Period 4

Period 4 of the periodic table refers to the fourth horizontal row of chemical elements, starting with potassium (atomic number 19) and ending with krypton (atomic number 36). Unlike the shorter periods above it, Period 4 contains a total of 18 elements – one for each of the 18 vertical groups in the table. This is the first period long enough to include the d-block of the periodic table, meaning it introduces the first row of transition metals. In other words, after the two light metals at the beginning of the period, a series of ten transition metals appears, before the period concludes with six elements in the p-block (on the right side of the table).

From an atomic structure perspective, the elements in Period 4 have electrons populating their fourth electron shell (energy level). The first two elements (potassium and calcium) fill the 4s orbital, and then the next ten (scandium through zinc) progressively fill the 3d orbitals while still having two electrons in 4s. After zinc, the 3d sub-level is fully filled (effectively becoming part of the inner electron core), and the remaining six elements (gallium through krypton) fill the 4p orbitals as their valence electrons. By the end of the period with krypton, all electron shells inside the fourth are completely filled – a situation unique to this period, as later periods involve additional inner orbitals (like f orbitals) that prevent complete filling of all lower shells. Period 4 is also notable as the last period in which every element is stable (non-radioactive); all 18 Period 4 elements have at least one stable isotope. Starting in Period 5, elements like technetium have no stable forms. Many Period 4 elements are in fact very common in the Earth's crust or core (for example, iron is extremely abundant in the Earth's core).

Chemically, Period 4 is a microcosm of the diversity of the periodic table. It begins on the far left with two highly reactive metals (an alkali metal and an alkaline earth metal), then spans a block of typical metals (the transition metals), and ends on the far right with a mix of metalloids and nonmetals, including a halogen and finally a noble gas. Because of this, the properties of Period 4 elements vary widely: from very soft, reactive metals to hard structural metals, from toxic substances to essential nutrients, and from solids to gases. There are periodic trends observable across the period – for instance, metallic character generally decreases from left to right, while electronegativity and ionization energies generally increase – but the presence of the transition metals causes some irregularities in these trends. Still, elements that lie directly above or below each other in the periodic table (for example, potassium above rubidium, or selenium below sulfur) show familial resemblances in their chemistry, illustrating the periodic law.

Another point of interest is how integral many Period 4 elements are to life and technology. Several of these elements play vital roles in biological systems or everyday materials. For example, calcium is crucial for bones and teeth, potassium and sodium are essential electrolytes in our nerves, and iron is necessary in blood hemoglobin. On the other hand, some Period 4 elements are famously toxic – arsenic and bromine, for instance, have poisonous effects. The period includes elements known since antiquity (like copper and iron) as well as others that were discovered in more recent history once the periodic table predicted their existence (gallium and germanium were predicted by Mendeleev before being found). Period 4 also contains many of the metals that enabled the Industrial Revolution and modern technology (iron, copper, nickel, etc.), as well as elements critical for electronics (such as germanium and gallium in semiconductors). In short, this row of the periodic table is extraordinarily rich in chemical and practical significance.

Period 4 Elements at a Glance

To give a quick comparison, the table below summarizes each of the 18 Period 4 elements, listing the name, symbol, atomic number, and one key use or notable property of each:

Element (Name)	Symbol	Atomic Number	Key Use or Notable Property
Potassium	K	19	Essential plant nutrient (major component of fertilizers)
Calcium	Ca	20	Builds strong bones and teeth (critical structural mineral in animals)
Scandium	Sc	21	Used in small amounts to strengthen aluminum alloys (e.g. in aerospace components)
Titanium	Ti	22	Strong, lightweight, and corrosion-resistant metal widely used in aircraft and medical implants
Vanadium	V	23	Added to steel alloys to increase strength (used in tools, axles); also used in large-scale batteries (vanadium redox flow batteries)
Chromium	Cr	24	Hard, shiny metal used for stainless steel and chrome plating; many compounds are vividly colored (name derives from Greek chroma, color)
Manganese	Mn	25	Important in steelmaking (prevents rust and adds hardness); also used in dry cell batteries (manganese dioxide cathodes)
Iron	Fe	26	Principal component of steel (the backbone of construction and manufacturing); also essential in blood (hemoglobin protein)
Cobalt	Co	27	Key ingredient in superalloys and magnets; cobalt compounds provide a brilliant blue pigment (cobalt blue) and Co is a trace element in vitamin B₁₂
Nickel	Ni	28	Corrosion-resistant metal used in stainless steel and coinage; majority of Earth’s core (with iron) is thought to be nickel-iron metal
Copper	Cu	29	Excellent electrical conductor widely used in wiring and motors (about 70% of global copper is used in electrical applications)
Zinc	Zn	30	Used to galvanize steel (protective zinc coatings to prevent rust) and in brass alloy (with copper); also an essential trace nutrient for the immune system
Gallium	Ga	31	Metal that melts at about 30 °C (it can melt in your hand); used in semiconductors and LEDs (gallium arsenide, gallium nitride)
Germanium	Ge	32	Metalloid used as a semiconductor (early transistors, fiber-optic systems); its existence was predicted as “eka-silicon” before its discovery
Arsenic	As	33	Poisonous metalloid historically used in pesticides and pigments; also used in semiconductor compounds (e.g. gallium arsenide)
Selenium	Se	34	Used in glassmaking (to remove color or create red glass) and was once common in photocopiers and light sensors; required in tiny amounts in nutrition (but toxic in excess)
Bromine	Br	35	A red-brown liquid halogen (one of only two elements liquid at room temperature) used in flame retardant chemicals and water purification compounds
Krypton	Kr	36	Inert noble gas used in certain fluorescent lamps and camera flash bulbs; an isotope of krypton was once used to define the length of a meter in science

Now, let's explore each of these elements in more detail, to understand their characteristics, interesting facts, common uses, and roles in science and everyday life.

Alkali and Alkaline Earth Metals (Group 1–2)

Potassium (K)

Main article: Periodic table/Alkali metals/Potassium

Potassium metal is so reactive that it must be stored under oil. In air it quickly tarnishes to a dull gray, and in water it reacts violently, often catching fire. Potassium is the first element in Period 4 and a classic alkali metal (group 1), sitting directly under sodium in the periodic table. In its pure form potassium is a silvery-white, extremely soft metal – soft enough to cut with a knife. It is also incredibly light (less dense than water), so a chunk of potassium would actually float – though it cannot safely float for long because potassium reacts explosively on contact with water. In fact, potassium is one of the most reactive elements known. Exposed to air, it tarnishes within seconds by reacting with oxygen, forming a pale peroxide coating. This intense reactivity means potassium is never found as a native metal in nature; instead it occurs only in salts and minerals.

Despite (or rather because of) its reactivity, potassium is an essential element for life. Potassium ions (K^+) are vital electrolytes in all living cells, helping to transmit nerve signals and regulate cellular processes. Both plants and animals require potassium in substantial amounts. For instance, in agriculture, potassium is a major nutrient given to crops in the form of potash fertilizers – in fact, around 90–95% of all manufactured potassium compounds (like potassium chloride) are used for fertilizers worldwide. Everyday sources of potassium for humans include foods like bananas, potatoes, and many others. In the human body, potassium (along with sodium) maintains fluid balance and is critical for heart and muscle function; too little or too much can cause serious health issues.

Historically, potassium was the first metal ever isolated by electrolysis (in 1807 by Sir Humphry Davy, who obtained it from potash, the ashes of plants). The element’s name comes from “potash” (potassium carbonate in wood ashes), and its symbol K derives from kalium, the Latinized version of an Arabic word for potash. In everyday experience, one seldom encounters metallic potassium for safety reasons – but its compounds are everywhere, from the salt substitute “potassium chloride” used in low-sodium diets to potassium nitrate in fireworks. When burned, potassium salts produce a lilac-colored flame, a distinguishing feature in flame tests. In short, potassium is a highly reactive metal that is nonetheless crucial for biological nutrition and widely used to support plant growth.

