Periodic Table — Interactive Element Reference
Explore the full periodic table with detailed element data including atomic number, mass, electron configuration, electronegativity, density, melting point, and category. Search by name, symbol, or atomic number.
C
Carbon
Z = 6 · Nonmetals
Solid at room tempAtomic Properties
Physical Properties
Discovery
What Is the Periodic Table — Interactive Element Reference?
This interactive reference tool gives you instant access to detailed data for all 118 confirmed chemical elements. Click any element in the grid or type a name, symbol, or atomic number into the search box to explore its properties.
- ›Full periodic table grid — all 118 elements in the standard 18-column layout, with lanthanides and actinides displayed below the main grid, color-coded by element category.
- ›Real-time search — find any element instantly by name (e.g. "carbon"), symbol (e.g. "C"), or atomic number (e.g. "6").
- ›Comprehensive element panel — click an element to see atomic mass, period, group, block, category, state, density, melting/boiling points, electron configuration, electronegativity, and discovery information.
- ›Category coloring — elements are visually grouped by category (alkali metals, noble gases, transition metals, etc.) for rapid visual orientation.
- ›Accurate IUPAC data — atomic masses and properties reflect current IUPAC recommendations and standard reference values.
Formula
Periodic Law
Properties of elements are a periodic function of their atomic number (Z)
Electron Configuration Notation
1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ ... (Aufbau principle)
Aufbau Orbital Filling Order
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p
→ 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Mass Number
A = Z + N (Z = protons, N = neutrons)
Electronegativity (Pauling Scale)
Range: 0.7 (Fr) to 4.0 (F) Increases: left→right, bottom→top
| Symbol | Name | Description |
|---|---|---|
| Z | Atomic Number | Number of protons in the nucleus; defines the element |
| A | Mass Number | Total protons + neutrons (for most stable isotope) |
| Symbol | Element Symbol | 1–2 letter abbreviation assigned by IUPAC |
| Period | Period | Row of the periodic table (1–7); equals the outermost electron shell |
| Group | Group | Column of the periodic table (1–18); determines valence electrons |
| Block | Block | s, p, d, or f — determined by the outermost orbital being filled |
| χ | Electronegativity | Ability to attract bonding electrons (Pauling scale) |
How to Use
- 1Search for an element: Type any part of an element name (e.g. "carbon"), its symbol (e.g. "C"), or its atomic number (e.g. "6") in the search box. Matching elements are highlighted instantly.
- 2Click an element cell: Click any element in the periodic table grid to open the detailed info panel below the grid. The selected element is highlighted with an orange border.
- 3Review element properties: The info panel shows all available data: atomic mass, period, group, block, category, physical state, density, melting point, boiling point, electron configuration, electronegativity, atomic radius, and discovery info.
- 4Browse by category: Use the color legend at the bottom of the grid to identify element categories. Click any element of the same color to compare properties within a group.
Example Calculation
Example: Carbon (C) — Element 6
Symbol: C Name: Carbon Atomic Number: 6
Atomic Mass: 12.011 u
Period: 2
Group: 14 (IVA)
Block: p
Category: Nonmetal
Electron config: [He] 2s² 2p²
Electronegativity: 2.55 (Pauling)
Covalent radius: 77 pm
Melting point: 3550°C (graphite sublimes)
Density: 2.267 g/cm³ (graphite)
Discovered: Ancient / Prehistory
Why carbon is unique
Carbon forms more compounds than any other element — over 10 million known organic compounds — due to its ability to form stable chains and rings with itself (catenation) and to form four covalent bonds. It is the basis of all known life on Earth and exists in multiple allotropes: diamond (hardest natural substance), graphite, graphene, fullerenes, and carbon nanotubes.
Understanding Periodic Table — Interactive Element Reference
How the Periodic Table Is Organized
The modern periodic table arranges all 118 known elements in order of increasing atomic number (number of protons). Elements are placed in 7 horizontal rows called periods and 18 vertical columns called groups. The period number equals the highest electron shell occupied by that element's electrons. Elements in the same group have the same number of valence electrons and thus similar chemical behaviour.
