Periodic Table
Interactive periodic table of elements. Click any element to see its properties including atomic mass, category, and more.
Key features
- All 118 elements
- Color-coded categories
- Element detail cards
- Search functionality
Guide
The periodic table organizes all known chemical elements by atomic number, electron configuration, and recurring chemical properties. It is the most important reference tool in chemistry, providing at a glance the relationships between elements and predicting their behavior in chemical reactions. This guide covers the structure of the periodic table, what the data for each element means, periodic trends, and how to use element properties for practical purposes. Each element occupies one cell in the table, identified by its atomic number (the number of protons in the nucleus), chemical symbol (one or two letters, like H for hydrogen or Fe for iron), and atomic mass (the weighted average mass of naturally occurring isotopes, measured in atomic mass units). The atomic number uniquely defines an element. Carbon always has 6 protons. If an atom has 6 protons, it is carbon. The atomic mass includes contributions from all stable and long-lived isotopes weighted by their natural abundance. The table arranges elements in rows called periods and columns called groups (or families). Period 1 contains hydrogen and helium. Period 2 contains lithium through neon. Each period corresponds to the filling of a new electron shell. Group 1 contains the alkali metals (lithium, sodium, potassium, etc.). Group 18 contains the noble gases (helium, neon, argon, etc.). Elements in the same group share similar chemical properties because they have the same number of valence electrons (electrons in the outermost shell). Valence electrons determine how an element behaves in chemical reactions. Group 1 elements have 1 valence electron and tend to lose it, forming +1 ions. Group 2 elements have 2 valence electrons and form +2 ions. Group 17 elements (halogens) have 7 valence electrons and tend to gain 1 electron, forming -1 ions. Group 18 elements have a full outer shell (8 valence electrons, except helium with 2) and are chemically inert under normal conditions. This pattern explains why sodium (group 1) reacts vigorously with chlorine (group 17) to form sodium chloride. Element categories divide the periodic table into metals, nonmetals, and metalloids. Metals occupy the left and center of the table. They conduct electricity and heat, are malleable and ductile, have metallic luster, and tend to lose electrons in reactions. Nonmetals occupy the upper right. They are poor conductors, brittle as solids, and tend to gain electrons. Metalloids (boron, silicon, germanium, arsenic, antimony, tellurium, polonium) sit along the dividing line and have intermediate properties, making them useful as semiconductors. More specific categories further classify the elements. Alkali metals (group 1, excluding hydrogen) are soft, highly reactive metals that must be stored away from air and water. Alkaline earth metals (group 2) are reactive but less so than alkali metals. Transition metals (groups 3 through 12) include familiar metals like iron, copper, gold, and silver. They often form colored compounds and can have multiple oxidation states. Lanthanides and actinides (the two rows below the main table) include rare earth elements and radioactive elements like uranium and plutonium. Electronegativity measures an atom's ability to attract electrons in a chemical bond. It increases from left to right across a period and decreases from top to bottom within a group. Fluorine has the highest electronegativity (3.98 on the Pauling scale) and francium has the lowest (0.7). The difference in electronegativity between two bonded atoms determines the bond type: small difference (0 to 0.4) means nonpolar covalent, moderate difference (0.5 to 1.7) means polar covalent, and large difference (greater than 1.7) typically means ionic. Atomic radius generally increases from top to bottom within a group (as new electron shells are added) and decreases from left to right across a period (as increasing nuclear charge pulls electrons closer). Lithium has a larger atomic radius than carbon, even though both are in period 2, because carbon's 6 protons pull its electrons tighter than lithium's 3 protons pull its electrons. This trend affects how elements bond, their density, and their physical properties. Ionization energy is the energy required to remove an electron from a neutral atom in the gas phase. It increases from left to right across a period (atoms hold their electrons more tightly as the nucleus gains protons) and decreases from top to bottom within a group (outer electrons are farther from the nucleus and easier to remove). Noble gases have the highest ionization energies in each period because their full outer