Calcium (Ca)

Main article: Periodic table/Alkaline earth metals/Calcium

Calcium is the second element in Period 4 and belongs to the alkaline earth metals (group 2), directly below magnesium on the table. Like potassium, pure calcium metal is quite reactive (though a bit less violently so) and is never found uncombined in nature. Any elemental calcium exposed will quickly react, especially with water or moisture, to form calcium hydroxide. In fact, if you drop pure calcium metal into water, it will fizz as it generates hydrogen gas (though calcium’s reaction with water is less vigorous than that of the alkali metals). Because of this reactivity, metallic calcium must be stored carefully (often in an inert atmosphere or oil). However, calcium compounds are extremely common on Earth – calcium is a major component of minerals like limestone (calcium carbonate), gypsum (calcium sulfate), and many others.

Calcium’s most famous role is in biology. It is one of the most abundant elements in the human body (by mass, the most abundant mineral element in our bodies) and absolutely essential to living organisms. Calcium compounds provide structural strength to bones and teeth – the rigid framework of our skeleton is largely calcium phosphate (in the form of hydroxyapatite). In shells of marine organisms and eggshells of birds, calcium carbonate provides hardness and protection. Beyond structural uses, calcium ions in solution serve as important cellular messengers, helping trigger muscle contractions, nerve impulses, and other cellular processes. We get calcium from foods like dairy products, leafy greens, or fortified foods, and a long-term calcium deficiency can lead to weak bones (osteoporosis).

In industry and everyday life, calcium compounds have a variety of uses. Lime (calcium oxide) has been used since ancient times to make mortar and plaster; mixed with water it forms slaked lime (calcium hydroxide), which is a key ingredient in cement and whitewash. Chalk and marble are forms of calcium carbonate; chalk was traditionally used for writing and drawing, and marble is a prized building and sculpting material. Calcium is also used as a reducing agent in metallurgy to extract other metals from their ores, and calcium carbonate is employed to neutralize acidic soils in agriculture. While you won’t encounter metallic calcium casually (it would readily corrode and react), the element’s compounds are all around us – in concrete, in antacid tablets (calcium carbonate neutralizes stomach acid), and of course, in our own bones. Calcium truly lives up to its name’s Latin root calx meaning lime (the mineral), underscoring how fundamental it is to both geology and biology.

Transition Metals (Group 3–12)

Scandium (Sc)

Main article: Periodic table/Transition metals/Scandium

Scandium is the first of the transition metals in Period 4 (and indeed the very first transition metal in the periodic table, as it inaugurates the 3d series). It has atomic number 21 and lies in group 3, just under yttrium. As a metal, scandium is silvery-white and relatively soft. It’s classified as a transition metal, but its chemistry often resembles that of the rare earth elements (lanthanides) due to a similar ionic size. Scandium is a rather scarce element in Earth’s crust; it typically occurs dispersed in trace amounts within rare minerals. This scarcity, combined with its chemical behavior, made it hard to isolate and find applications for scandium.

One interesting fact is that scandium was one of the elements predicted by Dmitri Mendeleev – he called the undiscovered element “eka-boron.” Scandium was subsequently discovered in 1879 in minerals from Scandinavia (hence its name, from Scandia, Latin for Scandinavia). Despite being moderately abundant in the Sun and certain stars, scandium is not commonly encountered on Earth and is considered a rare metal commercially. It is difficult to extract and purify, which historically limited its use.

In terms of uses, scandium has very few major commercial applications. The most notable use is in aluminum-scandium alloys. Adding a small fraction of scandium (just a percent or less) to aluminum can significantly increase the strength and reduce grain size of the alloy. Such Al-Sc alloys are used in high-performance sports equipment (like some baseball bats, lacrosse sticks, or bicycle frames) and in aerospace components (for example, Russian MIG fighter jets were reportedly built with scandium-alloyed aluminum for strength). Scandium oxide (scandia) is also used as an additive in high-intensity discharge lamps – scandium iodide mixed with mercury vapor produces a light that mimics sunlight, useful in stadium lighting. Other than these niche uses, scandium remains a curious element with limited presence in daily life. Its rarity and cost keep it something of a laboratory and specialty material. Nonetheless, scandium’s discovery was crucial historically as it helped validate the predictive power of the periodic table (filling the gap Mendeleev had pointed out), cementing the idea that the periodic arrangement could forecast new elements.

Titanium (Ti)

Main article: Periodic table/Transition metals/Titanium

Titanium (atomic number 22) is a well-known transition metal, famous for its combination of light weight and strength. It resides in group 4, under carbon (if counting by group number, but chemically under zirconium) and is often associated with strength-to-weight excellence. Pure titanium is a silvery-gray metal that is remarkably strong for its low density – it’s as strong as some steels but 45% lighter. Titanium is also highly resistant to corrosion; it forms a thin oxide layer that protects it from rusting or reacting, even in harsh conditions. Because of these properties (lightweight, strong, corrosion-proof), titanium has been dubbed a “wonder metal” for many advanced applications.

Titanium is the ninth most abundant element in Earth’s crust, but it’s mostly found in the form of oxides (like ilmenite and rutile). Extracting pure titanium metal is an energy-intensive process (often via the Kroll process), which historically made titanium quite expensive. Despite the cost, titanium’s properties have made it indispensable in certain fields. One major use is in the aerospace industry: titanium alloys are used extensively in aircraft frames, jet engine components, and spacecraft, where high strength and low weight are critical. For example, the skin of the SR-71 Blackbird spy plane and parts of modern jet engines contain significant titanium.

Another everyday use is in medical implants: titanium is biocompatible (it doesn’t react with body tissues and is not rejected by the immune system), so it’s used for artificial bone pins, joint replacements, dental implants, and even in pacemaker casings. If you have a titanium screw in a repaired bone or a titanium hip joint, it’s because the metal is strong, light, and doesn’t corrode inside the body. Titanium’s corrosion resistance also means it can be used in seawater environments (like in ship parts or ocean engineering) and for chemical plant equipment that handles corrosive substances.

Titanium dioxide (TiO₂) is a very common compound of titanium – it’s a brilliant white pigment used in paints, sunscreen, and food coloring. In fact, if you see white paint or toothpaste, chances are it contains TiO₂. The metal’s name comes from the Titans of Greek mythology, reflecting its robust nature. You might also encounter titanium in consumer products: for instance, high-end sports equipment (golf club heads, bicycle frames, tennis rackets) and even jewelry and watch cases use titanium for its strength, lightness, and attractive finish. It’s not an exaggeration to say titanium is a material where science fiction became reality – once as rare as unobtainium in everyday life, now it’s in everything from airplanes to eyeglass frames, enabling strong yet lightweight designs.

Vanadium (V)

Main article: Periodic table/Transition metals/Vanadium

Vanadium (atomic number 23) is a transition metal in group 5, sitting just under niobium. It’s a hard, silver-gray metal that often flies under the radar in general discussions, yet it has some important industrial roles. Pure vanadium metal is not commonly encountered (and is rarely used in pure form). In nature, vanadium is typically found in minerals and ores, often alongside other metals. It is never found as a free element, because it readily forms stable compounds; vanadium minerals can be found in certain iron ores and in minerals like vanadinite.

One of vanadium’s claims to fame is the colorful chemistry of its ions – vanadium can exist in several oxidation states (from +2 to +5), each of which can impart a different color to its solutions (purple, green, blue, yellow, etc.). However, a more practical significance of vanadium is its use in metallurgy. About 80% (or more) of produced vanadium is used as an alloying element in steels. Even a small addition of vanadium (on the order of a few tenths of a percent) can significantly increase the strength, toughness, and wear resistance of steel. Vanadium-alloyed steels, such as high-speed tool steels or certain high-strength low-alloy steels, are invaluable for making cutting tools, crankshafts, gears, and other critical components. The famously tough vanadium steel chassis of the Ford Model T automobile (early 1900s) was an early example of vanadium’s utility in improving steel. When you see tools labeled “chrome-vanadium steel,” the vanadium is there to enhance hardness and durability.

