ChemistryClassification of elements and periodicityClass 11

Periodic table & periodic trends: interactive simulation

The modern periodic table organizes chemical elements by atomic number. Repeating electronic configurations give rise to predictable, fundamental periodic properties — atomic radius, ionization energy, electron affinity and chemical reactivities.

Useful for: JEE Main, JEE Advanced, NEET, CBSE Boards · 3 modes in one lab

What you can explore

  1. 1

    Periodic Trends Heatmap

    Color-coded periodic table heatmap for radius, ionization enthalpy and electronegativity.
  2. 2

    Z_eff & Slater’s Rules

    Core shielding shells and effective nuclear pull calculator.
  3. 3

    Acidic & Basic Oxides

    Basic to amphoteric to acidic transition across Period 3 oxides.

Try this

Explore periodic trends across the s, p and d blocks with an interactive color heatmap, calculate effective nuclear charge Z_eff via Slater’s rules, uncover why Nitrogen has higher IE than Oxygen, and trace oxide acidities from basic Na₂O to acidic Cl₂O₇.

  1. 1

    Select Ionization Enthalpy and compare Nitrogen (Z=7) with Oxygen (Z=8).

    What you should see

    Nitrogen (1402 kJ/mol) has a HIGHER 1st IE than Oxygen (1314 kJ/mol) because half-filled 2p³ is extra stable!

  2. 2

    Compare Beryllium (Z=4) with Boron (Z=5).

    What you should see

    Beryllium (899 kJ/mol) > Boron (801 kJ/mol) because removing an electron from 2s² takes more energy than from 2p¹.

  3. 3

    In Period 3 Oxides, inspect Al₂O₃.

    What you should see

    Al₂O₃ is amphoteric — it reacts with both strong acids (forming Al³⁺) and strong bases (forming aluminate [Al(OH)₄]⁻).

Formulas behind it

Effective nuclear charge (Slater’s rule)

Zeff=Z−σ(σ=screening / shielding constant)Z_{\text{eff}} = Z - \sigma \quad (\sigma = \text{screening / shielding constant})

Ionization enthalpy periodic trend

IE1∝Zeffr(Anomalies: N>O,  Be>B)\text{IE}_1 \propto \frac{Z_{\text{eff}}}{r} \quad (\text{Anomalies: } \text{N} > \text{O}, \; \text{Be} > \text{B})

Electronegativity relation (Pauling scale)

XA−XB=0.102Δ(Δ=EA−B−EA−AEB−B)X_A - X_B = 0.102 \sqrt{\Delta} \quad (\Delta = E_{A-B} - \sqrt{E_{A-A} E_{B-B}})

Common mistakes

  • Assuming Chlorine has lower electron gain enthalpy than Fluorine — Chlorine has the HIGHEST negative electron gain enthalpy in the entire periodic table because 2p orbital in Fluorine suffers high interelectronic repulsion.
  • Thinking atomic radius increases across a period — radius DECREASES across a period because nuclear charge Z rises faster than shielding.
  • Overlooking the Lanthanoid Contraction — poor shielding by 4f electrons causes 4d (Zr) and 5d (Hf) elements to have virtually identical covalent radii (~160 pm).
Stuck on a periodic table & periodic trends question?Snap it and get a whiteboard video solution.

FAQ

Why is the second ionization enthalpy of Sodium so enormous?

Because removing the 2nd electron breaks into the stable, noble-gas closed-shell configuration of Neon (2s² 2p⁶).

What is Slater’s screening constant?

A quantitative measure of how much inner electrons shield the valence electrons from the full positive pull of the nucleus.

Why does electronegativity decrease down a group?

Because the atomic radius increases and additional electron shells place the valence bonding electrons further from the nucleus.

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