shells are very stable. Alkali metals have the lowest because removing one electron gives them a stable noble gas configuration. Electron affinity is the energy change when an atom gains an electron. Halogens have the most negative (most favorable) electron affinities because gaining one electron fills their outer shell. Noble gases have near-zero or positive electron affinities because they already have full outer shells and resist gaining electrons. This property, combined with ionization energy, explains why certain elements form cations (positive ions) and others form anions (negative ions). Oxidation states (oxidation numbers) indicate the charge an atom would have if all bonds were ionic. Some elements have fixed oxidation states: alkali metals are always +1, alkaline earth metals are always +2, fluorine is always -1, and oxygen is almost always -2 (except in peroxides where it is -1). Transition metals can have variable oxidation states. Iron can be +2 (ferrous) or +3 (ferric). Copper can be +1 (cuprous) or +2 (cupric). Knowing common oxidation states helps predict compound formulas and balance redox equations. Melting and boiling points show patterns across the periodic table. Metals generally have high melting points, with tungsten holding the record at 3,422 degrees Celsius. Noble gases have extremely low boiling points (helium boils at -269 degrees Celsius). Across a period, melting points tend to increase from group 1 to around the middle, then decrease. Down a group, trends depend on the category. Alkali metal melting points decrease going down the group, while carbon group elements show a more complex pattern. Density varies widely across the periodic table. Lithium, the lightest metal, has a density of 0.534 grams per cubic centimeter (it floats on water). Osmium, the densest naturally occurring element, has a density of 22.59 grams per cubic centimeter. Generally, density increases toward the middle and bottom of the table where elements have high atomic masses packed into relatively small atomic volumes. Isotopes are atoms of the same element with different numbers of neutrons. Carbon-12 has 6 neutrons, carbon-13 has 7, and carbon-14 has 8. The atomic mass on the periodic table is the weighted average across all naturally occurring isotopes. Carbon's atomic mass of 12.011 reflects that most carbon is carbon-12 with a small percentage of carbon-13 and trace carbon-14. Isotopes matter in radioactive dating, nuclear medicine, NMR spectroscopy, and nuclear energy. Electron configuration describes how electrons are distributed among an atom's orbitals. The periodic table structure directly reflects electron configurations. Period 1 fills the 1s orbital. Period 2 fills the 2s and 2p orbitals. Transition metals fill the d orbitals. Lanthanides and actinides fill the f orbitals. Knowing the electron configuration helps predict bonding behavior, magnetic properties, and spectral characteristics. The shorthand notation uses the previous noble gas in brackets: iron's configuration is [Ar] 3d6 4s2. Practical applications of periodic table knowledge span many fields. Materials science uses element properties to design alloys (mixing metals for desired strength, weight, and corrosion resistance). Pharmacology considers how different elements and their compounds interact with biological systems. Electronics relies on semiconductor properties of silicon and germanium. Environmental science tracks heavy metal contamination (lead, mercury, cadmium). Agriculture uses knowledge of essential plant nutrients (nitrogen, phosphorus, potassium, and micronutrients like iron, zinc, manganese). The elements discovered most recently are all synthetic and highly radioactive, existing for fractions of a second. Oganesson (element 118), the heaviest known element, was first synthesized in 2002. These superheavy elements are created by smashing lighter nuclei together in particle accelerators. While they have no practical applications, their properties test theoretical predictions about nuclear stability and the limits of the periodic table. The WebRecast interactive periodic table displays all 118 elements with their key properties. Click or tap any element to see its atomic number, symbol, name, atomic mass, electron configuration, electronegativity, ionization energy, density, melting and boiling points, and common oxidation states. Elements are color-coded by category (alkali metals, halogens, noble gases, transition metals, etc.) for quick visual identification. The table runs entirely in your browser, requires no downloads, and works on both desktop and mobile devices. Use it as a quick reference for homework, lab work, exam preparation, or professional applications in chemistry and related sciences.
Frequently asked questions
What info is shown?
Symbol, name, atomic number, atomic mass, and element category.
