Period 2

Period 2 is the second row of the periodic table, comprising eight chemical elements with atomic numbers 3 through 10. In the periodic table, a period is a horizontal row of elements. All elements in the same period have the same number of atomic electron shells (energy levels). For Period 2 elements, this means each atom has two electron shells occupied by electrons. As we move across Period 2 from left to right (from lithium to neon), we add more protons to the nucleus and more electrons to the outer shell, which leads to trends in their properties. Notably, the atomic size decreases across the period while the tendency to attract electrons (electronegativity) increases, and it becomes harder to remove electrons (higher ionization energy). In simpler terms, elements on the left of Period 2 are larger, more metallic, and easily lose electrons, whereas those on the right are smaller, non-metallic, and tend to gain or hold onto electrons more tightly.

Period 2 is remarkable for the diversity of element types it contains. It starts with two light metals – lithium and beryllium – and then includes a metalloid (boron) followed by four nonmetals (carbon, nitrogen, oxygen, fluorine), and ends with a noble gas (neon). In fact, Period 2 has five nonmetals (B, C, N, O, F), the highest number of nonmetals in any single period. This row showcases a dramatic range of chemical behavior: lithium is a reactive metal, fluorine is the most reactive nonmetal, and neon is an almost completely inert gas. Indeed, many Period 2 elements represent extremes in their respective groups – for example, fluorine is the most reactive halogen, neon is the most inert noble gas, and lithium is the least reactive alkali metal (though lithium will still readily react, it is mild compared to its heavier alkali metal cousins).

In terms of electron configuration, Period 2 corresponds to filling the second electron shell (the L-shell). This shell can hold up to 8 electrons. The progression from lithium to neon involves gradually filling this shell: lithium atoms have just 1 electron in the second shell, while neon atoms have a full complement of 8, achieving a stable "octet". This underpins why neon is so unreactive – it has a full outer shell – whereas elements earlier in the period tend to react in ways that help them attain a full outer shell (for instance, fluorine aggressively gains one electron to complete its octet, while lithium easily loses its single outer electron to drop back to a full first shell).

Overall, Period 2 is where we first really see the classic periodic trends and a rich variety of chemistry. The table below summarizes the Period 2 elements and highlights one key use or property of each:

Element	Symbol	Atomic Number	Key Use or Notable Property
Lithium	Li	3	Lightest metal; used in rechargeable batteries
Beryllium	Be	4	Lightweight metal; used in aerospace materials (satellites, jets)
Boron	B	5	Hard metalloid; used in borosilicate glass (e.g. Pyrex cookware)
Carbon	C	6	Nonmetal; basis of organic life (found in all living organisms)
Nitrogen	N	7	Nonmetal gas; makes up 78% of Earth’s air (essential for plant growth)
Oxygen	O	8	Nonmetal gas; makes up 21% of air, needed for breathing and burning
Fluorine	F	9	Reactive nonmetal; in toothpaste (fluoride prevents cavities)
Neon	Ne	10	Noble gas; used in neon signs (glows orange-red when electrified)

Table: The eight elements of Period 2, with their symbols, atomic numbers, and one prominent use or property of each.

Now, let's take a closer look at each of these elements in Period 2, from lithium, the soft silvery metal at the start of the row, to neon, the colorless noble gas at the end.

Lithium (Li)

Main article: Periodic table/Alkali metals/Lithium

Lithium is a soft, silvery-white metal with atomic number 3. It belongs to the alkali metals (Group 1) and is the lightest metal in the entire periodic table. In fact, lithium’s density is so low (about half that of water) that it can float on light oils. Freshly cut lithium is shiny silver, but it tarnishes quickly in moist air as it reacts with oxygen or water vapor. Like other alkali metals, lithium is highly reactive, though it is actually the least reactive alkali metal – it will fizz and slowly corrode in water (producing hydrogen gas), whereas its heavier cousins like sodium or potassium can explode on contact with water. Because its atoms have only one electron in the outer shell, lithium readily loses that electron to form positive Li⁺ ions. This tendency underpins many of lithium’s chemical behaviors.