Vanadium’s most important compound is vanadium pentoxide (V₂O₅). This compound is used as a catalyst in the chemical industry, notably in the production of sulfuric acid via the contact process. Vanadium pentoxide also finds use in ceramics and as a pigment. Another fascinating application of vanadium emerging in recent years is in vanadium redox flow batteries – these are large rechargeable batteries for energy storage on power grids. Vanadium’s multiple ionic states allow energy to be stored and released in solutions, making it useful for large-scale storage of renewable energy (like solar or wind power).

It’s worth noting that vanadium compounds tend to be toxic. All vanadium compounds are considered at least somewhat toxic, and some (in high oxidation states) are highly poisonous. Fortunately, vanadium is needed only in trace amounts biologically (if at all – it’s not firmly established as essential for humans, though some enzymes in other organisms use vanadium). The body handles vanadium poorly and too much can cause health issues. On the flip side, a curious fact: certain sea creatures, like some ascidians (sea squirts), concentrate vanadium from seawater to astonishing levels, puzzling scientists as to why.

Vanadium was discovered in 1801 by Andrés Manuel del Río, but it was later rediscovered and confirmed in 1831 by Nils Gabriel Sefström, who named it after “Vanadis,” a name of the Norse goddess Freyja, because of the element’s multicolored compounds. So while you might not hear about vanadium as often as iron or copper, whenever you drive a car or use a power tool, vanadium-strengthened steel likely plays a part. It’s a silent strengthener – making our metals tougher behind the scenes.

Chromium (Cr)

Main article: Periodic table/Transition metals/Chromium

Chromium (atomic number 24) is a shiny, steely-gray transition metal in group 6 (beneath molybdenum). The name chromium comes from the Greek chroma, meaning "color," because chromium compounds are famously colorful. For example, potassium dichromate is bright orange, chromium(III) oxide is green, and many gemstones (like rubies and emeralds) owe their colors to traces of chromium. But beyond pretty colors, chromium metal has significant practical value, especially known for its ability to resist corrosion.

Chromium is probably best recognized for its use in stainless steel. Stainless steel is an alloy that contains at least around 10–13% chromium (often along with iron, carbon, nickel, etc.). The chromium in the steel forms a thin, protective oxide layer on the surface, preventing rust (iron oxide) from forming, thereby “stain-less” steel remains shiny and rust-free. This application is ubiquitous – from kitchen sinks and cutlery to skyscraper cladding and surgical instruments, stainless steel’s durability is thanks to chromium. Chromium metal is also used for chrome plating: a thin layer of chromium can be electroplated onto other metals (like steel or plastic) to give a shiny, mirror-like finish that is both decorative and protective. The gleaming chrome trim on classic cars, motorcycle parts, or bathroom fixtures is essentially metallic chromium.

Another major use of chromium is in pigments. As mentioned, chromium’s compounds can be intensely colored. Lead chromate, for example, was the classic “chrome yellow” pigment used in school buses and paint (though its toxicity has reduced its use). Green chromium oxide is a stable pigment for paints and was used in military camouflage. Chromium compounds have also been used in tanning leather (chromium(III) helps crosslink collagen fibers).

On the flip side, some chromium compounds, particularly hexavalent chromium (Cr(VI)) compounds like chromates, are toxic and carcinogenic. This was highlighted in the famous Erin Brockovich case, where groundwater contaminated with hexavalent chromium posed a health hazard. In contrast, the trivalent form (Cr(III)) is less toxic and in trace amounts is even considered an essential nutrient for sugar metabolism in the human diet (though the necessity is debated).

In its pure form, chromium is quite hard and has a high melting point. It’s one of the hardest of the pure metals. When you see the dazzling shine of a car bumper or the reflective surface of a factory machine part, that’s chromium doing its job – protecting and beautifying. Discovered at the end of the 18th century in a mineral that produced colorful compounds, chromium today is a workhorse of metallurgy and materials science, valued both for its aesthetic glitter and its very real ability to protect and strengthen materials around us.

Manganese (Mn)

Main article: Periodic table/Transition metals/Manganese

Manganese (atomic number 25) is a transition metal in group 7, sitting just under technetium (and above rhenium). It is a gray-white metal, hard and brittle in pure form. Manganese doesn’t get as much public attention as some of its Period 4 neighbors, but it is absolutely essential in metallurgy and also important in biological systems.

One of manganese’s primary uses is in steel production. In fact, manganese is often found in combination with iron in nature (many iron ores contain manganese). When steel is made, manganese is usually added (in the form of ferromanganese or silico-manganese alloy) to serve several purposes: it helps remove oxygen and sulfur impurities, and it improves the strength and toughness of the steel. Hadfield steel, also known as manganese steel, contains around 12% manganese – it is extraordinarily tough and resistant to wear (used in railway tracks, rock crushers, and safes). Even ordinary steel typically has a small percentage of manganese (0.5–1%) to improve its hardness and to counteract brittleness. Manganese also helps steel resist rusting to some extent (though not as effectively as chromium does in stainless steel). In stainless steels, a combination of manganese and nitrogen is sometimes used to replace nickel for a more cost-effective alloy. Without manganese, we wouldn’t have the high-quality steels we rely on in construction and manufacturing.

Another everyday application of manganese is in dry cell batteries. The common alkaline battery (and the older zinc-carbon battery) uses manganese dioxide (MnO₂) as the cathode material. When you use a AA battery, the electric current is being facilitated by manganese dioxide reacting inside. Manganese dioxide is excellent for this because it’s a strong oxidizer and can accept electrons from zinc during the discharge of the battery. So every time you turn on a flashlight or TV remote using alkaline batteries, manganese is quietly at work.

Biologically, manganese is a trace element that organisms need. It’s found in some enzymes and is important in bone formation and metabolic processes. However, manganese can be toxic in high amounts. There is a condition known as manganism – a neurological disorder resembling Parkinson’s disease – that can afflict workers who inhale a lot of manganese dust or fumes (for example, in mining or welding without proper protection). Inhaling excess manganese over time can cause irreversible nervous system damage. Manganese compounds in certain oxidation states can also be poisonous if ingested in quantity. That said, normal dietary manganese (from nuts, grains, leafy vegetables, etc.) is both safe and essential.

An interesting historical use of manganese dioxide was in glassmaking – Roman glassmakers used it to remove the greenish tint from glass (caused by iron impurities), referring to MnO₂ as “glassmaker’s soap.” Also, manganese dioxide is what gives the deep purple color to amethyst gemstones. The element’s name comes from the Latin magnes (“magnet”), because pyrolusite ore (manganese dioxide) was once confused with a magnetic iron ore (loadstone). Pure manganese isn’t magnetic, but some of its alloys are.

In summary, manganese might not be a household word, but it’s a backbone element in industry – crucial for making steel tough and batteries functional – and it lurks in the background of our daily technologies and even in our bodies’ biochemistry.

Iron (Fe)

Main article: Periodic table/Transition metals/Iron

Iron is perhaps the most famous of the Period 4 elements – atomic number 26, group 8 – and one of the most important metals in human civilization. It’s not an exaggeration to say we live in the “Iron Age” even today, as iron (mostly in the form of steel) is the material pillar of modern infrastructure. Pure iron is a lustrous, silvery-gray metal that is reasonably soft and malleable, but it usually doesn’t stay pure for long outside of controlled conditions; iron reacts with oxygen and water to form rust (iron oxide), which is why unprotected iron will corrode.

Iron is incredibly abundant on Earth – it’s the fourth most abundant element in the Earth’s crust and by far the main component of the Earth’s core. In fact, among the Period 4 elements, iron is the most common element on Earth overall (by mass). It’s believed that the Earth’s solid inner core and liquid outer core are composed chiefly of iron (with some nickel). This abundance made iron accessible (via its ores) to ancient peoples, leading to the Iron Age when techniques to smelt iron from ore were developed.

