How many elements are in the periodic table?

As of 2024, the periodic table contains 118 confirmed elements. Elements 1–94 occur naturally on Earth (though some only in trace amounts). Elements 95–118 are synthetic — created in nuclear reactors or particle accelerators. The last four elements to be officially named were Nihonium (113), Moscovium (115), Tennessine (117), and Oganesson (118), confirmed by IUPAC in 2016.

What are groups and periods on the periodic table?

Groups are the 18 vertical columns of the periodic table, numbered 1–18. Elements in the same group have the same number of valence electrons, giving them similar chemical properties. Periods are the 7 horizontal rows. Elements in the same period have the same number of electron shells. Moving across a period (left to right), atomic number increases, electronegativity generally increases, and atomic radius decreases. Moving down a group, atomic radius increases and ionisation energy decreases.

What is electron configuration and how do you read it?

Electron configuration describes how electrons are distributed among atomic orbitals. It is written as a sequence of subshell notations: 1s² 2s² 2p⁶ 3s¹ (sodium). The number prefix is the principal quantum number (shell), the letter is the subshell type (s, p, d, f), and the superscript is the number of electrons in that subshell. Noble gas notation abbreviates the filled core: [Ne] 3s¹ for sodium, where [Ne] represents the electron configuration of neon.

What is electronegativity?

Electronegativity is a measure of an atom's tendency to attract electrons in a chemical bond. The Pauling scale (0–4) is the most widely used: Fluorine is the most electronegative element (3.98), while Francium and Caesium are the least (≈0.70–0.79). The electronegativity difference between bonded atoms (Δχ) determines bond character: Δχ 1.7 = ionic.

What is the difference between metals, metalloids, and nonmetals?

Metals (the majority of elements) are shiny, malleable, ductile, and good conductors of heat and electricity. They tend to lose electrons and form positive ions. Nonmetals are typically brittle in solid form, poor conductors, and tend to gain electrons. Metalloids (B, Si, Ge, As, Sb, Te, At) have intermediate properties — they can conduct electricity under certain conditions (semiconductors) but not as efficiently as true metals. Silicon, the most important metalloid, is the foundation of modern electronics.

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Interactive Periodic Table of Elements

Q: Who invented the periodic table?

The periodic table was independently developed by Dmitri Mendeleev (Russian chemist) and Lothar Meyer (German chemist) in 1869. Mendeleev's version is considered more significant because he left gaps for undiscovered elements and correctly predicted their properties — a landmark demonstration of the scientific method. The element Mendelevium (Md, atomic number 101) is named in his honour.

Q: What are periodic trends?

Periodic trends are systematic patterns in element properties that repeat across periods and down groups. The main trends are: (1) Atomic radius — decreases across a period, increases down a group. (2) Ionisation energy — increases across a period, decreases down a group. (3) Electronegativity — increases across a period (highest: Fluorine at 3.98), decreases down a group (lowest: Francium at 0.70). (4) Electron affinity — generally increases across a period. These trends arise from changes in effective nuclear charge and electron shielding.

Q: What are the noble gases and why are they unreactive?

The noble gases (Group 18: He, Ne, Ar, Kr, Xe, Rn, Og) are chemically inert under normal conditions because they have completely filled valence electron shells — helium has 2 valence electrons (full 1s shell), while neon, argon, krypton, xenon, and radon all have 8 valence electrons (full ns² np⁶ configuration). This stable octet configuration means they have no tendency to gain, lose, or share electrons with other atoms.

Explore all 118 chemical elements. Click any element to see its atomic number, atomic mass, electron configuration, electronegativity, and physical state. Filter by element category or search by name or symbol.

Built by He Loves Math — the most detailed free periodic table for chemistry students, teachers, and curious minds.

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Quick Reference — Key Quantities

The periodic table encodes three fundamental quantities per element:

$$Z = \text{atomic number (proton count)}, \quad M = \text{atomic mass (u)}, \quad e^- = \text{electrons}$$

For a neutral atom: electrons = protons = Z. Atomic mass ≈ protons + neutrons for stable isotopes. The periodic trends follow from Z and electron shell filling via the Aufbau principle.