Despite its reactivity, lithium has become extremely important in modern technology. Rechargeable lithium-ion batteries rely on lithium ions moving between electrodes to store and release energy, and these batteries power our smartphones, laptops, and electric cars. The proliferation of portable electronics and electric vehicles has made lithium a highly sought-after resource. Lithium metal is also used in certain high-performance alloys; for example, adding lithium to aluminum or magnesium alloys can produce strong, lightweight materials used in aircraft and spacecraft. Another interesting use is in medicine: various lithium compounds (like lithium carbonate) are used as mood stabilizers to treat bipolar disorder and depression. This medical application was one of the first uses of lithium, dating back to the mid-19th century, and today lithium salts are still prescribed (in carefully controlled doses) to help stabilize mood swings. Overall, lithium is an element that bridges chemistry and everyday life – from the battery in your phone to important medications – all while being a fundamental member of the periodic table’s second row.

Beryllium (Be)

Main article: Periodic table/Alkaline earth metals/Beryllium

Beryllium, atomic number 4, is a hard, grayish metal that is much less well-known in daily life than lithium. It is an alkaline earth metal (Group 2 of the periodic table). Beryllium is relatively lightweight (less than twice the density of water) and has a very high melting point for a light metal. In its pure form, beryllium is steel-gray and brittle. One of its remarkable qualities is stiffness: beryllium is exceptionally rigid for such a light metal, and it maintains its dimensions across a wide temperature range (meaning it doesn’t expand or contract very much with heat). This combination of light weight, stiffness, and stability makes beryllium valuable in high-tech applications. For example, beryllium is used in aerospace and defense industries: structural components in high-speed aircraft, missiles, and communication satellites often contain beryllium alloys. It’s also used in the aerospace sector for making lightweight precision instruments and in the construction of the James Webb Space Telescope’s mirrors (which are made of beryllium coated with gold, taking advantage of beryllium’s lightness and stability).

Beryllium’s usefulness doesn’t end there. When alloyed with copper (forming beryllium-copper alloy), it creates a material that is both strong and a very good conductor of electricity and heat. Beryllium-copper springs and electrical contacts are used in applications that demand reliability and conductivity, like in electrical connectors and precision instruments. Another interesting use of beryllium is in X-ray equipment: thin sheets of beryllium are used as windows for X-ray tubes and detectors. Beryllium is transparent to X-rays, so it allows X-rays to pass through while blocking visible light and other unwanted radiation. This property makes it ideal for medical and scientific X-ray instruments.

On a more everyday note, beryllium is present in some gemstones. The mineral beryl is a beryllium-containing mineral (beryllium aluminum silicate), and its precious varieties include emerald (the green color in emerald comes from traces of chromium, but the crystal structure is built on beryllium) and aquamarine. So, beryllium is literally at the heart of those gemstones. It’s worth mentioning that beryllium and its compounds are toxic if inhaled or ingested – prolonged exposure can cause a serious lung condition called chronic berylliosis. Because of this, beryllium is handled with care in industrial settings. While beryllium doesn’t play a biological role in living organisms, its unique physical and chemical properties secure it a small but significant place in technology and industry.

Boron (B)

Main article: Periodic table/Boron group/Boron

Boron (atomic number 5) is a metalloid – an element with properties in between metals and nonmetals. Pure boron is uncommon; when obtained, it can exist as a dark amorphous powder or in a hard, crystalline form. In fact, crystalline boron is extremely hard (comparable to diamond in hardness) and has a very high melting point. Boron is perhaps most famous for its compounds, which have a wide array of uses in daily life and industry. Even if people don’t encounter pure boron, they certainly come across boron compounds regularly.

One of the most important boron compounds is borosilicate glass, which contains boron oxide. If you’ve used a Pyrex baking dish or a laboratory beaker, you’ve used borosilicate glass. The addition of boron makes glass resistant to thermal shock – it won’t crack under rapid temperature changes – and more chemically durable. This heat-resistant glass is invaluable for cookware and laboratory equipment. Another common boron compound is borax (sodium borate), historically used as a laundry booster and cleaning product. Borax and its cousin boric acid have mild antiseptic and cleansing properties; boric acid is found in some eye drops and antiseptics, and borax has been used in detergents and even as a household remedy.

Boron’s usefulness extends to tougher applications as well. Boron combined with carbon forms boron carbide, a ceramic material that is extraordinarily hard – it’s used in bulletproof vests and tank armor to provide lightweight protection. Boron is also a component in advanced neodymium magnets (NdFeB magnets, which include neodymium, iron, and boron) – these are the super-strong permanent magnets found in everything from computer hard drives to cordless power tools and electric vehicle motors. Additionally, boron compounds are used as flame retardants, in electronics (as doping material for semiconductors), and even in nuclear reactors: the isotope boron-10 is very good at capturing neutrons and is used in control rods and radiation shields in the nuclear industry.