The most significant use of iron is in steel. Steel is an alloy of iron with a small percentage of carbon (and often other elements). The addition of carbon (even less than 2%) dramatically increases iron’s strength, making steel the backbone of buildings, bridges, railways, automobiles, tools, and countless other things. Cast iron (with more carbon) has been used historically for pots and pipes, while refined steels are used for everything from razor blades to skyscrapers. Iron and steel are so ubiquitous that our daily life is full of them – the frame of your car, the rebar in concrete buildings, the machinery in factories, and so on. Iron’s importance in industry is impossible to overstate.

Beyond its structural role, iron is central to biology as well. Iron is a key component of hemoglobin, the protein in red blood cells that carries oxygen from our lungs to our tissues. The iron in hemoglobin binds oxygen molecules; each hemoglobin can carry four oxygen molecules attached to iron atoms. This is what gives blood its red color (iron-oxygen complexes are red). Without enough iron in our diet, we become anemic and weak because our blood can’t transport oxygen efficiently. Iron is also present in many enzymes and is vital for plants (in enzymes for photosynthesis).

An interesting fact from nuclear physics: iron-56 (one of iron’s isotopes) has the highest binding energy per nucleon of any element’s isotopes, which means it is the most stable nucleus. This has cosmic implications – when massive stars fuse elements in their core, they can fuse up to iron, but no further without inputting energy. So iron is the last element produced in the core of a star before a supernova; heavier elements are only formed in the violent supernova explosion or other extreme events. In essence, iron is the “ash” of stellar fusion – the end product of energy-releasing fusion processes. This is why iron is so common in the universe and in meteorites (and thus early humans first encountered iron in meteorites as "gift from the heavens").

Culturally, iron has been known and used for thousands of years (the Iron Pillar of Delhi is a testament to ancient iron-working skill, rust-resistant after 1600 years). Its symbol Fe comes from the Latin ferrum. Whether it’s the steel framework of the building you’re in, the cast-iron skillet in your kitchen, or the hemoglobin in your blood, iron is literally and figuratively a foundational element of both civilization and life.

Cobalt (Co)

Main article: Periodic table/Transition metals/Cobalt

Cobalt (atomic number 27) is a hard, lustrous, silver-gray metal in group 9, sitting between iron and nickel on the periodic table. While not as widely talked about as iron or copper, cobalt has some fascinating attributes and uses. One of cobalt’s historical and well-known aspects is its contribution to brilliant blue colors. Cobalt compounds (like cobalt aluminate, also known as "cobalt blue") have been used for centuries to impart a deep blue color to glass, ceramics, and paints. The word “cobalt” itself comes from the German Kobold, meaning goblin – early miners found cobalt ores troublesome because when smelted they produced toxic arsenic fumes and no useful copper (they were looking for copper), so they blamed goblins for the ore’s “deceptive” nature. But eventually, those ores yielded the element cobalt, which became invaluable for coloring glass a rich blue (think of the deep blue of some wine bottles or cathedral stained glass – that’s often cobalt).

In modern times, cobalt’s value extends far beyond pigments. It is a component of high-performance alloys. For instance, certain superalloys used in jet engine parts, gas turbines, and rocket engines contain cobalt for its ability to retain strength at high temperatures. Cobalt is also ferromagnetic (like iron and nickel), and it’s used in making strong permanent magnets (e.g., alnico magnets contain aluminum, nickel, cobalt). Hard facing alloys (like Stellite) that resist wear and corrosion often contain cobalt. In steel tools, adding a bit of cobalt improves performance at high cutting temperatures (some high-speed steels have 5–12% cobalt).

Cobalt finds use in rechargeable battery technology as well. Lithium-ion batteries, common in smartphones, laptops, and electric cars, often use lithium cobalt oxide in their cathodes (though efforts are underway to reduce cobalt due to cost and ethical sourcing concerns). This connection to battery tech has made cobalt a strategically important element in recent years.

Biologically, cobalt is unique among Period 4 elements because it is an essential part of vitamin B₁₂ (cyanocobalamin). Vitamin B₁₂ is a complex organometallic molecule that contains a cobalt ion at its core. Our bodies (and those of other animals) need B₁₂ for proper nerve function and blood cell formation, which means we all need a little bit of cobalt in our diet – typically obtained from meat, eggs, or B₁₂ supplements, since plants do not make B₁₂. This is why cobalt is considered an essential trace dietary element, albeit indirectly (through B₁₂). Cobalt deficiency per se isn’t a thing in humans (it would manifest as B₁₂ deficiency), but in grazing animals like sheep and cattle, cobalt-deficient soils can lead to health issues, so sometimes cobalt is added to fertilizers.

On the other hand, a radioactive isotope of cobalt, cobalt-60, is notable (and notorious). Cobalt-60 is produced in nuclear reactors and emits intense gamma rays. It has been used in radiation therapy for cancer (cobalt “teletherapy” machines) and for sterilizing medical equipment. However, cobalt-60 is also a component of nuclear fallout and can be hazardous if dispersed in the environment. In theory (and fiction), a “cobalt bomb” could maximize lingering radioactive contamination. Fortunately, Co-60 has a half-life of about 5.3 years, so it doesn’t last incredibly long in the environment, but while active it’s quite dangerous.

In summary, cobalt is a multifaceted element: historically giving us beautiful blue arts, industrially strengthening our metals and driving battery chemistry, and biologically sitting at the heart of a vital vitamin. From the blue of medieval stained glass to the lithium-ion battery in your phone, cobalt has a colorful and critical presence in our world.

Nickel (Ni)

Main article: Periodic table/Transition metals/Nickel

Nickel (atomic number 28) is a transition metal in group 10, known for its strength, shininess, and resistance to corrosion. In appearance, nickel is a silvery metal with a slight golden tinge. It’s hard, ductile, and ferromagnetic (nickel is one of the few elements that is magnetic at room temperature, the others being iron and cobalt). Nickel’s major claim to fame is as a component in alloys, especially stainless steel and other corrosion-resistant alloys. Just like chromium, nickel is commonly added to steel; in the case of austenitic stainless steel (e.g., 304 or 316 stainless), nickel (around 8–10%) is present along with chromium. The nickel stabilizes the steel’s crystal structure and adds toughness and corrosion resistance. Many stainless steels owe their strength and non-rusting nature partly to nickel.

Another everyday association with nickel is in coins – indeed, the US 5-cent coin is called a “nickel” because it historically contained nickel metal. Many countries’ coins are made of nickel alloys or nickel-plated steel, since nickel is hard and takes a good shine. For example, cupronickel (copper-nickel alloy) is used for coins (like US nickels are actually 25% nickel, 75% copper). Nickel’s usage in coins has declined somewhat (due to cost), but the legacy is there – even the name of the coin. Earlier in history, the metal was not widely used until the 18th century when it was identified properly; before that, some nickel ores were misidentified (the German miners who named kupfernickel or “Devil’s copper” were dealing with nickel ore that looked like copper ore but yielded none).

Nickel is also crucial in electroplating. A layer of nickel can be plated onto other metals to provide a protective, shiny coating (often as a base layer before chrome plating). Many bathroom fixtures, for instance, are nickel-plated beneath the chrome. Nickel plating is used on coins, jewelry, and industrial parts to prevent corrosion and wear.

In the field of batteries, nickel has had important roles: Ni-Cd (nickel-cadmium) and Ni-MH (nickel-metal hydride) batteries, which have been used in portable electronics and hybrid cars, rely on nickel in their electrodes. The iconic rechargeable AA batteries were often Ni-Cd or Ni-MH before lithium-ion took over. Nickel is also a component of some catalysts (e.g., Raney nickel, used in hydrogenation of organic compounds like in margarine production).

From a geological perspective, nickel is noteworthy because, along with iron, it is one of the two main elements in the Earth's core. Most of the Earth’s nickel sank to the core during planet formation (hence relatively less in the crust). Nickel is relatively rare in the Earth’s crust (much rarer than iron), which is one reason pure nickel metal was historically hard to get (most nickel comes from sulfide ores like pentlandite or laterite ores). Interestingly, much of the nickel humans first used came from meteorites – iron meteorites are typically iron-nickel alloy (around 5–20% nickel). The presence of nickel in a metal is a good test for meteorite vs. manmade iron, because pre-industrial iron rarely had nickel, whereas meteorites always do.

