The right side of the periodic table is home to the nonmetals, metalloids, halogens, and noble gases. Roughly speaking, everything from about group 13 onward (and certainly groups 15 through 18) belongs to this territory, which includes some of the most familiar substances on Earth: the oxygen you breathe, the nitrogen that makes up most of the atmosphere, the carbon in every living molecule, and the neon glowing in storefront signs. These elements share a broad tendency to gain or share electrons rather than give them up, which makes their chemistry fundamentally different from the metals that dominate the left and center of the table.
What Counts as the “Right Side”
The periodic table is arranged so that elements with similar chemical behavior fall into vertical columns called groups. Moving from left to right across any row, elements shift from strongly metallic to nonmetallic. There is no official column number where “right” begins, but the most common understanding places the boundary somewhere around group 13 or 14. By the time you reach groups 15 through 18, you are firmly in nonmetal territory for the lighter elements, though heavier elements lower in those same groups can behave more like metals.
The key groups on the right side, working from left to right, include group 14 (the carbon group), group 15 (the nitrogen or pnictogen group), group 16 (the oxygen or chalcogen group), group 17 (the halogens), and group 18 (the noble gases). A diagonal band of metalloids, elements that straddle the line between metal and nonmetal, runs through several of these groups. Below that band, post-transition metals like tin, lead, and bismuth sit in the same columns but behave quite differently from their lighter cousins.
The Reactive Nonmetals
The nonmetals clustered in the upper-right portion of the table are among the most chemically active elements known. Carbon, nitrogen, oxygen, phosphorus, sulfur, selenium, fluorine, chlorine, bromine, and iodine all fall into this zone. What unites them is their tendency to attract electrons from other atoms, forming covalent bonds with each other or ionic bonds with metals. That electron-grabbing ability, broadly called electronegativity, increases as you move up and to the right across the table, which is why fluorine, perched in the top-right corner of the reactive elements, is the single most electronegative element in existence.
These nonmetals are extraordinarily versatile. Oxygen and nitrogen are gases at room temperature and make up about 99 percent of dry air by volume. Carbon forms more known compounds than all other elements combined, thanks to its ability to bond with itself in long chains and rings. Sulfur and phosphorus are solids under normal conditions, and both play critical roles in industry and biology alike. The elements essential for life as we know it are frequently summarized by the acronym CHNOPS, standing for carbon, hydrogen, nitrogen, oxygen, phosphorus, and sulfur, and every one of those except hydrogen sits on the right side of the table.1Frontiers in Cell and Developmental Biology. Metals Are Integral to Life as We Know It
One thing that sometimes surprises people is how different these nonmetals are from one another despite being neighbors. Carbon is a solid that can form diamond. Nitrogen is an almost inert gas at room temperature. Oxygen is a gas that supports combustion so aggressively it can corrode iron. Their shared trait is not a shared personality but a shared position on the electronegativity scale: they all prefer to take electrons rather than surrender them.
The Halogens in Group 17
Group 17, one column from the far right, contains fluorine, chlorine, bromine, iodine, astatine, and tennessine. These are the halogens, a name derived from Greek words meaning “salt-forming,” because they readily react with metals to produce salts like sodium chloride (table salt). The halogens have compact, strongly electron-attracting open valence electron shells, which explains why they are among the most reactive nonmetals.2Frontiers in Chemistry. Understanding Periodic and Non-periodic Chemistry in Periodic Tables
Each halogen needs just one more electron to fill its outer shell, making them ferocious oxidizers. Fluorine is so reactive that it attacks almost anything it touches, including materials that resist other chemicals. Chlorine is a pale green gas used in water treatment because it destroys bacteria and other pathogens. Bromine is one of only two elements that are liquid at room temperature (mercury is the other). Iodine is a dark solid that sublimes into a distinctive purple vapor. The heavier halogens, astatine and tennessine, are radioactive and exist only in tiny quantities or for fractions of a second, so their practical chemistry is essentially academic.