Interactive Periodic Table — All 118 Elements

Click any element tile for a detailed property panel below the table. Use the search box or category filter buttons to highlight specific elements.

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Element Details

What Is the Periodic Table of Elements?

The periodic table of elements is the most fundamental organisational tool in all of chemistry. It is a systematic arrangement of all known chemical elements, ordered by increasing atomic number (the number of protons in an atom's nucleus), in a grid that groups elements with similar chemical properties into vertical columns called groups and organises elements with the same number of electron shells into horizontal rows called periods.

Every element on the table is a pure substance made of atoms with a unique number of protons. Hydrogen (H), with 1 proton, is element 1. Helium (He) has 2. Carbon (C) has 6. Gold (Au) has 79. Oganesson (Og), the heaviest confirmed element, has 118. The number of protons — the atomic number Z — is what defines an element: change the number of protons and you have a completely different element.

The periodic table was not designed arbitrarily. Its layout is a direct consequence of quantum mechanics: specifically, the way electrons fill atomic orbitals according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle. Elements in the same group share the same valence electron configuration, which is why they behave similarly in chemical reactions. Group 1 elements (lithium, sodium, potassium etc.) all have one valence electron and react vigorously with water. Group 18 elements (noble gases) all have full valence shells and are chemically inert under normal conditions.

How to Read the Periodic Table

Each element tile in the periodic table displays standardised information. In most standard formats, a tile shows:

Atomic number (Z) — top centre or top left. An integer from 1 to 118. Equals the number of protons (and in a neutral atom, the number of electrons).
Chemical symbol — the largest, most prominent text. One or two letters (first always capitalised). Derived from English or Latin names (Fe = Ferrum for Iron, Au = Aurum for Gold).
Element name — full English name below the symbol.
Atomic mass (M) — the standard atomic weight in unified atomic mass units (u or Da). For natural elements, this is the weighted average of all stable isotopes. For synthetic or radioactive elements, the number in parentheses indicates the mass number of the most stable known isotope.

Atomic Mass Relationship $$M \approx Z \cdot m_p + (A - Z) \cdot m_n \qquad \text{where } A = \text{mass number, } m_p \approx m_n \approx 1\text{ u}$$

Groups (Columns)

The 18 vertical columns are numbered 1–18 from left to right. Elements in the same group have the same number of valence electrons, producing similar chemical behaviour across a wide range of conditions. The most chemically important group families are:

Group 1 — Alkali Metals: Li, Na, K, Rb, Cs, Fr. All have 1 valence electron. Extremely reactive with water, producing hydrogen gas and a metal hydroxide.
Group 2 — Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra. Two valence electrons. Less reactive than Group 1 but still reactive.
Groups 3–12 — Transition Metals: The d-block. Variable oxidation states, coloured compounds, catalytic activity.
Group 17 — Halogens: F, Cl, Br, I, At, Ts. Seven valence electrons — one short of a full shell. Highly electronegative and reactive.
Group 18 — Noble Gases: He, Ne, Ar, Kr, Xe, Rn, Og. Complete valence shells — chemically inert under standard conditions.

Periods (Rows)

Period 1 contains only 2 elements (H, He — filling the 1s orbital). Period 2 has 8 elements (filling 2s and 2p orbitals). Period 3 has 8 elements (3s and 3p). Periods 4 and 5 each have 18 elements (including d-block transition metals). Periods 6 and 7 have 32 elements each (including both d-block and f-block lanthanides/actinides).

The 10 Element Categories

🔴 Alkali Metals (Group 1)

Six highly reactive metals: Li, Na, K, Rb, Cs, Fr. Soft, low-density, low melting point. React explosively with water releasing H₂ gas. Always have +1 oxidation state in compounds. Reactivity increases down the group.

🟠 Alkaline Earth Metals (Group 2)

Be, Mg, Ca, Sr, Ba, Ra. Harder and denser than alkali metals. Two valence electrons — form +2 ions. Ca is crucial for bones; Mg in chlorophyll. Reactivity increases down the group.