On the human health side, nickel is not known to be essential for us (though some microorganisms and plants use it in enzymes). However, nickel metal and salts can cause allergy in some people – nickel allergy is fairly common, causing contact dermatitis (rash) if nickel-containing jewelry touches the skin. That’s why many earrings and watch backs are now “nickel-free.”

Overall, nickel’s role in alloys is where it shines the most (sometimes literally). The term “German silver” (which contains nickel, copper, zinc) and things like Monel metal (a nickel-copper alloy) highlight its utility. Nickel provides materials with toughness, lustre, and longevity. Whether in the coins jingling in pockets, the steel frame of a skyscraper, or the core of our planet, nickel is a quietly ubiquitous element that lends resilience and polish to the world.

Copper (Cu)

Main article: Periodic table/Transition metals/Copper

Copper (atomic number 29) is one of the best-known metals in the world, prized for its electrical conductivity and malleability. Copper has been used by humans for thousands of years – it was one of the first metals ever extracted and utilized (giving its name to the Copper Age). In pure form, copper is a distinctive reddish-orange metal, which sets it apart from the usual silvery-gray metals. It’s relatively soft, very ductile (easy to draw into wires), and an excellent conductor of both heat and electricity. In fact, among pure metals at room temperature, only silver has higher electrical conductivity than copper, but silver is too expensive for general use – so copper is the metal of choice for wiring.

The most impactful use of copper is in electrical wiring and electronics. Our modern electrically powered world runs on copper: the wires in the walls of houses, the windings in electric motors, the circuits in appliances – they all rely on copper. Roughly 70% of all copper in use goes into electrical applications (wires, cables, motors, transformers). Copper’s combination of high conductivity, reasonable cost, and durability makes it ideal for carrying electric current. If you look at a power cord or the guts of a computer, the reddish metal you see is copper.

Copper is also a major component of many alloys. Two famous copper alloys are bronze (copper + tin) and brass (copper + zinc). Bronze was so important that an entire era of human history is named after it. It’s harder than pure copper and was used for tools, weapons, and art for millennia. Brass has a golden color and is used in musical instruments, decorative hardware, and coins. Copper is also mixed with nickel to make cupronickel (used in some coins like US nickels and the outer layer of quarters and dimes), and with nickel and zinc to make “German silver.”

Another big use of copper is in plumbing and building materials. Copper pipes have been common for household water supply lines for decades (though plastic pipes are increasing). Copper roofing, gutters, and architectural details are valued because copper resists corrosion – it slowly forms a green patina (copper carbonate, like on the Statue of Liberty) which actually protects it from further corrosion. That green patina is often seen on old domed roofs and statues.

Beyond infrastructure, copper plays a role in biological systems too. It is an essential trace element for most living organisms. In humans, copper is part of enzymes and is needed for iron metabolism and other functions. Certain animals even use copper in their blood in place of iron – for example, horseshoe crabs and octopuses have blue blood because they use a copper-based oxygen carrier (hemocyanin) instead of hemoglobin. But in higher quantities, copper can be toxic (which is why copper-based fungicides and anti-microbial surfaces work – they kill microbes). Historically, people noticed that water stored in copper vessels tended to be safer to drink; indeed, copper surfaces have natural antibacterial properties (leading modern hospitals to experiment with copper alloy touch surfaces to reduce infections).

Culturally and historically, copper has an illustrious place. The symbol Cu comes from cuprum, Latin for Cyprus, the island famed for copper mining in Roman times. Ancient Egyptians and Sumerians were smelting copper by around 5000 BC. The Bronze Age (starting around 3300 BC) was enabled by copper-tin alloys. Fast forward to today, copper is literally driving the Information Age, hidden in the walls and circuits that connect our world. It’s also present in everyday items: coins (like pennies, which are mostly zinc with copper plating nowadays), cookware (copper bottom pans for even heating), and the electronics we carry. Copper’s reddish shine and versatility ensure that it remains one of the most cherished and useful metals. And unlike many elements, copper can be found in nature in metallic form – nuggets of pure copper were found in places like the Keweenaw Peninsula (Michigan), making it easily accessible to early humans. So as an element, copper bridges our ancient past and high-tech present, conducting electricity and history with equal ease.

Zinc (Zn)

Main article: Periodic table/Transition metals/Zinc

Zinc (atomic number 30) is the last of the ten transition metals in Period 4. It’s a bluish-silver metal that is not often encountered in pure shiny form by the general public, yet zinc’s compounds and alloys are extremely common. Chemically, zinc behaves somewhat like a “late” transition metal or even a post-transition metal, often preferring to stay in the +2 oxidation state (Zn²⁺) in compounds. One of zinc’s hallmark properties is its ability to protect other metals from corrosion – this is the principle behind galvanization, where iron or steel is coated with a layer of zinc to prevent rusting.

If you’ve seen a piece of galvanized steel (like a roadside guardrail or a roofing nail), you might notice a spangled crystalline pattern on the surface – that’s zinc coating. When exposed to air, zinc forms a thin protective layer of zinc carbonate, which unlike rust, doesn’t flake off easily, thus shielding the iron underneath from water and oxygen. Galvanized steel is used for buckets, outdoor structures, car bodies, and countless other items. Galvanization is probably the most widespread use of zinc by volume; it has been said that about half of all zinc produced goes into galvanizing iron and steel.

Zinc is also famously used in alloys, most notably brass (copper–zinc alloy). Brass has been in use since antiquity (the Romans were making brass coins and ornaments). It has a pleasant gold-like color and good workability. Depending on the ratio of copper to zinc, brasses can be soft for easy shaping or harder for durability. Musical instruments (trumpets, trombones) and many hardware pieces are made of brass. Another alloy, zamak, combines zinc with aluminum and other elements for die-casting toys, carburetors, etc. Historically, an earlier zinc alloy called “pot metal” was used for inexpensive cast items.

As a pure metal, zinc is not structurally strong, but it has a low melting point (~420 °C), which makes it easy to cast or form into shapes and to alloy with other metals. It’s also used in making batteries: the common zinc-carbon battery (the old-style flashlight batteries) uses a zinc can as the anode and a carbon rod as the cathode with MnO₂ inside. Modern alkaline batteries similarly use zinc powder as the anode (with manganese dioxide cathode). Zinc is also part of rechargeable systems like zinc-air batteries and some newer flow battery concepts.

In the realm of biology and health, zinc is an essential trace element. Almost 2 billion people worldwide suffer from zinc deficiency, which can lead to weakened immune function and other health problems. Our bodies require zinc for the proper function of over 300 enzymes and many proteins. It plays roles in immune response, DNA synthesis, wound healing, and more. This is why zinc supplements or zinc lozenges are often marketed to support immunity (such as for colds). We get zinc from foods like meat, shellfish, legumes, and seeds. Conversely, too much zinc can interfere with other minerals (excess zinc can cause copper deficiency, for example), so balance is key.

Zinc compounds also have some interesting uses: zinc oxide is a white powder used in paints, rubber, cosmetics (traditional calamine lotion and some sunscreens have ZnO – it’s that white paste lifeguards put on their nose). It’s also used in making varistor electronics and as an additive in food (breakfast cereals are often “fortified with zinc”). Zinc sulfide is used in some luminous paints (glow-in-the-dark), and historically in cathode ray tube screens.

The name zinc comes from German Zink, of uncertain origin, possibly referring to the sharp tooth-like appearance of zinc metal crystals. Metallurgically, zinc was recognized as a distinct metal relatively late (Indian metallurgists were producing zinc by distillation by the 12th century, but in Europe pure zinc wasn’t common until the 18th century). One quirky fact: pennies in the United States since 1982 are actually 97.5% zinc with only a thin copper coating – a reversal from older pennies which were solid copper. So, that “copper” coin in your pocket is mostly zinc inside.

In summary, zinc may not have the glamour of gold or the notoriety of lead, but it’s a workhorse element. It extends the life of steel through galvanization, forms useful alloys like brass, powers batteries, and even helps keep us healthy. In the grand lineup of Period 4, zinc fittingly ends the transition metals with a solid contribution to both industry and biology.

