Interactive Periodic Table
Click any element for its atomic weight, electron configuration, electron shell diagram, oxidation states, density, and melting point. Search by name, symbol, or number, and color the table by category, block, or state.
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What it does
The Interactive Periodic Table displays all 118 elements with their key properties. Click any element to see its atomic weight, electron configuration, oxidation states, density, and melting point. It serves students learning the table and researchers who need quick property lookup.
How it works
Elements are ordered by atomic number, which equals the number of protons in the nucleus, and arranged into periods (rows) and groups (columns) so that elements in the same group share similar valence-electron patterns and chemical behavior. The table is color-coded by category such as alkali metals, halogens, and noble gases. Lanthanides and actinides are placed in a separate block below the main body.
Worked example
Click carbon (atomic number 6): atomic weight 12.011, electron configuration 1s2 2s2 2p2, common oxidation states -4, +2, +4, density about 2.27 g/cm3, melting point 3550 °C. Chlorine (17) shows oxidation states -1, +1, +3, +5, +7 and is a halogen.
When to use it
Use it to look up an atomic weight before a molar mass calculation. Use it to check the typical oxidation states of an element when writing a formula. Use it to compare categories and periodicity across a row or group.
Periodic trends
The periodic table is arranged so that properties recur in a predictable way, and those trends are what let you answer chemistry questions without memorising 118 entries. Electronegativity and ionisation energy rise left to right across a period and fall down a group; atomic radius does the opposite, shrinking across a period and growing down a group. The reason is that moving right adds protons without adding a shell, pulling electrons closer, while moving down adds a shell that puts the outer electrons further from the nucleus. Two comparisons come up constantly. Fluorine is the most electronegative element at 3.98, above oxygen's 3.44, because it sits one group further right in the same period. And the radius jump from period 2 to period 3 is much larger than from period 3 to period 4, because although a shell is added each time, the intervening d-block and f-block electrons screen poorly and the rising nuclear charge pulls the outer shell in, producing the d-block and lanthanide contractions. Once you have two electronegativity values, their difference classifies the bond: under about 0.4 nonpolar covalent, 0.4 to 1.7 polar covalent, and above about 1.7 ionic. Open an element's panel to read its electronegativity, then compare two elements to judge a bond.
FAQ
- How is the periodic table organized?
- Elements are listed in order of atomic number and placed in rows called periods and columns called groups, so elements with similar chemistry sit in the same column.
- What is the difference between atomic number and atomic weight?
- Atomic number is the count of protons and defines the element, while atomic weight is the average mass of its atoms including isotopes, shown in g/mol for molar calculations.
- Why are lanthanides and actinides shown separately?
- They are the f-block elements with 14 members each; placing them under the main table keeps the layout compact while still preserving group and period order.
- What do the colors in the periodic table mean?
- Colors mark element categories such as alkali metals, alkaline earth metals, transition metals, metalloids, nonmetals, halogens, and noble gases, making trends easy to see.
- Why is fluorine more electronegative than oxygen?
- Electronegativity rises from left to right across a period because the nuclear charge increases while electrons are added to the same shell, so the pull on the bonding electrons grows. Oxygen (group 16) and fluorine (group 17) sit in the same period, but fluorine's larger effective nuclear charge and smaller radius mean it attracts a shared pair of electrons more strongly, which is why fluorine's Pauling value is 3.98 against oxygen's 3.44. Fluorine is the most electronegative element, so it is the only one that beats oxygen in this contest — but note the effect is not uniform down a group: electronegativity falls as you go down a column because the outer shell is further from the nucleus and better shielded.
- Why is the atomic-radius jump from period 2 to period 3 much larger than from period 3 to period 4?
- Atomic radius grows within a group mainly because each new period adds a whole shell, so the outer electrons sit further from the nucleus. Going from period 2 (e.g. lithium, shell n = 2) to period 3 (sodium, n = 3) adds an entire principal shell of about 100 pm, so the jump is large. From period 3 to period 4 (potassium and calcium, n = 4) the shell is also added, but the eighteen 3d electrons that intervene screen the outer 4s electrons poorly; the increased nuclear charge pulls the shell in, so the increase is much smaller and radii even dip slightly across the first transition series. This 'd-block contraction' — and the stronger lanthanide contraction — is why heavier elements are not as large as a simple shell-count model would predict.
- How do I use electronegativity difference to judge bond polarity or bond type?
- Subtract the two Pauling values and read the difference. Below about 0.4 the bond is essentially nonpolar covalent (H–H, C–H); roughly 0.4 to 1.7 is polar covalent, with the more electronegative atom carrying the partial negative charge δ− (H–Cl at 0.9, O–H at 1.4); above about 1.7 the bond is usually ionic, because one atom pulls the electron pair so completely that discrete ions form (NaCl at 2.1, CsF at 3.2). The cutoffs are boundaries on a continuum rather than hard rules, and small differences matter biologically: a C–H bond at 0.4 is nonpolar while an N–H bond at 0.9 is strongly polar, which is why amide and hydroxyl groups form hydrogen bonds and alkyl chains do not.