In the natural world, boron is a bit scarce but not absent. It’s an essential trace element for plants, needed in small amounts for proper cell wall formation and other functions. In humans and animals, boron is not classified as essential, but researchers have found it may play a role in bone health and other processes (though we typically get all needed boron from our diet without thinking about it). To sum up, boron is a hard-working element – from kitchenware and soap to armor plating and electronics, its compounds find use across a surprising spectrum of activities.

Carbon (C)

Main article: Periodic table/Carbon group/Carbon

Carbon, atomic number 6, is often called the "king of the elements" for life, and with good reason: carbon is the foundational element for all known life on Earth. Every molecule in your body that is part of living tissue – DNA, proteins, carbohydrates, fats – contains carbon. In fact, about 45–50% of the dry mass of living organisms is carbon. Carbon has a unique ability to form stable bonds with itself and with many other elements, creating an incredible variety of complex molecules. Over 10 million different carbon compounds have been documented, which is the domain of organic chemistry, and many millions more are theoretically possible. This versatility underlies carbon’s central role in biology and also in a vast range of materials and chemicals around us.

In its pure elemental form, carbon shows very different faces. It exists as diamond, one of the hardest substances known (transparent and sparkling, used in jewelry and cutting tools), and as graphite, one of the softest (the black, slippery material in pencil "lead" that marks on paper). These vastly different properties arise because of different atomic arrangements: in diamond, each carbon atom is bonded strongly to four others in a 3D lattice; in graphite, carbon atoms form flat sheets that can slide over one another. Modern science has revealed even more forms of carbon, like graphene (a single-atom-thick sheet of carbon, incredibly strong and conductive) and fullerenes (soccer-ball-shaped carbon molecules), which are leading to new technologies in electronics and nanotechnology.

Carbon’s importance extends beyond biology. Fossil fuels such as coal, oil, and natural gas are primarily carbon-based and have powered the industrial world (though at the cost of releasing carbon dioxide, CO₂, which contributes to climate change). Plastics and synthetic fibers are made of carbon compounds (usually derived from petrochemicals). Steel, the backbone of infrastructure, is an alloy of iron with a small percentage of carbon – the carbon dramatically increases the strength of iron to make steel. In the atmosphere, carbon appears in the form of carbon dioxide (CO₂), which plants use for photosynthesis (producing oxygen and sugars) and which helps regulate Earth’s temperature via the greenhouse effect. There’s also carbon in the oceans, rocks (limestone is calcium carbonate), and so on – the circulation of carbon through living things, air, water, and earth is known as the carbon cycle.

In everyday life, carbon is everywhere: the graphite in pencils, the diamond in an engagement ring, the rubber in tires (made from carbon-rich polymers), the gasoline in cars, the carbon fiber in a high-end bicycle frame, the fizz in a soda (carbonation is CO₂ gas). Few elements touch our lives as pervasively as carbon. It’s often said that carbon is common (which it is, being the fourth most abundant element in the universe) but also extraordinary. Its ability to connect and form structures underlies the complexity of life and the diversity of materials that define modern civilization.

Nitrogen (N)

Main article: Periodic table/Pnictogens/Nitrogen

Nitrogen is a nonmetal element with atomic number 7, and under normal conditions it is a transparent, colorless gas. If you take a breath of air, nearly four-fifths of it is nitrogen – roughly 78% of Earth’s atmosphere is nitrogen gas (N₂). This makes nitrogen the most abundant uncombined element in the air. However, nitrogen gas is quite inert (non-reactive) at room temperature, because the two nitrogen atoms in N₂ are bonded by a very strong triple bond. We breathe in nitrogen with every breath, but it simply goes in and out of our lungs without being used directly (our bodies have no easy way to break that strong bond). Nonetheless, nitrogen is absolutely essential to life: it’s a critical component of amino acids and proteins, of DNA and RNA, and of other biomolecules. Plants and animals cannot use nitrogen gas directly; instead, nitrogen must be "fixed" (converted into compounds like ammonia or nitrates) for organisms to incorporate it. In nature, certain bacteria can fix atmospheric nitrogen, and phenomena like lightning can also create reactive nitrogen compounds.