Post-Transition and P-Block Elements (Group 13–18)

Gallium (Ga)

Main article: Periodic table/Boron group/Gallium

Gallium (atomic number 31) marks the start of the “p-block” elements in Period 4, and it’s quite a fascinating metal. Gallium is perhaps most famous for a single quirky property: its low melting point. Solid gallium metal will melt in your hand, since its melting point is about 29.8 °C (85.6 °F). That’s right – on a warm day, gallium turns to liquid. Hold a small piece of gallium, and soon you’ll have silvery liquid metal dripping through your fingers (safe to touch, as gallium is not toxic like mercury, but it will stain). This surprising property makes gallium a great conversation piece. It also finds practical use in specialized thermometers or temperature references – for instance, gallium can be used in high-temperature thermometers where mercury would boil (gallium’s boiling point is extremely high, ~2400 °C). An alloy of gallium called Galinstan (gallium, indium, tin) is liquid at room temperature and is used as a mercury replacement in some thermometers.

Pure gallium is a soft, silvery metal (solid gallium is so soft it can be cut with a knife). It has an unusual trait of expanding upon freezing, which means gallium should not be stored in glass or metal containers that it could break when it solidifies (similar to water’s expansion on freezing).

Beyond its neat melting trick, gallium is extremely important in the realm of electronics and semiconductors. Gallium doesn’t occur as a free element in nature, but it’s obtained as a byproduct of aluminum and zinc refining. One of its key uses is in compounds like gallium arsenide (GaAs) and gallium nitride (GaN), which are III-V semiconductors. Gallium arsenide is used to make high-speed and optoelectronic devices – for example, many laser diodes (like those in CD/DVD/Blu-ray drives) and LED lights are made of GaAs or related materials. GaAs can convert electricity into laser light and is also used in some solar panels (notably in satellites or Mars rovers where high efficiency is needed). Gallium nitride is another superstar – it’s used for blue and green LEDs and for modern high-power, high-frequency transistors (like those in 5G base stations). The invention of GaN blue LEDs was so important it won a Nobel Prize in Physics in 2014, as it enabled white LED lighting. So whenever you use an LED flashlight, television, or energy-efficient lamp, gallium likely plays a role in the semiconductor that produces light.

Gallium’s story also has a Mendeleev connection: it was discovered in 1875 by Paul-Émile Lecoq de Boisbaudran, and its existence had been predicted by Mendeleev a few years earlier as “eka-aluminium.” The discovered properties matched Mendeleev’s predictions quite closely, helping validate the periodic table. The element was named “Gallia” for France (Latin Gallia), and amusingly also as a pun on the discoverer (Lecoq = “the rooster,” and the Latin for rooster is gallus).

In everyday life, gallium isn’t something you handle (unless you buy a bit for curiosity’s sake). But its presence is felt in the tech we use. For instance, many mobile phone chips use gallium compounds. Gallium is also used in some medical imaging agents and high-temperature thermometers. In research, a gallium alloy (Galinstan) is even being explored for use in flexible electronics and cooling.

So, gallium stands out as an element of surprises: a metal that melts in your hand yet helps emit the light from a smartphone screen or a Blu-ray laser. It’s a bridge between the tangible novelty of melting metal and the intangible realm of electrons and photons in modern technology.

Germanium (Ge)

Main article: Periodic table/Carbon group/Germanium

Germanium (atomic number 32) is a metalloid – it sits in group 14, right beneath silicon on the periodic table. In many ways, germanium is chemically and physically similar to silicon, which is not surprising given their alignment in the table. Germanium was a noteworthy element historically because Dmitri Mendeleev predicted its existence and properties (calling it “eka-silicon”) before it was discovered. Indeed, in 1886 Clemens Winkler discovered germanium in a new mineral and found that its properties (atomic weight, density, etc.) closely matched Mendeleev’s predictions, which was a triumph for the periodic law.

In its pure form, germanium is a shiny, hard, brittle, gray-white metalloid. It’s perhaps most famous for its role in early semiconductor technology. Along with silicon, germanium was one of the first semiconductors used in transistors. In fact, the very first transistor (invented in 1947 at Bell Labs) was made of germanium. Throughout the 1950s, germanium was commonly used in transistors and diodes in electronics. However, silicon eventually overtook germanium as the preferred semiconductor (silicon devices work better at higher temperatures, and silicon dioxide is an excellent insulator, which germanium’s oxide is not). Still, germanium has not vanished from the tech world. It’s used in some high-speed devices and often alloyed with silicon in modern chips (Silicon-germanium alloys are used in high-frequency integrated circuits and radio devices).

One significant application of germanium is in fiber optic cables and infrared optics. Germanium oxide added to the glass of fiber-optic cables helps increase the refractive index in the fiber core, allowing efficient light transmission for telecommunications. Germanium is also transparent to infrared light, so germanium lenses are used in thermal imaging cameras (like night-vision systems, IR cameras for the military or for firefighting) – those black glassy lenses on infrared cameras are often germanium.

In chemistry, germanium forms organogermanium compounds somewhat akin to organosilicon. These have been investigated for various applications but are not nearly as widespread as, say, organosilicon polymers (silicones).

Germanium’s presence in everyday life is subtle. You might not directly handle germanium, but if you use the internet (fiber optics) or have an infrared thermometer, germanium is likely involved. Some low-energy light bulbs and LEDs also involve germanium in their semiconductor materials. In metallurgy, adding germanium to alloys can improve their properties (e.g., some aluminum alloys have a bit of Ge).

Germanium is considered a semiprecious material. It’s not super common – extracted largely from the byproducts of zinc ore processing or from certain coal deposits – and thus it is relatively expensive. Each year only a few dozen tons of germanium are produced. Unlike silicon (which is abundant sand), germanium is scarce.

Biologically, germanium has no known essential role (though there was a fad of “organic germanium” supplements decades ago, those are not proven and could be harmful if overused). Germanium compounds can irritate eyes, skin, lungs, but in general germanium is not highly toxic (comparable to silicon in low toxicity, but certain organogermanium compounds did cause kidney damage in people who overdosed on those supplements).

In short, germanium might be thought of as a shadow silicon – doing many of the same things on a smaller scale. It played a starring role in the dawn of semiconductor electronics and continues to support modern optics and electronics in important but less visible ways. And not to forget, its discovery was a pivotal moment for the periodic table’s acceptance, as it beautifully confirmed Mendeleev’s predictive power.

Arsenic (As)

Main article: Periodic table/Pnictogens/Arsenic

Arsenic (atomic number 33) is an element with a split personality of sorts – infamous as a poison, yet also a dopant in cutting-edge electronics, and even a possible medicinal in past eras. It sits in group 15 (pnictogens) under phosphorus, and it’s typically classed as a metalloid. Arsenic has several allotropes, but the most common form is a steel-gray, brittle metallic-looking solid.

The notoriety of arsenic comes largely from its toxicity. Compounds like arsenic trioxide (white arsenic) have been used as poisons for centuries – it was historically dubbed the "king of poisons" and the poison of kings. It’s odorless, tasteless (in food or drink), and deadly, which made it a favored substance for nefarious deeds in history. The symptoms of arsenic poisoning (before modern forensics) could be mistaken for cholera or food poisoning, allowing murders to go undetected. One macabre bit of trivia: arsenic was used by some as a cosmetic or tonic in small doses (so-called Fowler’s solution in Victorian times) – obviously a dangerous practice, but it persisted due to ignorance of long-term effects.

Arsenic is poisonous to most multicellular life because it disrupts essential metabolic pathways. It can mimic phosphorus (being in the same group), which is a key element in DNA and energy molecules, thereby interfering with biochemistry. In fact, arsenate (AsO₄³⁻) can substitute for phosphate in ATP, but then those compounds are unstable – essentially sabotaging cellular energy. This is why arsenic in groundwater (as happens in parts of Bangladesh, for example) is a serious health hazard causing organ damage, skin lesions, and cancers.

