The rows in the periodic table are called periods. There are seven of them, numbered 1 through 7, and each one represents a new “layer” of electrons surrounding the atoms of the elements it contains. The term is easy to confuse with “groups,” which are the columns, but periods and groups organize elements in fundamentally different ways, and understanding the distinction unlocks a lot of the table’s usefulness.
Where the Name Comes From
The word “period” in this context comes from the same root as “periodic,” meaning recurring at regular intervals. When Dmitri Mendeleev presented his arrangement of the elements to the Russian Chemical Society in March 1869, his key insight was that elemental properties repeat in a predictable pattern when the elements are listed by increasing atomic weight.1The Royal Society. The periodic law of the chemical elements Line them up that way and you notice something striking: reactive metals show up, then less reactive metals, then nonmetals, then a noble gas, and then the cycle starts over. Each complete cycle of that repeating pattern is one period. The table’s name and its row name share a root because the rows are what make the repetition visible.
Mendeleev was not the only chemist to notice the pattern. Others, including John Newlands and Lothar Meyer, spotted similar regularities around the same time. But Mendeleev’s version endured because he was bold enough to leave gaps for elements that had not yet been discovered and to predict their properties. Those predictions turned out to be remarkably accurate, which cemented the periodic table as one of chemistry’s most powerful tools.
What Determines Which Period an Element Sits In
Every element in a given period has the same number of electron shells. Hydrogen and helium, both in Period 1, have one shell. Carbon, nitrogen, and oxygen are in Period 2 and have two shells. Sodium through argon sit in Period 3 with three shells, and so on all the way down to the heaviest known elements in Period 7, which have seven shells. The period number is, in effect, a direct readout of how many layers of electrons wrap around an atom’s nucleus.
This is why periods are horizontal rather than vertical. Moving across a period from left to right, you are adding electrons to the same outermost shell. Moving down a column, you are jumping to a new shell entirely. The two directions encode different kinds of information: periods track the filling of a single shell, while groups (the columns) collect elements whose outermost shells are filled in a similar way, which is why elements in the same group tend to behave alike chemically.
Periods Versus Groups
The distinction trips people up because both are just ways of slicing the same table. The short version: groups sort elements by what they do, and periods sort elements by how big they are. Elements in Group 1, the alkali metals, are all soft, reactive metals that explode on contact with water. They share that behavior because they all have one electron in their outermost shell, even though those shells are at different distances from the nucleus. Elements in Period 3, by contrast, range from the metal sodium to the noble gas argon. They do not behave similarly at all. What they share is the same atomic “size class,” meaning the same number of electron shells.
Teachers often emphasize groups because the chemical similarities are dramatic and easy to demonstrate. But the trends that run across periods are just as important. Atomic radius shrinks as you move left to right within a period, because protons are being added to the nucleus and pulling the electron cloud inward. Electronegativity, the tendency of an atom to attract electrons in a chemical bond, generally increases across a period. Metallic character decreases: the left side of most periods is dominated by metals, the middle by metalloids or transitional elements, and the right side by nonmetals and noble gases. These left-to-right trends are period trends, and they repeat in every row of the table.
Why the Periods Are Not All the Same Length
If you glance at the periodic table, the rows are clearly not uniform. Period 1 has just two elements. Periods 2 and 3 each have eight. Periods 4 and 5 each have eighteen. Periods 6 and 7 each have thirty-two. That uneven pattern puzzles a lot of people who expect a tidy grid, but it falls directly out of the way electron shells are structured.
Each shell is subdivided into smaller regions where electrons can live. The first shell has room for only two electrons, so Period 1 ends after two elements (hydrogen and helium). The second shell can hold eight electrons across two types of subregion, so Period 2 runs for eight elements. By the time you reach Period 4, a third type of subregion becomes available, and the shell can accommodate eighteen electrons. Periods 6 and 7 introduce yet another subregion, expanding to thirty-two elements each. The widening of the periods is not arbitrary; it mirrors the increasing complexity of the electron cloud.
