The periodic table tells you the identity, weight, and chemical behavior of every known element, all 118 of them, through a compact set of numbers and a carefully designed layout. Each element’s square contains specific data points, while the element’s position on the table reveals how it behaves, what it reacts with, and how its properties compare to neighboring elements.
What Each Element Square Contains
Every square on the periodic table displays a few core pieces of information. The atomic number, typically shown at the top of the square, tells you how many protons are in the nucleus of that element. Hydrogen is 1, helium is 2, carbon is 6, and so on up to oganesson at 118. Because atoms are electrically neutral, the atomic number also tells you the number of electrons orbiting the nucleus.
Below or beside the atomic number, you’ll find the element’s chemical symbol (H for hydrogen, Fe for iron, Au for gold) and its full name. At the bottom of the square sits the atomic mass, measured in atomic mass units. This number represents the weighted average mass of all the element’s naturally occurring forms. Carbon, for example, has an atomic mass of 12.011. If you round that mass to the nearest whole number, you can figure out how many neutrons the element’s most common form contains by subtracting the atomic number. Carbon: 12 minus 6 equals 6 neutrons.
Some modern periodic tables also color-code squares to indicate whether an element is a metal, nonmetal, or metalloid, and whether it’s a solid, liquid, or gas at room temperature. Only two elements, mercury and bromine, are liquids under normal conditions. Eleven elements are gases, including all six noble gases plus hydrogen, nitrogen, oxygen, fluorine, and chlorine. Everything else is solid.
Why Atomic Weights Sometimes Appear as Ranges
If you look closely at a current periodic table, you may notice that twelve elements list their atomic weight as an interval rather than a single number. This happens because those elements have multiple stable forms (isotopes) whose proportions vary depending on where in the world the sample was collected. A chunk of sulfur from a volcanic vent may have a slightly different atomic weight than sulfur from seawater. IUPAC, the international authority on chemical nomenclature, adopted interval notation to reflect this natural variability rather than give a single number that could be misleading for precise work.
What Rows and Columns Tell You
The table’s structure is where the real power lies. Horizontal rows are called periods, and vertical columns are called groups (or families). Each carries distinct information about an element’s electron structure and chemical personality.
The period number tells you how many electron shells an element has. Sodium sits in period 3, so its electrons occupy three energy levels. Cesium is in period 6 and has six. As you move down a group, each new period adds another shell, which means atoms get physically larger.
Groups are arguably more useful for predicting behavior. Elements in the same group have the same number of electrons in their outermost shell, which is what determines how an element reacts. Group 1 elements (the alkali metals like lithium, sodium, and potassium) all have one outer electron and are extremely reactive, readily giving it up to form compounds. Group 17 (the halogens, including fluorine, chlorine, and bromine) have seven outer electrons and are also highly reactive, aggressively pulling in one more electron to complete their shell. Group 18, the noble gases (helium, neon, argon, krypton, xenon, radon), already have a full outer shell of eight electrons, which is why they almost never react with anything.
For the main-group elements in groups 1, 2, and 13 through 18, the group number directly indicates how many electrons an element will lose or gain when forming chemical bonds. A group 2 element like magnesium loses two electrons. A group 16 element like oxygen gains two.
The Shape of the Table Maps Electron Structure
The periodic table’s distinctive shape, wide in some places and narrow in others, isn’t arbitrary. It reflects four types of electron orbitals, labeled s, p, d, and f, which describe the regions where electrons are likely to be found.
Groups 1 and 2, the two leftmost columns, form the s-block. Elements here have their outermost electrons in s orbitals. Groups 13 through 18, on the right side, are the p-block, where the outermost electrons fill p orbitals. The transition metals in groups 3 through 12 make up the d-block, where electrons fill d orbitals one energy level below the outermost shell. The two long rows typically separated out at the bottom of the table (the lanthanides and actinides) are the f-block, where electrons fill f orbitals two levels below the outermost shell.
This means you can determine an element’s electron configuration just by reading its position. An element in the fourth period of the d-block is filling its 3d orbital. An element in the sixth period of the p-block is filling its 6p orbital. The table itself is essentially a map of how electrons arrange themselves.
Metals, Nonmetals, and the Staircase Between
A diagonal staircase line cuts through the right side of the periodic table, roughly from boron down to astatine. Everything to the left of this line is a metal. Everything to the upper right is a nonmetal. Elements touching the line, such as silicon, germanium, and arsenic, are metalloids, meaning they share properties of both categories.
This division follows a pattern: within any group in the p-block, elements become more metallic as you move down. Carbon at the top of group 14 is a nonmetal. Silicon below it is a metalloid. Tin and lead at the bottom are metals. Moving left to right across a period, the opposite happens: elements shift from metallic to nonmetallic behavior. This gradient is one of the periodic table’s most practical features, letting you estimate whether an unfamiliar element conducts electricity, forms positive or negative ions, or behaves more like a metal or an insulator.
Predictable Trends Across the Table
Beyond individual element data, the table reveals patterns in how properties change from one element to the next. Three trends are especially useful.
Atomic size decreases as you move left to right across a period. This happens because each step to the right adds a proton to the nucleus, pulling the electron cloud inward, while the shielding effect from inner electrons stays roughly the same. Moving down a group, atomic size increases because each new period adds another electron shell farther from the nucleus. Cesium, in the bottom left, is one of the largest atoms. Fluorine, in the upper right, is one of the smallest.
Ionization energy, the energy needed to strip away an electron, follows the opposite pattern. It increases across a period (electrons are held more tightly as nuclear charge grows) and decreases down a group (outer electrons sit farther from the nucleus and are easier to remove). This is why alkali metals at the bottom left of the table lose electrons so readily, while noble gases at the upper right hold onto theirs tightly.
Electronegativity, an element’s ability to attract electrons in a chemical bond, also increases across a period and decreases down a group. Fluorine, in the upper right corner of the reactive elements, is the most electronegative element on the table. Francium, at the bottom left, is the least. When two elements with very different electronegativities bond, the more electronegative one hogs the shared electrons, creating a polar or ionic bond. You can estimate this just by checking how far apart the two elements sit on the table.
Putting It All Together
Pick any element on the table and its position alone tells you its number of protons and electrons, how many electron shells it has, what type of orbital its outermost electrons occupy, whether it’s a metal or nonmetal, how large its atoms are relative to neighbors, how reactive it is, and which other elements it behaves most like. The numbers printed in its square let you calculate its neutron count, predict its mass in chemical reactions, and identify its isotopic variability. No other single chart in science packs this much information into so compact a format.