Beyond salt-making, halogens have found uses ranging from Teflon coatings (which rely on fluorine-carbon bonds) to pharmaceutical design. In catalysis research, halogen-bond interactions are increasingly used to speed up chemical reactions. Studies have shown that halogen-bond donors can lower energy barriers in certain organic reactions, making them useful tools for chemists designing new synthetic routes.3PubMed. Organocatalysis by Halogen, Chalcogen, and Pnictogen Bond Donors in Halide Abstraction Reactions: An Alternative to Hydrogen Bond-Based Catalysis
The Noble Gases in Group 18
The far-right column of the periodic table, group 18, belongs to the noble gases: helium, neon, argon, krypton, xenon, radon, and oganesson. Their discovery in the late 19th century by William Ramsay and others added an entirely new group to the periodic table and provided a keystone to understanding how electrons bind atoms into molecules.4PubMed Central. Sir William Ramsay and the noble gases Before their discovery, no one realized a complete family of elements had been hiding in plain sight, too chemically aloof to participate in ordinary reactions.
The noble gases earned their name because they were once believed to be entirely unreactive, like aloof royalty that refused to mingle. Their outer electron shells are completely full, which means they have no strong drive to gain, lose, or share electrons. In practice, however, “unreactive” turned out to be an overstatement. In 1962, Neil Bartlett showed that xenon could form compounds with fluorine, cracking open a small but genuine field of noble-gas chemistry. Krypton compounds have since been made as well, though helium, neon, and argon remain stubbornly inert under all but the most extreme laboratory conditions.
Despite their low reactivity, noble gases are immensely useful. Argon fills incandescent light bulbs and is used as a shielding gas in welding. Helium cools superconducting magnets in MRI machines and, of course, fills party balloons. Neon lights up signs with its characteristic red-orange glow, while xenon is used in high-intensity lamps and ion propulsion engines for spacecraft. Radon, the radioactive member of the family, is a health hazard that seeps into basements from natural uranium decay in soil and rock, making it a leading environmental cause of lung cancer.
Metalloids Along the Staircase
A zigzag line, sometimes drawn as a staircase, cuts diagonally through the right side of the table separating metals below and to the left from nonmetals above and to the right. The elements that sit right along this boundary, including boron, silicon, germanium, arsenic, antimony, and tellurium, are called metalloids. They share some properties with metals (they can conduct electricity, for instance) and some with nonmetals (they tend to be brittle and form covalent bonds).
Silicon and germanium are the most famous metalloids because they are semiconductors, the backbone of the modern electronics industry. Silicon dominates computing, but germanium has historically been important in transistor technology and infrared optics. Recent computational research has explored novel crystal structures for germanium that could give it a direct bandgap, a property that would open the door to optoelectronic applications that its standard form cannot support. One study identified a hexagonal crystal structure for germanium with energy only slightly above its ground state, yielding a direct bandgap that could someday be tuned by alloying with silicon to rival the performance of established semiconductor compounds.5Computational Materials Today. Unveiling Novel Direct Bandgap Allotropes of Germanium: A Computational Exploration
Arsenic, despite its reputation as a poison, is used in small quantities in semiconductor manufacturing as well. Antimony hardens lead alloys for batteries. Tellurium appears in advanced solar cells and thermoelectric devices. These metalloids are technologically essential precisely because they live in the gray zone between metal and nonmetal, letting engineers control how and when they conduct electricity.
The Chalcogens and Pnictogens
Groups 15 and 16 deserve their own discussion because they illustrate a phenomenon unique to the right side of the table: the same column can contain a gas, a metalloid, and a metal. Group 15, the pnictogens, includes nitrogen (a gas), phosphorus (a reactive solid nonmetal), arsenic (a metalloid), antimony (another metalloid with increasing metallic character), and bismuth (a post-transition metal). Group 16, the chalcogens, includes oxygen (a gas), sulfur (a nonmetal solid), selenium (a nonmetal with some metalloid traits), tellurium (a metalloid), and polonium (a radioactive metal).
This vertical transition from nonmetal to metal within a single group is one of the defining features of the right side of the table. It happens because as you move down a group, atoms get larger, their outer electrons are held less tightly, and metallic character increases. The pattern is much less dramatic on the left side, where even the lightest elements like lithium and beryllium are already metals.