🟡 Lanthanides (f-block)

Elements 57–71 (La to Lu). The first row of the f-block. Similar chemical properties (3+ oxidation state dominant). Used in magnets, phosphors, lasers, and catalysts. Sometimes called "rare earth metals" (though many are not rare).

🟠 Actinides (f-block)

Elements 89–103 (Ac to Lr). All are radioactive. Include Uranium (U) and Plutonium (Pu) — key to nuclear energy and weapons. Thorium has potential as a nuclear fuel. Elements beyond 94 (Pu) are entirely synthetic.

🔵 Transition Metals (d-block, Groups 3–12)

38 elements including Fe, Cu, Zn, Ag, Au, Pt, Ti. Characterised by partially filled d-orbitals. Multiple oxidation states, form coloured complexes, excellent electrical conductors, catalysts. Form the bulk of engineered metal structures.

🔷 Post-Transition Metals

Al, Ga, In, Sn, Tl, Pb, Bi, Fl. Softer and with lower melting points than transition metals. Show more covalent bond character. Tin and lead have been used since antiquity; aluminium is the most abundant metal in Earth's crust.

🟢 Metalloids

B, Si, Ge, As, Sb, Te, At. Intermediate properties between metals and nonmetals. Silicon and germanium are essential semiconductors forming the basis of all modern electronics and solar cells.

🔵 Diatomic Nonmetals

H, N, O, F, Cl, Br, I. Exist as diatomic molecules in pure form (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂). Essential to life: O₂ for respiration, N₂ (78% of atmosphere), H₂ for energy.

🟢 Polyatomic Nonmetals

C, P, S, Se. Form complex structures: carbon forms diamond, graphite, graphene, and fullerenes. Sulphur forms S₈ rings. All essential for biological molecules.

🟣 Noble Gases (Group 18)

He, Ne, Ar, Kr, Xe, Rn, Og. Complete valence shells — chemically inert under most conditions. He used in MRI cooling and balloons. Ar in light bulbs. Xe in xenon arc lamps and anaesthesia. Rn is a radioactive health hazard in buildings.

Periodic Trends — Mathematical Patterns Across the Table

The brilliance of the periodic table is that element properties do not vary randomly — they follow predictable trends governed by quantum mechanics. Understanding these trends allows chemists to predict the behaviour of any element, even without experimental data.

1. Atomic Radius

Atomic radius decreases across a period (left to right) because increasing nuclear charge pulls electrons closer. Atomic radius increases down a group because each new period adds an additional electron shell.

Bohr Model Orbital Radius $$r_n = \frac{n^2 a_0}{Z_{\text{eff}}} \qquad a_0 = 0.529\text{ Å (Bohr radius)}$$

Where $n$ is the principal quantum number and $Z_\text{eff}$ is the effective nuclear charge. As $Z_\text{eff}$ increases across a period, r_n decreases at a given n, explaining why the atomic radius shrinks.

2. Ionisation Energy

The first ionisation energy (IE₁) is the energy required to remove one electron from a gaseous atom. It increases across a period (harder to remove electrons from increasingly attractive nucleus) and decreases down a group (outer electrons are farther from nucleus and more shielded).

Bohr Model Energy Levels $$E_n = -\frac{Z^2 \cdot 13.6\text{ eV}}{n^2} \qquad \text{IE}_1 = -E_1 = \frac{Z_{\text{eff}}^2 \cdot 13.6\text{ eV}}{n^2}$$

The hydrogen atom (Z=1, n=1) has the ionisation energy calculated as $E = 13.6\text{ eV}$. The Bohr formula is exact for hydrogen-like (one-electron) atoms and approximate for multi-electron systems after applying effective nuclear charge corrections (Slater's rules).

3. Electronegativity

Electronegativity (χ) is an atom's tendency to attract bonding electrons. The Pauling scale sets fluorine as 3.98 (highest) and assigns other values relative to it. Noble gases are typically omitted (undefined) since they rarely form bonds.