However, arsenic is not only a poison. It has uses in modern technology. A significant one is in semiconductors: when added in tiny amounts to silicon, arsenic acts as an n-type dopant, providing extra electrons for conduction. Many silicon-based microchips are doped with arsenic to create the desired electrical properties. Additionally, as mentioned with gallium, gallium arsenide (GaAs) is a crucial semiconductor for lasers and high-frequency electronics. Arsenic is also used in some specialized alloys; for instance, small percentages of arsenic can be added to lead (like in lead shot for shotgun shells) to harden it.

In agriculture, arsenic compounds were historically used as pesticides and herbicides. Lead arsenate was used in orchards to kill insects in the early 20th century, and some arsenic-based herbicides were used for weed control. Most of these uses have been phased out due to toxicity and environmental persistence. There is also chromated copper arsenate (CCA), which was widely used to pressure-treat lumber to prevent rot (the arsenic kills insects and fungi). This gave greenish-tinted wood that was common in decks and playgrounds, but concerns about arsenic leaching led to its reduction in residential uses.

On the plus side, certain organisms have developed tolerance or even use for arsenic. Some bacteria can metabolize arsenic compounds, and there was a contested 2010 report of a bacterium that might incorporate arsenic into its DNA in place of phosphorus (though that claim has been largely debunked or at least not widely accepted). Also, interestingly, arsenic in very tiny doses has been investigated for medicinal purposes (some leukemia treatments use arsenic trioxide to induce remission).

Another form of arsenic is arsenic sulfide, historically known as orpiment (yellow) and realgar (red), which were used as pigments and in traditional Chinese medicine. Unfortunately, those too can be toxic.

In terms of elemental forms, there’s gray arsenic (metallic form), yellow arsenic (molecular nonmetallic form like P₄ structure), etc. Gray arsenic is more stable. If heated, arsenic doesn’t melt into a liquid at normal pressure; instead, it sublimates (goes directly from solid to vapor).

In the periodic context, arsenic’s presence in Period 4 adds a nonmetallic flavor among the metals. It forms the bridge between metals and nonmetals in this period. Its chemistry is predominantly covalent. Arsenic’s neighbors are interesting: Germanium (semiconductor), Selenium (nonmetal), showing the gradual transition.

In summary, arsenic is an element of infamy and utility: deadly poisons and beneficial semiconductors. It reminds us that what makes an element “good” or “bad” often depends on dosage and context. While we avoid arsenic in our drinking water, we welcome it (in minute quantities) in our smartphones and electronics. Arsenic’s dual nature has secured it a permanent place in both the dark annals of crime and the bright forefront of technology.

Selenium (Se)

Main article: Periodic table/Chalcogens/Selenium

Selenium (atomic number 34) is a nonmetal in group 16 (the chalcogens), sitting just below sulfur. It doesn’t get as much popular press as its lighter cousin sulfur, but selenium is quite interesting in its own right. In pure form, selenium can exist in several allotropes. There’s red selenium (amorphous or monoclinic Se₈ rings, similar to sulfur’s S₈ rings) which is powdery, and there’s gray (or metallic) selenium, which is more stable – a gray, brittle, lustrous form that conducts electricity better in light than in dark. This property – higher electrical conductivity under illumination – is called photoconductivity and it made selenium very important in early electronics.

One of the most historically significant uses of selenium was in photocopiers (Xerography). The very first xerographic copiers in the mid-20th century used a selenium-coated drum. How did that work? A selenium layer would be given an electrostatic charge in the dark. Then a light image of the page to be copied was projected onto the drum. In the bright areas (corresponding to white paper), selenium became conductive and the charge leaked away; in dark areas (text), selenium stayed insulating and held its charge. Then toner particles (charged oppositely) would stick to the areas that still had charge (forming the text image), and then be transferred to paper and fused. Selenium’s photoconductive property was key – in darkness an insulator, in light a conductor. Later photocopiers moved to organic photoconductors, partly because handling selenium fumes can be toxic, but selenium’s role in launching that technology was pivotal.

Another major use of selenium is in glass manufacturing. Small amounts of selenium can counteract the green tint caused by iron impurities in glass, effectively decolorizing glass to make it more clear. Or, adding a bit more can give glass a red hue (selenium ruby glass). So selenium is used both to make clear glass and to produce certain red/pink glasses. It’s also used in some pigments (like cadmium sulfoselenide yields a brilliant red pigment used in ceramics, known as cadmium red).

Selenium is also used in electronics such as rectifiers (early electronics used selenium rectifier stacks to convert AC to DC before silicon diodes took over). Selenium is still used in some solar cells, often alloyed with other elements (for example, copper indium gallium selenide – CIGS – is a thin-film photovoltaic material).

Biologically, selenium is an essential trace element for humans and many other organisms. It is a component of certain enzymes, most notably glutathione peroxidase, which helps protect cells from oxidative damage. Our bodies need only a very small amount of selenium – on the order of tens of micrograms per day. Foods like Brazil nuts (which are famously high in selenium), seafood, and grains can provide selenium. However, the gap between necessary and toxic is not huge; too much selenium leads to a condition called selenosis (with symptoms like hair loss, nail brittleness, and neurological damage). Livestock can also suffer from selenium deficiency or toxicity depending on soil selenium levels.

Interestingly, some plants will accumulate selenium (notably some species in the American West), which can poison grazing animals. But other organisms need selenium: for example, selenium is necessary for proper thyroid function and in reproduction.

Selenium was discovered by Jöns Jacob Berzelius in 1817. He named it after the Greek moon goddess Selene, as it was chemically similar to tellurium (named for the earth). In industry, selenium often comes as a byproduct of copper refining.

In summary, selenium is a sort of under-the-radar element with a wide range of roles: it helped usher in the era of photocopiers and laser printers, it quietly makes our glass clearer and our pigments brighter, and within our bodies it protects our cells in minute doses. But like many things, the dose makes the poison – too much selenium is harmful. Selenium’s placement in the periodic table – the first nonmetal after a stretch of metals – marks a clear change in behavior: unlike the metals before it, selenium forms negative ions (selenide) and covalent compounds rather than metallic alloys. Its presence in Period 4 highlights the diversity of this period, spanning from reactive metals to essential nonmetals.

Bromine (Br)

Main article: Periodic table/Halogens/Bromine

Bromine (atomic number 35) is a unique element on the periodic table as one of only two elements that are liquid at room temperature (the other being mercury). Bromine is a halogen, in group 17 right below chlorine. At room temperature, pure bromine is a reddish-brown liquid with a noxious, bleach-like odor (the name bromine comes from the Greek bromos, meaning “stench” or “bad odor”). It readily evaporates into a red-brown vapor. Bromine’s boiling point is about 58.8 °C, so on a warm day it will slowly boil away; its melting point is −7.2 °C, so even in a freezer it stays liquid. Seeing bromine is quite striking – it’s a heavy, red “fuming” liquid, often kept in glass ampoules because it is corrosive and its vapors are very toxic.

Because of its reactivity, elemental bromine is not found in nature – it exists as bromide salts, mostly in sea water or brine pools. Bromine is about bromine is found in the ocean as bromide ions (Br⁻), which can be concentrated in salt lakes and brine wells. The Dead Sea, for example, is a source of bromine.

Bromine’s main uses exploit its reactivity to form compounds. A major application is in flame retardants. Many bromine-containing compounds (like polybrominated diphenyl ethers, PBDEs, or newer organobromine compounds) are added to plastics, furniture foam, electronics casings, etc., to make them less flammable. When these materials are heated, the bromine compounds decompose and release bromine radicals that interfere with the combustion reactions, slowing or stopping the fire. For decades, brominated flame retardants have been widespread in consumer products; however, some have raised environmental and health concerns, leading to their phase-out and replacement with alternatives.

Another use of bromine is in certain water treatment chemicals. For instance, bromine tablets are used to disinfect swimming pools and hot tubs (often in combination with chlorine). Bromine is less harsh on the skin and stays effective at higher temperatures, which is why hot tubs often use bromine sanitizer. Bromine (in the form of bromamines) also has a different smell than chlorine-treated pools (some people find it gentler, though bromine has its own strong odor).