The importance of nitrogen for plants is immense – it’s often the limiting nutrient for plant growth, which is why farmers and gardeners add nitrogen in the form of fertilizers. In the early 20th century, chemists developed the Haber-Bosch process to capture nitrogen from the air by combining it with hydrogen to make ammonia (NH₃). This was a game-changer: now we produce over 150 million tons of ammonia fertilizers each year. These fertilizers (ammonium salts, urea, nitrates) have vastly increased agricultural productivity, helping to feed the world’s growing population. However, overuse of nitrogen fertilizers can cause environmental issues, like water pollution (runoff of nitrates can lead to algal blooms in waterways).

In industry and technology, nitrogen gas’s inertness is very useful. Factories use nitrogen to create oxygen-free environments when needed – for example, food packaging often contains nitrogen gas to displace oxygen and keep foods fresh (preventing oxidation or spoilage). Similarly, manufacturers use nitrogen to blanket flammable liquids or sensitive electronics during production. Liquid nitrogen (which is nitrogen cooled to –196 °C, so that it becomes a liquid) is widely used as a cryogenic fluid. Because it is so cold, it can rapidly freeze foods (locking in texture and nutrients in frozen goods) and is used to preserve biological samples like sperm, eggs, or tissues for medical use. You might have seen liquid nitrogen in action when making gourmet ice cream or in science demonstrations where it causes objects like flowers or balloons to freeze and shatter.

Nitrogen also finds its way into explosives – compounds like nitroglycerin, TNT, and ammonium nitrate all contain nitrogen, which can rapidly form gas products and release energy. Moreover, nitrous oxide (N₂O, “laughing gas”) is a nitrogen compound used as an anesthetic and analgesic in dentistry, and also as an engine oxidizer in racing. And let’s not forget the humble blueprint of life: nitrogenous bases in DNA (adenine, guanine, cytosine, thymine) are so-called because they contain nitrogen. In summary, nitrogen is a bit of a paradox: as a gas it is inert and passive, making up the bulk of our atmosphere without fueling our breathing, yet in compounds it is vital, dynamic, and sometimes explosive. Life on Earth and modern agriculture heavily depend on this element that, in its free form, quietly swirls around us every day.

Oxygen (O)

Main article: Periodic table/Chalcogens/Oxygen

Oxygen, atomic number 8, is a nonmetal gas that is famously reactive and indispensable for life as we know it. Oxygen gas (O₂) makes up about 21% of Earth’s atmosphere – it’s the second most abundant component after nitrogen. The fact that our atmosphere contains so much free oxygen is actually extraordinary from a chemical standpoint: oxygen readily reacts with many substances, so it wouldn’t remain in the air at such levels if it weren’t continuously produced by photosynthesis in plants and algae. The oxygen you breathe is largely a product of green plants releasing O₂. For humans and most other animals, oxygen is essential for respiration – our cells use oxygen to help convert the sugars from our food into energy (ATP) in a process that also produces carbon dioxide and water. Without enough oxygen, we cannot survive; that’s why we breathe continuously and why at high altitudes (where oxygen is sparse) or in medical situations, supplemental oxygen may be needed. Pure oxygen in a hospital setting can be literally lifesaving for patients with respiratory issues or during surgery.

Oxygen’s reactive nature is evident in how it supports combustion. Anything that “burns” is reacting with oxygen. If you’ve watched a fire, adding oxygen (for instance, by blowing on coals or using a bellows) makes it burn hotter and brighter; conversely, smothering a fire by cutting off oxygen will extinguish it. Oxygen is the reason iron rusts (forming iron oxide) and why fruits brown when cut (enzymatic reactions with oxygen). It forms compounds with nearly every other element – oxides are ubiquitous in minerals (sand, for example, is silicon dioxide). In fact, by mass, oxygen is the most abundant element in the Earth’s crust (mostly in the form of oxides like silica, alumina, etc.). It also makes up about two-thirds of the mass of the human body (mostly in water).