Electronegativity and Bond Character $$\Delta\chi = |\chi_A - \chi_B|$$ $$\Delta\chi < 0.5 \implies \text{nonpolar covalent}$$ $$0.5 \leq \Delta\chi \leq 1.7 \implies \text{polar covalent}$$ $$\Delta\chi > 1.7 \implies \text{ionic}$$

4. Electron Affinity

Electron affinity is the energy released (or absorbed) when an electron is added to a neutral atom in the gas phase. Generally increases across a period and varies irregularly down a group. Chlorine (Cl) has the highest electron affinity among stable elements.

Electron Configuration and the Periodic Table

The structure of the periodic table directly reflects how electrons fill atomic orbitals. The Aufbau ("building up") principle states that electrons fill the lowest-energy available orbital first.

Orbital Filling Order (Madelung Rule) $$1s \to 2s \to 2p \to 3s \to 3p \to 4s \to 3d \to 4p \to 5s \to 4d \to 5p \to 6s \to 4f \to 5d \to 6p \to 7s \to 5f \to 6d \to 7p$$

This filling order explains the periodic table's structure:

s-block (Groups 1–2): valence electrons in s-orbitals → periods 1–7, 2 columns
p-block (Groups 13–18): valence electrons in p-orbitals → periods 2–7, 6 columns
d-block (Groups 3–12): filling d-orbitals → transition metals, 10 columns
f-block (below main table): filling f-orbitals → lanthanides & actinides, 14 columns

Two important exceptions to the Aufbau rule occur because of the extra stability of half-filled and fully-filled d-orbitals:

Chromium (Cr, Z=24): Expected [Ar] 3d⁴ 4s²; actual [Ar] 3d⁵ 4s¹ (half-filled d shell)
Copper (Cu, Z=29): Expected [Ar] 3d⁹ 4s²; actual [Ar] 3d¹⁰ 4s¹ (fully filled d shell)

Key Chemistry Formulas Rooted in the Periodic Table

Molar Mass and Amount of Substance

$$n = \frac{m}{M} \qquad N = n \times N_A \qquad N_A = 6.022 \times 10^{23}\text{ mol}^{-1}$$

Where $n$ = amount of substance (mol), $m$ = mass (g), $M$ = molar mass (g/mol, numerically equal to atomic mass in u), $N$ = number of particles, and $N_A$ = Avogadro's constant. The molar mass of any compound is the sum of the atomic masses of its constituent elements (read directly from the periodic table).

Rydberg Formula (Spectral Lines)

Rydberg Formula for Hydrogen $$\frac{1}{\lambda} = R_H \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right) \qquad R_H = 1.097 \times 10^7\text{ m}^{-1}$$

This formula predicts the wavelengths of light emitted or absorbed when a hydrogen electron transitions between energy levels $n_1$ and $n_2$. Each element has a unique emission spectrum — the foundation of spectroscopy — which is how Mendeleev's predicted elements were confirmed by their spectra before they were physically isolated.

de Broglie Wavelength

$$\lambda = \frac{h}{mv} = \frac{h}{p} \qquad h = 6.626 \times 10^{-34}\text{ J·s}$$

The wave nature of electrons — expressed through the de Broglie relation — underpins the quantum mechanical model that the modern periodic table is built upon. Electrons do not orbit in fixed paths; they occupy probability clouds (orbitals) defined by wave functions (ψ), and the periodic table's structure is a direct consequence of how these wave functions are filled.

History of the Periodic Table

The search for an organising principle of the elements spanned centuries. By the 1860s, chemists had identified about 60 elements and observed patterns in their properties. Several scientists made contributions — John Newlands noticed that elements showed repeating properties every 8th element (his "law of octaves"), and Lothar Meyer produced a periodic table based on atomic volume in 1868.

However, it was Dmitri Ivanovich Mendeleev (1834–1907) who presented the most complete and predictively powerful version on March 6, 1869, to the Russian Chemical Society. Crucially, Mendeleev left deliberate blank spaces for undiscovered elements and predicted their properties with remarkable accuracy. His predicted "eka-aluminium" was discovered as Gallium (Ga) in 1875 by Paul Émile Lecoq de Boisbaudran, with properties matching Mendeleev's prediction within a fraction of a percent. "Eka-silicon" became Germanium (Ge) in 1886. These discoveries cemented the periodic table as one of science's greatest theoretical frameworks.