Both groups have drawn interest in catalysis. The noncovalent interactions formed by chalcogen and pnictogen atoms, sometimes called sigma-hole interactions, can be used to catalyze organic reactions. In halide abstraction reactions, for example, pnictogen-bond donors have been found to outperform both chalcogen-bond and halogen-bond donors, with catalytic activity increasing as you move from lighter to heavier elements within each group.3PubMed. Organocatalysis by Halogen, Chalcogen, and Pnictogen Bond Donors in Halide Abstraction Reactions: An Alternative to Hydrogen Bond-Based Catalysis In other reaction types, the ordering can reverse: a separate computational study on a classic organic reaction found that halogen-bonded catalysts actually produced the lowest energy barriers, while pnictogen-bonded catalysts had the least catalytic effect.6PubMed Central. Lewis Acid-Catalyzed Diels-Alder Reactions: Reactivity Trends across the Periodic Table The takeaway is that these right-side elements are far from chemically boring; their interactions are subtle, tunable, and increasingly exploited in modern chemistry.
Post-Transition Metals on the Right Side
Below the metalloid staircase, you find a cluster of metals that behave differently from the familiar transition metals like iron and copper. These post-transition metals, including aluminum, gallium, indium, thallium, tin, lead, and bismuth, occupy groups 13 through 15 in the lower rows. They tend to be softer, have lower melting points, and conduct electricity less efficiently than transition metals. Several of them, particularly gallium, indium, and tin, have melting points low enough that they can be liquid near or below the boiling point of water.
Lead and tin have been used by human civilizations for thousands of years. Lead’s toxicity is now well understood, and its use has been curtailed in paints, gasoline, and plumbing. Tin remains important in soldering and as a coating for steel cans. Bismuth, the heaviest stable element, has found a second life in pharmaceuticals (it is the active ingredient in some stomach-relief medications) and as a lead replacement in shotgun pellets and fishing sinkers.
Gallium and indium are critical materials in the semiconductor industry. Gallium arsenide and gallium nitride are used in LEDs, laser diodes, and high-frequency electronics. Indium tin oxide is the transparent conducting film that makes touchscreens work. These post-transition metals may not command the same name recognition as iron or gold, but modern electronics would collapse without them.
Superheavy Elements at the Bottom Right
The very bottom of the right side of the periodic table includes a set of superheavy elements, those with atomic numbers 113 through 118: nihonium, flerovium, moscovium, livermorium, tennessine, and oganesson. All of them are synthetic, produced in particle accelerators, and most exist for only fractions of a second before decaying. Their positions on the table place them in the same columns as familiar elements like thallium, lead, bismuth, polonium, astatine, and radon, but there is good reason to think they do not behave the same way.
At these extreme atomic numbers, the electrons closest to the nucleus move at a significant fraction of the speed of light, which changes how electron shells fill and interact. Researchers expect that the chemical properties of these superheavy elements may not align with what their periodic-table positions suggest. Flerovium, element 114, is a striking example: it sits in the same column as lead, but theoretical calculations predict it may behave more like a noble gas because of how its outermost electrons pair up into an unusually stable configuration.7Journal of Physics G: Nuclear and Particle Physics. Comparing the fluorination of Pb+ and Po+: insights into p-block relativistic effects for superheavy elements If confirmed experimentally, this would mean the periodic table’s predictive power starts to break down at the heaviest end of its right side.
Oganesson, element 118, sits in the noble-gas column, but theoretical work suggests it may actually be a reactive solid rather than an inert gas. These predictions remain largely untested because the atoms decay too quickly and are produced in quantities too small for traditional chemistry experiments. The superheavy region of the right side is, in a real sense, the frontier where the periodic table’s organizing logic is being tested and, in places, found wanting.
Why Nonmetals Cluster on the Right
The pattern is not accidental. Moving left to right across any row of the periodic table, the number of protons in the nucleus increases while the outer electrons remain at roughly the same distance. This means each successive element holds its electrons more tightly. By the time you reach the right side, the atoms are small, their electron clouds are compact, and they strongly resist losing electrons. Instead, they tend to gain electrons or share them through covalent bonds. That is why the right side is dominated by nonmetals and why reactivity patterns flip compared to the left side: the most reactive metal (francium or cesium, at the bottom left) gives up electrons most easily, while the most reactive nonmetal (fluorine, at the top right) grabs them most aggressively.