This is also why the periodic table has that distinctive shape with a gap at the top and two long rows typically pulled out and placed below the main body. Those two detached rows, the lanthanides and actinides, are part of Periods 6 and 7 respectively. They are separated purely for visual convenience. If you slotted them back in where they belong, the table would be thirty-two columns wide and far harder to print on a poster or fit on a classroom wall.
Period 1 Is Genuinely Strange
Hydrogen and helium form the shortest period, and both elements are oddballs. Hydrogen is usually placed in Group 1 because it has one electron, but it behaves nothing like the alkali metals below it. Under extreme pressure it can act as a metal, but at everyday conditions it is a lightweight nonmetal gas. Some versions of the table place hydrogen above the halogens in Group 17 instead, or float it separately above the table, acknowledging that it does not fit neatly into any column. Helium, similarly, is chemically a noble gas and behaves like the other elements in Group 18, but its electron configuration would technically place it in Group 2 alongside beryllium and magnesium. Nearly every published table keeps helium with the noble gases because its chemistry matters more than its configuration for practical purposes, but the placement is a genuine point of debate among table designers.
Period 1’s weirdness is a reminder that the periodic table is a human-made organizational tool, not a law of nature itself. The law of nature is the periodicity. The table is one way, arguably the best way, of displaying that periodicity, but edge cases like hydrogen and helium reveal the limits of any two-dimensional grid for representing three-dimensional quantum reality.
The Superheavy Frontier of Period 7
Period 7 was finally completed in 2016 when the International Union of Pure and Applied Chemistry officially named elements 113, 115, 117, and 118: nihonium, moscovium, tennessine, and oganesson. These elements do not exist in nature. They are synthesized in particle accelerators by smashing lighter nuclei together, and they decay in fractions of a second. Studying their chemistry is extraordinarily difficult, but researchers have managed it for a few of them using gas chromatography techniques that work on just a handful of atoms at a time.
What makes Period 7 especially interesting to physicists and chemists is that the elements at its far end experience strong relativistic effects. Their inner electrons move at speeds approaching the speed of light, which alters the shape and energy of their outer electron orbitals in ways that change chemical behavior. Experiments on nihonium and moscovium, for example, showed that both elements interact more weakly with surfaces than the lighter elements above them in the same groups (thallium and bismuth), a departure from simple periodic trends that is well explained by relativistic calculations.2PubMed Central. Manifestation of relativistic effects in the chemical properties of nihonium and moscovium revealed by gas chromatography studies The periodic table still works for these elements, but the trends it suggests need to be corrected for relativistic distortions.
Could There Be an Eighth Period?
No element with an atomic number above 118 has been confirmed, but the question of whether Period 8 could exist is far from idle speculation. Quantum mechanical models predict that elements with atomic numbers 119 through roughly 168 could occupy an eighth period, extending the table significantly.3ResearchGate. The Extended Periodic Table of the Elements: Predicted Physical Constants, Relativistic Electronic Structure, and Laboratory Verification Protocols for Z = 119 to 172 Theoretical work suggests that somewhere around element 172, relativistic effects become so extreme that stable electron configurations may no longer be possible, setting an upper boundary on the periodic table itself.
Synthesizing element 119 is an active goal at several heavy-ion laboratories around the world. The challenge is immense. Each step up in atomic number requires more energetic collisions and produces atoms with shorter lifetimes. Even if element 119 is created, it may survive for only microseconds, making chemical study nearly impossible. Still, the attempt matters because it tests whether quantum mechanics continues to make correct predictions at the edge of known matter. If the chemistry of a Period 8 element deviates from predictions, it could reveal new physics.
The theoretical eighth period would also introduce new types of electron subregions that do not appear in any existing period. This means the table might need to widen further, or designers might need a fundamentally new layout. Some chemists have already proposed three-dimensional or spiral periodic tables to accommodate the growing complexity, though none has displaced the familiar rectangular version.