Beyond breathing and burning, oxygen has a multitude of uses. In industry, pure oxygen is used to boost or enable high-temperature processes. For example, steelmakers blast oxygen into blast furnaces to oxidize and remove impurities from molten iron, an important step in refining steel. Oxy-fuel welding and cutting torches combine oxygen with fuels like acetylene to produce extremely hot flames – a welding torch can cut through thick steel using the power of oxygen-assisted combustion. The medical field uses bottled or piped oxygen for patients. Astronauts and scuba divers rely on oxygen supplies when they’re in environments without breathable air. Oxygen is also a component of the liquid rocket fuel combination: many rockets use liquid oxygen (LOX) as the oxidizer paired with liquid hydrogen or kerosene fuel – when they react, they release enormous energy and thrust. Another interesting form of oxygen is ozone (O₃), a triatomic form. Ozone in the upper atmosphere forms a layer that absorbs harmful UV radiation, protecting life on Earth; at ground level, however, ozone is a pollutant and irritant.

It’s worth noting the double-edged sword of oxygen: we need it to live and it has enabled the evolution of complex life (because aerobic respiration yields much more energy than anaerobic processes). But oxygen is also reactive enough to cause oxidative damage to cells (this is why we need antioxidants in our diet, for example). The presence of oxygen dramatically changed the Earth’s chemistry roughly 2.4 billion years ago (the "Great Oxidation Event"), enabling new life forms while driving others extinct. In our day-to-day experience, though, oxygen is simply the part of air that keeps us alive and makes candles glow and campfires possible. It’s a vigorous element that, in its own way, keeps the world burning brightly.

Fluorine (F)

Main article: Periodic table/Halogens/Fluorine

Fluorine, atomic number 9, is a pale yellow gas and holds the distinction of being the most reactive element on the periodic table. It is a halogen (in the same group as chlorine, bromine, etc.), and even among halogens fluorine is extreme. A molecule of fluorine gas (F₂) will attack almost any other substance it comes in contact with – metals, glass, water, and even many noble gases under the right conditions. Early chemists had a terrible time isolating fluorine because it would react violently with the containers or chemicals used (many suffered injuries or worse in the 19th-century quests to produce elemental fluorine). It wasn’t until 1886 that fluorine gas was successfully isolated by Henri Moissan, who later won a Nobel Prize for that work. The reactivity of fluorine is due to it being the most electronegative element, meaning fluorine atoms attract electrons more strongly than any other atoms. In reactions, fluorine typically snatches electrons to form fluoride ions (F⁻). Because of this aggressive behavior, fluorine gas must be handled with extreme care – it can ignite fuels or cause explosions on contact with many materials.

However, once fluorine is bonded to other elements (forming fluoride compounds), it becomes far less dangerous and in fact incredibly useful. One everyday example is fluoride in toothpaste and drinking water. Compounds like sodium fluoride or tin(II) fluoride are added in tiny amounts to toothpaste and many municipal water supplies because fluoride ions help strengthen tooth enamel and prevent cavities. This practice has dramatically improved dental health in populations over the last several decades. Another ubiquitous fluorine-derived product is Teflon. Teflon is the trade name for PTFE (polytetrafluoroethylene), a polymer where carbon atoms are completely surrounded by fluorine atoms. The carbon-fluorine bond is extremely strong, making PTFE very stable, non-reactive, and slippery. That’s why frying pans coated with Teflon are non-stick (foods and liquids don’t readily adhere to it). Teflon can withstand high temperatures and harsh chemicals, so it’s used not just in cookware but also in cable insulation, waterproof fabrics (like Gore-Tex® which has a PTFE membrane), and many industrial applications.

Fluorine shows up in other surprising places. Certain high-performance plastics and refrigerants contain fluorine. Historically, chlorofluorocarbons (CFCs) were used in refrigerators and aerosol sprays – they are stable compounds that, unfortunately, were found to destroy the ozone layer and are now mostly phased out. In the pharmaceutical industry, many drugs contain fluorine atoms because introducing fluorine can make molecules more bioactive or help them interact with biological targets in useful ways. For example, the common antidepressant Prozac (fluoxetine) and many anesthetics have fluorine in their structures. Fluorine chemistry has also given us Sulfur hexafluoride (SF₆), an inert, non-toxic gas used as an electrical insulator in high-voltage equipment (and it’s so heavy that it can make your voice deep if inhaled, the opposite of helium’s effect – though one should not inhale SF₆ casually due to risks).