This also explains the staircase of metalloids. As you descend through any of the right-side groups, atoms get larger and their outer electrons are held more loosely, so metallic character gradually emerges. The transition from nonmetal at the top of a group to metal at the bottom is smooth, and the metalloids are simply the elements caught in the middle of that gradient.
Everyday Encounters with Right-Side Elements
You interact with right-side elements constantly, often without realizing it. The oxygen you breathe and the carbon dioxide you exhale are both made entirely of right-side elements (oxygen and carbon in one case, with the same two in the other). The nitrogen in fertilizers that grow your food, the chlorine that disinfects your tap water, the silicon in your phone’s processor, the sulfur in the vulcanized rubber of your car tires, and the iodine added to table salt for thyroid health are all right-side elements doing essential work.
Phosphorus deserves special mention. It appears in DNA and RNA as part of the sugar-phosphate backbone, in ATP (the energy currency of cells), in bones and teeth as calcium phosphate, and in the match heads you strike to start a fire. Few elements span such a range of roles, from the molecular machinery of life to industrial pyrotechnics.
Noble gases show up in less obvious ways. The argon in double-pane windows insulates your house. Xenon headlamps on high-end cars produce a bright, bluish-white light. Krypton is used in certain photographic flash equipment and in the fluorescent tubes of energy-efficient lighting. Even helium, best known for balloons and squeaky voices, is a nonrenewable resource critical to scientific instruments, cryogenics, and leak detection systems.
Common Misconceptions About the Right Side
One widespread misunderstanding is that all nonmetals are gases. In reality, only a handful are gases at room temperature: hydrogen (on the far left, an exception to the general rule), nitrogen, oxygen, fluorine, chlorine, and the noble gases. Carbon, phosphorus, sulfur, selenium, and iodine are all solids. Bromine, as mentioned, is a liquid. The state of matter depends on the strength of intermolecular forces, not just on whether an element is a metal or nonmetal.
Another misconception is that noble gases form no compounds at all. While helium, neon, and argon remain genuinely inert under any conditions achievable in a normal lab, xenon forms a variety of stable fluorides and oxides, and krypton difluoride can be made under the right conditions. The “noble” label indicates reluctance to react, not an absolute prohibition.
A third common confusion involves hydrogen. It appears in group 1 on most periodic tables, alongside the alkali metals, but it is a nonmetal that behaves nothing like lithium or sodium under ordinary conditions. Some table designs place hydrogen on both sides or float it above the table entirely to signal its ambiguous classification. In terms of its chemistry, hydrogen often acts more like a halogen than an alkali metal, gaining an electron to form a negative ion in metal hydrides. Its placement is one of the longest-running debates in periodic-table design.
How Right-Side Elements Interact with Metals
The most straightforward chemistry in the periodic table happens when elements from opposite sides meet. Metals on the left readily give up electrons; nonmetals on the right readily accept them. The result is ionic compounds: salts, oxides, and similar materials held together by the electrostatic attraction between positive and negative ions. Sodium chloride is the textbook example, with sodium from group 1 handing an electron to chlorine from group 17.
When right-side elements interact with each other, the bonding is covalent, meaning the atoms share electrons rather than transferring them outright. Water is a covalent compound of oxygen and hydrogen. Carbon dioxide is covalent. The organic molecules that make up your body are vast covalent networks built mostly from carbon, hydrogen, oxygen, nitrogen, phosphorus, and sulfur. This capacity for complex covalent bonding is one reason life depends so heavily on right-side elements: they can build intricate molecular architectures that ionic bonds simply cannot achieve.
Some right-side elements also form coordinate bonds and other noncovalent interactions that are increasingly being harnessed in materials science and drug design. The sigma-hole interactions mentioned in the context of halogen, chalcogen, and pnictogen bonds are a case in point: these relatively weak but tunable forces allow chemists to design catalysts, crystal structures, and molecular recognition systems with remarkable precision.