The quantum mechanical explanation of the periodic table came in the 20th century with Niels Bohr's atomic model (1913), the development of quantum mechanics by Heisenberg, Schrödinger, and Dirac, and the formulation of the Aufbau principle. Henry Moseley's 1913 work demonstrated that the true organising property was atomic number (proton count), not atomic mass — resolving several anomalies in Mendeleev's original arrangement.

The most recent additions are the four Period 7 elements confirmed by IUPAC in 2016: Nihonium (Nh, 113), Moscovium (Mc, 115), Tennessine (Ts, 117), and Oganesson (Og, 118). These elements were synthesised only for fractions of a second in particle accelerators.

Block	Groups	Orbital Type	Periods	Max Electrons
s-block	1–2	s (0–2 electrons)	1–7	2 per period
p-block	13–18	p (0–6 electrons)	2–7	6 per period
d-block	3–12	d (0–10 electrons)	4–7	10 per period
f-block	—	f (0–14 electrons)	6–7	14 per period

Frequently Asked Questions

How many elements are on the periodic table?

There are 118 confirmed elements on the current periodic table. Elements 1–94 are found naturally (though some only in trace amounts). Elements 95–118 are synthetic. The four most recently named elements (2016) are Nihonium (113), Moscovium (115), Tennessine (117), and Oganesson (118).

Who invented the periodic table?

Dmitri Mendeleev is credited with the most significant early periodic table (1869), which he arranged by atomic mass and left deliberate gaps predicting undiscovered elements. Lothar Meyer independently developed a similar system at the same time. The modern table is ordered by atomic number, a refinement due to Henry Moseley (1913).

What are periods and groups on the periodic table?

Periods (rows) number 1–7. Elements in the same period have the same highest principal quantum number (number of electron shells). Groups (columns) number 1–18. Elements in the same group have the same valence electron configuration, giving them similar chemical properties. For example, all Group 1 elements have one valence electron and form +1 ions.

What are periodic trends?

Periodic trends are patterns in properties that repeat across periods and down groups: (1) Atomic radius decreases across a period, increases down a group. (2) Ionisation energy increases across a period, decreases down a group. (3) Electronegativity increases across a period (highest: F = 3.98), decreases down a group (lowest: Fr ≈ 0.70). These trends arise from changing effective nuclear charge and electron shell count.

What is electron configuration?

Electron configuration describes how electrons are distributed in an atom's orbitals. Written as a sequence of subshell notations: e.g. sodium (Na) = 1s² 2s² 2p⁶ 3s¹, or abbreviated with noble gas core: [Ne] 3s¹. The superscript is the number of electrons in that subshell. Click any element in the table above to see its electron configuration.

Why are lanthanides and actinides shown separately?

Lanthanides (La–Lu, elements 57–71) and actinides (Ac–Lr, elements 89–103) are members of the f-block. Placing them within the main table would make it 32 columns wide, impractical for most use cases. They are shown separately below the main table with placeholder tiles at their logical positions in Groups 3 of Periods 6 and 7.

What does the atomic mass in parentheses mean?

Atomic mass values in parentheses — e.g. (98) for Technetium, (244) for Plutonium — indicate that the element has no stable isotopes. The number represents the mass number of the most stable known isotope (longest half-life). All elements from atomic number 84 (Polonium) onwards are radioactive, as are Technetium (43) and Promethium (61).

What are the noble gases and why are they unreactive?

Noble gases (Group 18: He, Ne, Ar, Kr, Xe, Rn, Og) have completely filled valence electron shells — helium has 2 (full 1s²), all others have 8 (ns² np⁶). This configuration is maximally stable: no tendency to gain, lose, or share electrons. Xenon can form compounds with highly electronegative elements (XeF₂, XeF₄), but this requires extreme conditions.