Davy in 1808, is a silvery-white, soft metal that is extremely reactive. In fact, barium reacts so readily with oxygen and moisture that it must be stored under oil or in an inert atmosphere. When exposed to air, barium metal will tarnish to a dark gray as it oxidizes, and if ignited, it burns with a pale green flame – a signature used in fireworks to produce green colors.

Barium sits toward the bottom of the periodic table’s main block, giving it a large atomic size and a lot of “heft” (atomic weight ~137). It forms only a 2+ ion in chemistry and typically exists in the +2 oxidation state in all its compounds, similar to the rest of the alkaline earths. Because of its reactivity, barium is never found in elemental form in nature, only as compounds. The two primary barium minerals are barite (BaSO₄) and witherite (BaCO₃). Barite is quite inert due to its low solubility, whereas witherite (barium carbonate) will react with acids and is somewhat toxic.

Despite (or rather, because of) its high reactivity, barium has a number of valuable commercial uses in compound form. Barium sulfate, in particular, is amazingly useful because it is so insoluble and dense. A fine suspension of barium sulfate is used as an X-ray contrast agent for imaging the human gastrointestinal tract. If you’ve heard of a “barium meal” or barium enema used in medical diagnostics, this refers to patients ingesting or being administered a slurry of barium sulfate before an X-ray. The heavy barium atoms absorb X-rays strongly, coating the outline of the stomach or intestines on an X-ray image, but since BaSO₄ doesn’t dissolve, the barium passes through the body without poisoning the patient. (It’s crucial that only insoluble barium sulfate is used for this; soluble barium compounds would release barium ions into the body, which are highly toxic.)

Barium sulfate’s combination of whiteness and heft also led to its use in materials. It is a component of the pigment lithopone, a brilliant white paint pigment (lithopone is a mixture of BaSO₄ and zinc sulfide) used historically in paints and inks. Because barite is so dense (around 4.5 g/cc), it was and still is used in oil and gas drilling fluids: ground barite is added to the drilling “mud” to make it heavy, which helps counter high underground pressures and prevent blowouts. Barite is also used in the production of high-quality glass, ceramics, and rubber, where its chemical inertness and weight are advantageous.

Other barium compounds have their niches. Barium nitrate (Ba(NO₃)₂) and barium chlorate (Ba(ClO₃)₂) are used in pyrotechnics to produce the vibrant green colors in fireworks and signal flares. Barium carbonate (witherite) has seen use as a rat poison (unfortunately, as it can be effective – it dissolves in the acid of a rodent’s stomach to release toxic Ba²⁺ ions). Barium titanate is a noteworthy electronic material – it’s a ferroelectric ceramic used in capacitors, sensors, and electromechanical devices due to its ability to change shape under an electric field and store electrical charge.

On the topic of toxicity: soluble barium salts are poisonous. Barium ions interfere with the proper function of muscles (including the heart) by affecting potassium ion channels. Ingesting a soluble barium compound can cause violent nausea, abdominal pain, and potentially fatal disturbances in heart rhythm. Historically, there have even been a few murders and accidental deaths via barium compounds. Thankfully, barium poisoning is rare, and in everyday life one does not encounter soluble barium – the insoluble sulfate is the common form. Our environment contains trace amounts of barium (the average adult has about 22 mg of barium in their body, mostly from food and water, but it remains in insoluble forms and serves no biological function). In fact, scientists can examine the barium in human teeth to learn about infant nutrition: a baby’s teeth incorporate barium in different amounts before and after weaning, so by analyzing barium levels across a tooth, researchers can tell when an infant transitioned to solid food – a fascinating application of analytical chemistry in anthropology.

In summary, barium is an element that we typically encounter in compound form rather than as a metal. It’s the secret behind certain paints, drilling fluids, and the life-saving contrasts on medical X-rays. And while it can be dangerous in the wrong form, in the right form – a dose of dense, chalky barium sulfate – it quite literally helps doctors see inside us. Just remember that the flashy green fireworks on holidays have a less flashy side in the lab and hospital, where barium quietly does its job as a heavy, useful helper.

Radium (Ra)

Main article: Periodic table/Alkaline earth metals/Radium

Radium is the last and heaviest alkaline earth metal, and it stands apart from the others for one key reason: intense radioactivity. In fact, radium is over a million times more radioactive than the same amount of uranium. This element was discovered in 1898 by Marie and Pierre Curie, who extracted it from the uranium ore pitchblende – a painstaking process since radium is present in uranium ores only in trace amounts. The Curies named the element radium from the Latin radius, meaning “ray,” because it emitted powerful rays of energy. Radium constantly undergoes radioactive decay, emitting alpha particles, beta particles, and gamma rays, and it glows with a faint blue light due to the intense radiation exciting the air around it. All isotopes of radium are radioactive; the most stable isotope, radium-226 (the one the Curies isolated), has a half-life of about 1600 years. This means any radium that was present when the Earth formed is long gone – today’s radium is continuously generated as a decay product of uranium and thorium. Radium-226 is produced in the decay chain of uranium-238, so wherever you find uranium, tiny bits of radium are being formed (and further decaying into radon gas, lead, and other elements).