Historically, silver bromide (AgBr) was crucial in photography. Silver bromide is light-sensitive; when exposed to light, it forms a latent image that can be developed into a photograph. All those old photographic films and papers depended on silver bromide (and silver chloride/iodide mixtures). That use has waned with digital photography, but it’s a key part of bromine’s legacy.

Ethylene dibromide was once used as an additive in leaded gasoline (to scavenge lead by forming lead bromide that left with exhaust), but with the phase-out of leaded fuel, that use ended.

Bromine compounds also have niche medical uses: in the 19th and early 20th centuries, bromide salts (like potassium bromide) were used as sedatives and anticonvulsants. In fact, the phrase "bromide" came to mean a boring person or platitude, because people on bromide medication were noted to be dulled or sedated. Today, other drugs have replaced bromides in medicine, except in veterinary contexts.

On the darker side, bromine is very toxic and corrosive. Contact with liquid bromine causes severe burns. Its vapors irritate eyes and throat and can damage respiratory organs. Proper precautions are needed when handling even small amounts. The element was discovered in 1826 by Antoine Jérôme Balard, who isolated it from salt water from Montpellier, France.

Chemically, bromine lies between chlorine and iodine in reactivity. It will react with many metals and organics. For instance, bromine will readily bleach dyes (though not as strongly as chlorine) and is used in bromination reactions in organic chemistry (adding bromine to double bonds, etc.). In the ocean, certain algae and sponges produce organobromine compounds (one of the reasons the ocean atmosphere has bromine, which contributes to ozone layer interactions).

In summary, bromine stands out in Period 4 as the only nonmetallic liquid, a heavy, smelly, dangerous but useful element. It’s used to slow down fires, clean up pools, and historically to capture images on film and calm jittery nerves. The vivid image of a red fuming liquid in a flask has made bromine a bit of a poster child for “hazardous chemical,” but in controlled forms it has undoubtedly benefited us in various technologies.

Krypton (Kr)

Main article: Periodic table/Noble gases/Krypton

Krypton (atomic number 36) is the last element in Period 4, residing in group 18 as a noble gas. The name krypton comes from the Greek kryptos, meaning "hidden," reflecting that it is a very rare component of the air and was difficult to discover. Krypton is present in Earth’s atmosphere at about 1 part per million, making it quite scarce (hence “hidden”). It was discovered in 1898 by Sir William Ramsay and Morris Travers during their study of liquefied air; after evaporating other gases, they found krypton in the residue.

In terms of characteristics, krypton is colorless, odorless, and inert (like the other noble gases). It does nothing remarkable chemically because its atoms almost never react or form compounds. Only under extreme conditions (like in the presence of powerful fluorine gas, or inside an electric discharge) will krypton form a few compounds, such as krypton difluoride (KrF₂). All known krypton compounds are unstable and decompose easily. So, for all practical purposes, krypton can be considered chemically inert.

Krypton’s uses stem largely from its physical properties rather than any chemistry. One notable use is in lighting. Krypton gas is used in certain fluorescent lamps, flashbulbs, and high-intensity discharge lamps. For example, older camera flash units (flash photography) often used krypton-filled flash tubes, which when pulsed with high voltage, emit bright white light. Krypton gas is also used in some types of energy-efficient fluorescent bulbs and in “neon” lamps to produce a bright whitish glow (neon produces red-orange light, so for other colors like bright white, other gases including krypton are used). Incandescent bulbs labeled as “krypton bulbs” contain krypton gas; because krypton is a heavy gas, it slows the evaporation of the filament, allowing the bulb to run a bit more efficiently and at higher color temperature than argon-filled bulbs. However, krypton is expensive, so these are usually specialty bulbs.

A very interesting historical role of krypton is that it was used to define the length of a meter for a period of time. From 1960 to 1983, the official standard meter was defined as 1,650,763.73 wavelengths of a specific orange-red spectral line emitted by krypton-86 gas. This spectroscopic definition was later replaced by the current definition (distance light travels in a vacuum in 1/299,792,458 of a second). But for those couple of decades, krypton had the honor of being the physical basis of the meter stick!

Krypton, being a noble gas, is also used in research and technology that exploits inert environments. For example, krypton isotopes can be used in nuclear medicine or in physics experiments. Krypton-85, a radioactive isotope, is released in small amounts from nuclear reprocessing and has been used to track air masses and even to estimate nuclear activities during the Cold War (since krypton-85 is a fission product, spikes in atmospheric krypton-85 could indicate plutonium separation).

In science fiction and pop culture, krypton is best known as the namesake of Superman’s home planet (Krypton) and the infamous "kryptonite". Real krypton gas, however, won’t weaken Superman – it’s harmless unless you’re trapped without oxygen. The association likely comes from the element’s mysterious, hidden nature and cool name.

One more niche use: because krypton has several sharp spectral lines, krypton-neon mixed gas lasers (and the krypton fluoride excimer laser, KrF) have been developed. Krypton fluoride (KrF) lasers emit deep ultraviolet and have been used in high-precision photolithography for making semiconductor chips, as well as in laser fusion research.

Summing up krypton: It’s inert, rare, and quietly useful in lighting and scientific measurement. You’re unlikely to ever encounter krypton gas directly, but its presence is felt in camera flashes, certain fluorescent lamp bulbs, and in the annals of scientific standards. As the full stop of Period 4, krypton ends the period not with a bang but with a silent glow – a fitting close to a row of elements that has everything from explosive metals to poisonous metalloids to invisible gases. Krypton’s very inertness underscores the completion of the valence shell trend across the period, achieving the coveted full-shell stability that the elements far to its left so energetically strive for.

Conclusion

Period 4 of the periodic table is remarkable for the sheer diversity it contains. In this single horizontal row, we encounter metals that can explode on contact with water (potassium), metals that form the backbone of our civilization’s steel (iron) and wiring (copper), subtle metalloids that enabled the transistor revolution (germanium), notorious poisons (arsenic and bromine), and a noble gas that quietly lit our lamps and calibrated our rulers (krypton). The journey from potassium to krypton illustrates the periodic law in action – properties shift from highly reactive to highly inert as we move across, and elements gradually transition from metallic to nonmetallic character.

For an educated observer, Period 4 offers a microcosm of chemistry and its interplay with human life. These 18 elements collectively are involved in critical biological roles (think of calcium in bones and iron in blood), industrial processes (from titanium in airplanes to chromium in stainless steel), technology (gallium and arsenic in LEDs and lasers), and even historical intrigue (the use of cobalt for blue pigments or arsenic in, alas, nefarious deeds). Each element has its story: whether it's the way manganese quietly fortifies our steel structures or how bromine-containing compounds help prevent fires in our homes.

In everyday life, we might not always realize it, but we are constantly interacting with Period 4 elements. When you flip on a light (tungsten filament glowing in an argon/krypton mix, wires of copper, a steel lamp base of iron/carbon, perhaps a brass switch of copper/zinc), or when you take a breath (a trace of krypton enters your lungs), or eat your breakfast (fortified with zinc and iron, containing potassium and calcium), you are benefiting from the unique properties of these elements. Period 4 truly showcases the richness of the periodic table – the way elemental building blocks combine and contribute to the world around us in countless seen and unseen ways.

In summary, Period 4 is where the periodic table really hits its stride: eighteen elements working their magic from the structural metal in skyscrapers to the flashing gases in photography, from the nutrients in our cells to the microchips in our devices. It is a period of workhorses and wonder, anchoring both the ancient and modern aspects of chemistry in our daily experience. Each element, with its key characteristics and uses outlined above, plays a part in the grand tapestry of science, technology, and life. The next time you encounter stainless steel, a blinking LED, a copper penny, or even a simple table salt (which contains a dash of potassium chloride along with sodium chloride), you might remember the contributions of these Period 4 elements – an illustrious lineup in the middle of the periodic table that keeps our world going.