Common Misconceptions About Periods
One widespread confusion is that elements in the same period should have similar properties. They usually do not. Sodium, a metal you can cut with a butter knife, and chlorine, a toxic yellow-green gas, are in the same period. What they share is a structural feature: three electron shells. Their differences come from how those shells are filled, which is governed by their position within the period (leftward versus rightward). If someone tells you two elements “are in the same period” as though that means they are chemically related, they are confusing periods with groups.
Another misconception is that the periodic table is fixed and settled. In reality, its layout is still debated. The placement of hydrogen and helium is contested, as mentioned earlier. The position of the lanthanides and actinides relative to the main body has been argued over for decades. Even the question of which element belongs in Group 3 underneath scandium and yttrium (is it lanthanum or lutetium?) has prompted formal recommendations from international chemistry bodies, most recently in 2021. The periodic law itself is rock solid, but the best way to represent it visually is still evolving.
Reading the Table Like a Map
Once you know that the rows are periods and the columns are groups, you can use the table as a quick-reference map for a surprising amount of information without memorizing anything. The period number tells you how many electron shells an atom has, and therefore roughly how large it is. The group number for the main-group elements tells you how many electrons are in the outermost shell, which largely determines how the element will react. An element’s position from left to right within its period tells you whether it is likely a metal (left), a metalloid (middle), or a nonmetal (right).
You can also estimate an element’s ionization energy, which is how hard it is to strip away an electron. That tends to increase across a period and decrease down a group. Electronegativity follows the same pattern. These are not perfectly smooth trends; there are bumps and exceptions, particularly among the transition metals in the middle of the table where the subregion-filling order gets complicated. But as a first approximation, knowing an element’s row and column tells you a lot about what it will do.
The Lanthanide and Actinide Rows
Those two floating rows at the bottom of most periodic table posters are probably the most misunderstood feature of the whole layout. They are not separate from the main table. The lanthanides (elements 57 through 71) belong in Period 6, and the actinides (elements 89 through 103) belong in Period 7. They are pulled out and displayed below for a purely practical reason: fitting thirty-two columns on a page or wall chart makes the table too wide to be useful.
The lanthanides are sometimes called rare earth elements, though most of them are not actually rare. Cerium, for example, is about as abundant in the Earth’s crust as copper. The name persists from an era when they were difficult to separate from one another, not because they were scarce. What makes the lanthanides distinctive is that they are filling a deep inner subregion of their electron shell, which means their outermost electrons barely change from one element to the next. This gives them famously similar chemical properties and made them a nightmare to isolate historically.
The actinides include the well-known radioactive elements uranium, plutonium, and thorium, along with a string of synthetic elements created in laboratories. All actinides are radioactive, though thorium and uranium are stable enough to persist in nature. The later actinides, like einsteinium and fermium, have been produced in such tiny quantities that their chemistry is barely characterized. They belong to Period 7, and their existence is part of what made completing that period’s map such a long and painstaking effort.
Why Some Periodic Tables Look Different
If you have ever searched for a periodic table online and noticed that some versions color-code differently, place hydrogen in a different spot, or stretch the table out to a full thirty-two columns, you are not imagining things. There is no single “official” version of the periodic table. The International Union of Pure and Applied Chemistry publishes a widely used version, but it does not mandate that everyone use the same layout. Over 700 alternative periodic tables have been proposed since Mendeleev’s time, including spiral versions, three-dimensional models, and even versions arranged in concentric circles.
The standard rectangular form persists because it is compact, easy to print, and good enough for most purposes. But it does obscure some relationships. A spiral table, for instance, can show more clearly that the periods are not discrete rows but one continuous sequence of elements that loops back on itself. A three-dimensional table can separate the s, p, d, and f blocks into different planes, making the electron-filling order more intuitive. None of these has caught on broadly, partly because the rectangular table is so deeply embedded in textbooks, lab walls, and cultural memory. Still, the variety of alternatives is a healthy reminder that the periodic table is a tool for organizing information, and like any tool, it can be redesigned when a better layout serves a particular purpose.