The table is also divided into four blocks (s, p, d, f) based on which type of orbital the outermost electrons occupy. The s-block (groups 1–2) and p-block (groups 13–18) form the main group elements. The d-block (groups 3–12) contains the transition metals. The f-block (lanthanides and actinides) is placed separately below the main table for space reasons.
Periodic Trends
- ›Atomic radius — increases down a group (more electron shells) and decreases left-to-right across a period (greater nuclear charge pulls electrons closer).
- ›Ionization energy — energy required to remove an electron. Increases left-to-right across a period and decreases down a group.
- ›Electronegativity — ability to attract bonding electrons. Highest at fluorine (4.0) in the top-right, lowest at francium (0.7) in the bottom-left.
- ›Electron affinity — energy released when an atom gains an electron. Generally increases across a period and decreases down a group, with notable exceptions.
- ›Metallic character — increases down a group and decreases left-to-right. Most metallic element: francium. Least metallic (most nonmetallic): fluorine.
Electron Configuration Rules
Three principles govern how electrons fill orbitals:
- ›Aufbau principle — electrons fill the lowest available energy orbital first. Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.
- ›Pauli exclusion principle — each orbital holds at most 2 electrons, which must have opposite spins (↑↓).
- ›Hund's rule — when multiple orbitals of equal energy are available, electrons occupy them singly before pairing up, and all single electrons have parallel spins.
Notable exceptions
Chromium (Cr, Z=24) and copper (Cu, Z=29) are classic exceptions to the Aufbau principle. Cr has configuration [Ar] 3d⁵ 4s¹ (not 3d⁴ 4s²) because a half-filled d subshell is extra stable. Cu has [Ar] 3d¹⁰ 4s¹ because a completely filled d subshell is extra stable. Similar exceptions occur in heavier transition metals.
Metals, Metalloids, and Nonmetals
Elements are broadly classified by their physical and chemical properties into three categories:
- ›Metals (left and centre of the table) — lustrous, malleable, ductile, good conductors of heat and electricity. Include alkali metals, alkaline earth metals, transition metals, and post-transition metals.
- ›Nonmetals (upper right of the table) — poor conductors, brittle in solid form, many are gases at room temperature. Include hydrogen, carbon, nitrogen, oxygen, halogens, and noble gases.
- ›Metalloids (semimetals) — elements with properties intermediate between metals and nonmetals. Include boron, silicon, germanium, arsenic, antimony, tellurium, and polonium. Silicon is the basis of semiconductor technology.
Element Categories and Examples
| Category | Count | Examples | Key Characteristic |
|---|---|---|---|
| Alkali metals | 6 | Li, Na, K, Rb, Cs, Fr | Highly reactive, 1 valence electron, soft |
| Alkaline earth | 6 | Be, Mg, Ca, Sr, Ba, Ra | 2 valence electrons, less reactive than alkali |
| Transition metals | 38 | Fe, Cu, Zn, Au, Ag, Pt | Variable oxidation states, coloured compounds |
| Post-transition | 7 | Al, Ga, In, Sn, Tl, Pb, Bi | Weaker metallic properties than transition metals |
| Metalloids | 7 | B, Si, Ge, As, Sb, Te, Po | Semiconductors; properties of both metals/nonmetals |
| Nonmetals | 7 | C, N, O, P, S, Se | Poor conductors, high electronegativity |
| Halogens | 5 | F, Cl, Br, I, At | Very reactive, 7 valence electrons |
| Noble gases | 6 | He, Ne, Ar, Kr, Xe, Rn | Stable, 8 valence electrons (He: 2), inert |
| Lanthanides | 15 | La–Lu | 4f orbital filling, rare earth elements |
| Actinides | 15 | Ac–Lr | 5f orbital filling, mostly radioactive |
History of the Periodic Table
The periodic table's modern form emerged from two centuries of discovery and refinement. Dmitri Mendeleev (1869) is credited with the first widely accepted arrangement, ordering known elements by atomic weight and leaving gaps for undiscovered elements — predictions that were confirmed when gallium, scandium, and germanium were later discovered. Henry Moseley's 1913 work established that atomic number (proton count), not atomic weight, is the true organizing principle.