In nature, radium is extraordinarily scarce. Its abundance in Earth’s crust is on the order of 1 part per trillion. Essentially, it’s only found mixed in uranium or thorium minerals, and even there in minute quantities. There are no “radium ores” as such; historically radium was obtained by processing tons of uranium ore residues to isolate just milligrams of radium – a testament to the Curies’ dedication that they managed to isolate it at all.

Radium’s properties are in some ways typical of an alkaline earth (for instance, it forms a 2+ ion, and radium compounds such as radium chloride and radium sulfate resemble the corresponding barium compounds), but its radioactivity dwarfs its normal chemical behavior. For example, radium metal is silvery and will react with air and water like barium does, forming radium oxide and radium hydroxide; however, handling enough radium metal to observe those reactions is impractical and dangerous due to radiation. Notably, radium’s intense radioactivity causes it to self-heat (a sample will keep itself warm) and to luminesce (glow) by activating phosphors.

In the early 20th century, radium was treated like a miraculous new substance. People were mesmerized by its glow and its mysterious “rays.” This led to some unfortunate uses of radium in everyday products – before the dangers were fully understood, radium was mixed into things like luminous paints, clock dials, even health tonics and toothpaste. Radium-based paint was used to make watch and clock hands and numerals glow in the dark, which was very useful before battery-powered lights or tritium existed. Factories hired young women to paint watch dials with radium-laced paint; these workers, later known as the "Radium Girls", would often lick their brushes to keep a fine point, unwittingly ingesting radium in the process. Over time, many of them developed radiation poisoning – bone deterioration, lesions, cancers – as the radium deposited in their bones delivered destructive radiation internally. Their plight in the 1920s became a landmark case in occupational health and led to improved safety standards and the eventual halt of radium dial painting. Radium was also marketed in quack “health” products: for example, radium water (water with trace radium) was sold as a rejuvenating tonic. One infamous case involved a wealthy industrialist, Eben Byers, who drank so much radium-laced water (believing it boosted vitality) that his jaw fell apart from necrosis, and he died in 1932 – a cautionary tale that finally dampened the radium craze.

Today, we recognize that radium is extremely hazardous. The intense radiation from radium (mostly from radium-226 and its decay products like radon gas) can cause cancer and radiation sickness. Marie Curie herself suffered health effects from handling radium – she eventually died from aplastic anemia likely caused by radiation exposure. Because of this, radium no longer has consumer applications, and strict protocols surround its use. The luminous paints were replaced first by less harmful phosphorus-based glow paints, and later by tritium (a radioactive hydrogen isotope that emits much weaker radiation).

However, radium did have one very positive application: it was the first element used in cancer radiation therapy. The gamma rays emitted by radium could kill cancer cells, so in the early 1900s doctors used radium in sealed needles or tubes as a treatment for tumors (a technique called brachytherapy). Radium-226 was implanted or applied to the tumor site to irradiate it. While this did sometimes work, it also exposed medical staff to radiation, and the supply of radium was limited and extremely costly. By the mid-20th century, radium was largely supplanted by isotopes like cobalt-60 and cesium-137, which are easier to produce and handle. Interestingly, a modern resurgence has occurred in the form of radium-223 (trade name Xofigo), a short-lived isotope used to treat prostate cancer that has spread to bones – it delivers high-energy alpha particles to bone tumors and then decays away quickly.

In summary, radium is a fascinating but daunting member of the alkaline earth family. Chemically, it behaves as a heavier cousin of barium, but its radioactivity dominates any discussion. In the span of a few decades, radium went from a celebrated wonder element lighting up watch dials and touted as a cure-all, to a notorious killer and symbol of radioactive danger. Its story is intertwined with the birth of nuclear physics and the implementation of safety regulations in industry. If calcium is the element of life, radium is the element of lethal rays. Today, aside from specialized medical uses and research, radium is kept out of the public sphere – a glowing reminder of why understanding chemistry (and heeding caution) is so important.

Conclusion

From the sturdy calcium in our bones and buildings to the brilliant magnesium lighting up fireworks, from the quiet strontium in glow-in-dark paints to the heavy barium revealing our inner organs on X-rays, the alkaline earth metals are truly a diverse bunch. They illustrate how elements in the same group can share similar chemistry – forming 2+ ions, creating alkaline solutions, imparting flame colors – yet serve wildly different roles in our world. These Group 2 metals might not be as notorious as the explosive alkali metals or as famous as the noble gases, but they are no less significant. In fact, they are woven into the fabric of daily life and human progress: essential nutrients, industrial workhorses, and even historical curiosities. The tale of the alkaline earths stretches from the sweet shimmer of beryllium in gemstones to the eerie glow of radium in a dark room, underscoring the beauty and complexity of the periodic table. Each of these metals, in its own way, continues to impact science, technology, and life every day – a testament to the rich chemistry of the alkaline earth family.

References: The information in this article is drawn from reputable sources including the Encyclopædia Britannica for historical and chemical details, chemistry textbooks and databases for properties of Group 2 elements, and educational resources such as Live Science for interesting facts and applications of each element. The summary table and specific examples are supported by these sources to ensure accuracy and up-to-date context.