It’s interesting to note that while fluorine is so reactive, the element neon right next to it is completely inert – a beautiful illustration of how one extra proton and electron (from F to Ne) can change chemical behavior completely. In nature, fluorine is found in minerals like fluorite (calcium fluoride) and fluorapatite (in bones and teeth). We never find free fluorine gas naturally on Earth – it’s just too reactive. To summarize fluorine: it’s ferociously reactive in elemental form, yet when tamed in compounds, it contributes to many beneficial products, from protecting our teeth to cooking our breakfast with ease. Just don’t mess with it in the raw state!

Neon (Ne)

Main article: Periodic table/Noble gases/Neon

Neon, atomic number 10, is the last element in Period 2 and is a member of the noble gases, renowned for being extremely inert. Neon gas is colorless, odorless, and lighter than air. It forms no stable compounds under normal conditions – neon atoms prefer to stay single and rarely bond to other elements. In fact, neon was historically considered completely inert; only in recent decades have chemists created a few exotic neon compounds under extreme conditions, and even those are very unstable. Because neon doesn’t react with anything under standard conditions, any neon present on Earth has mostly floated away into space over geological time (neon is rare on Earth’s crust and atmosphere, about 18 parts per million in air). It’s more abundant in the universe at large (for example, in stars).

Neon’s claim to fame in everyday life is its brilliant reddish-orange glow in neon signs. When a high-voltage electric current is passed through a low-pressure tube filled with neon gas, the neon atoms get "excited" and then emit light as they return to their normal state. The light from neon is a vivid orange-red color – this is the classic neon sign color seen in old-fashioned advertising signs and city lights. In the early 20th century, neon signs were a huge sensation; Georges Claude, a French engineer, introduced them around 1910, and by the 1920s they were lighting up cities (the famous Las Vegas and Times Square signage owes much to neon). The term "neon lights" has become a general term, though many modern "neon" signs use other gases or LED technology to produce different colors. Pure neon gives off that unmistakable red-orange glow – if you see a neon sign glowing blue or green, it’s typically neon gas mixed with other gases or different noble gases like argon, or uses colored coatings.

Apart from signage, neon is used in some electronic applications. Neon-filled glow lamps were once common as indicator lights (small neon bulbs that glow orange at about 100 volts) in appliance displays and power strips, before LEDs took over. Neon is also one of the gases used in helium-neon lasers, which gave us the familiar red laser pointers and barcode scanners (many of the early low-power lasers were He-Ne lasers emitting a red beam). These have largely been replaced by semiconductor lasers now, but He-Ne lasers are still used in some research and holography because of their stable output. Additionally, neon, due to its inertness, is used as a cryogenic refrigerant in some cases. Liquid neon, which boils at around –246 °C, can provide cooling; it has about 40 times more refrigerating capacity per unit volume than liquid helium, for example. However, neon is more expensive than helium (since neon is only obtained by extracting it from air in small quantities), so it’s used only in specialized applications.

Neon doesn’t play a role in biology – we don’t need it, and it doesn’t interact with living tissue (which is good, as it’s non-toxic and just passes through). Culturally, neon is often associated with a kind of retro-futuristic aesthetic – the neon glow evokes mid-20th century downtowns, art deco, and noir atmospheres, as well as modern artistic signage. In summary, neon sits at the peaceful end of Period 2: a light, noble gas that refuses to react, but earnestly glows when we electrify it, finding its niche in lighting up our night life and technological devices with a warm orange-red hue.

Conclusion

Period 2 of the periodic table may be short, but it packs a punch in terms of chemical diversity and significance. In this single row, we encounter a soft silvery metal that powers our batteries, a stiff gray metal hiding in satellites and emeralds, a hard brown metalloid that helps make glass cookware and clean our laundry, the elemental backbone of all life on Earth, the gas that dominates our atmosphere and nourishes plants, the gas that we breathe for survival and that makes fires roar, the most aggressive element on the planet (tamed for our benefit in toothpaste and non-stick pans), and an inert gas that quietly makes neon signs glow. These eight elements — lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and neon — collectively illustrate the beauty of the periodic table: as atomic number increases, element properties transition in a systematic way from metallic to nonmetallic, from highly reactive to utterly inert. Each element in Period 2 plays a role in the world, whether in the grand scheme of biology and geology or in the conveniences of everyday human life. Together, they remind us how a small step across a period can be a giant leap in chemical behavior, and how the second period’s little atoms have a big impact on our science, technology, and daily experiences.