Frequently Asked Questions
How many elements are in the periodic table?
There are 118 confirmed elements in the periodic table:
- ›Elements 1–94: occur in nature (some only in trace radioactive quantities)
- ›Elements 95–118: synthetic, produced in nuclear reactors or particle accelerators
- ›Elements 113, 115, 117, 118: officially named by IUPAC in 2016 (Nh, Mc, Ts, Og)
- ›Searches for elements 119 and 120 are ongoing in multiple research facilities
What do groups and periods mean?
- ›Group (column 1–18): same number of valence electrons → similar chemical behaviour
- ›Period (row 1–7): same number of occupied electron shells → similar atomic radii within a row
- ›Group 1 (alkali metals): 1 valence electron, highly reactive
- ›Group 18 (noble gases): 8 valence electrons (2 for He), very stable, inert
- ›Period 1 (H, He): only 1 electron shell; simplest elements
Why are elements arranged by atomic number and not atomic mass?
Mendeleev originally ordered by atomic mass, which worked well but had anomalies:
- ›Argon (Ar, mass 39.9) would follow potassium (K, mass 39.1) if ordered by mass, placing Ar in the wrong group
- ›Moseley's 1913 X-ray experiments showed atomic number is the true organizing principle
- ›Atomic number = protons = nuclear charge = the fundamental property that determines electron configuration and thus all chemical behaviour
- ›Isotopes of the same element have different atomic masses but the same atomic number — they belong in the same box
Why are noble gases so unreactive?
- ›Helium: 1s² — 2 electrons fill the n=1 shell completely
- ›Neon, Argon, etc.: 8 valence electrons fill the outermost s and p orbitals
- ›Very high ionization energies (most energy needed to remove an electron)
- ›Near-zero electron affinity (little tendency to gain electrons)
- ›Heavy noble gases (Kr, Xe) do form some compounds (e.g., XeF₂) under forcing conditions
What are transition metals and why are they special?
Transition metals are distinct because their d orbitals are partially filled:
- ›Multiple oxidation states: Fe²⁺ and Fe³⁺; Cu⁺ and Cu²⁺; Mn²⁺ through Mn⁷⁺
- ›Coloured ions and compounds: the d–d electron transitions absorb visible light
- ›Catalytic activity: Fe in Haber process, Pt in catalytic converters, Ni in hydrogenation
- ›Hard, high-melting-point metals with good electrical conductivity
- ›Form complex ions (coordination compounds) with ligands
What are lanthanides and actinides?
- ›Lanthanides (Z=57–71): 4f orbital fills; called "rare earth elements" (though most are not truly rare)
- ›Uses: neodymium magnets (Nd), europium phosphors (Eu), yttrium lasers (Y)
- ›Actinides (Z=89–103): 5f orbital fills; uranium (U) and thorium (Th) occur naturally
- ›Most actinides beyond uranium are synthetic (transuranium elements)
- ›Applications: uranium/plutonium in nuclear fission, americium in smoke detectors (Am-241)
How does electronegativity trend across the periodic table?
- ›Increases left-to-right across a period: nuclear charge increases, pulling bonding electrons closer
- ›Decreases top-to-bottom in a group: atomic radius increases, so the nucleus is farther from the bonding electrons
- ›Most electronegative: F (4.0), O (3.44), N (3.04), Cl (3.16)
- ›Least electronegative: Cs (0.79), Fr (0.7), Ba (0.89)
- ›Difference in electronegativity between bonded atoms predicts bond polarity and type (ionic vs. covalent)