Element Property Lookup

Two elements show similar chemistry mainly because they have the same number of valence electrons in their outermost shell. Since chemical reactions involve only valence electrons (forming, breaking and rearranging bonds), elements with the same valence-shell configuration tend to form similar types of compounds, similar oxidation states and similar bonding patterns. That is precisely why the periodic table groups these elements together. For example, lithium, sodium and potassium all have one valence electron and therefore behave alike.

Oxygen sits in Group 16 (chalcogens) of the periodic table. Its group-mates — sulfur (S), selenium (Se), tellurium (Te) and polonium (Po) — all have 6 valence electrons (ns2 np4) and therefore share oxygen's chemical character: they form divalent anions (O2-, S2-, Se2- …), make compounds like H2O, H2S, H2Se with hydrogen, and burn metals to give similar oxides/sulphides. Oxygen is the most electronegative of them and is uniquely small, which sometimes makes its properties slightly different (e.g., H-bonding in H2O).

Lithium is in Group 1 (alkali metals), so its closest cousins are sodium, potassium, rubidium and cesium — they all have a single ns1 valence electron and form +1 cations, react with water to give MOH + H2, and burn in air to form oxides/peroxides. However, lithium also resembles magnesium (Group 2) due to the diagonal relationship: similar size and charge density. Both Li and Mg burn in nitrogen to form nitrides (Li3N, Mg3N2) — a unique behaviour among other Group 1 metals.

Elements in the same group of the periodic table have the same number and type of valence electrons, simply at different principal energy levels. Since chemical properties depend almost entirely on valence electrons (it is they that take part in bond formation, ionisation and reactions), atoms with similar valence configurations react in similar ways. For example, Group 17 (halogens — F, Cl, Br, I) all have 7 valence electrons, so each forms one bond and a -1 ion. The pattern is so reliable that Mendeleev predicted properties of unknown elements simply from group position.

Beryllium belongs to Group 2 (alkaline-earth metals), so its group-mates magnesium, calcium, strontium and barium have the most similar properties: all form +2 ions, divalent oxides (MgO, CaO, etc.) and ionic chlorides. However, due to the diagonal relationship, beryllium also strongly resembles aluminium (Group 13). Both form covalent chlorides (BeCl2, AlCl3), amphoteric oxides (BeO, Al2O3) and form complex ions. So both Mg and Al are good answers, depending on context.

Lead has been used in bullets for centuries because of its high density (~11.3 g/cm3), low melting point (327 °C) and softness. The closest practical substitutes are tungsten (very dense, 19.3 g/cm3 — used in modern lead-free bullets), bismuth (almost the same density as lead, 9.8 g/cm3, non-toxic), and copper (used for full metal jacket bullets). For environmentally friendly ammunition, bismuth-tin and tungsten-polymer composites are now preferred over toxic lead. Of all of these, bismuth is chemically most similar to lead — both are heavy Group 14/15 post-transition metals.

That property is called lustre (also written “luster”). It is the way a polished surface reflects light, giving it a characteristic shine. All metals have lustre because their delocalised free electrons absorb and re-emit photons across the entire visible spectrum. Gold, silver and aluminium are particularly lustrous and that is why they are used for jewellery and mirrors. Non-metals (except iodine and graphite) generally lack lustre and appear dull.

Transition elements (the d-block, Groups 3–12) show many fascinating common properties: (1) high melting/boiling points and densities; (2) variable oxidation states (Fe2+/Fe3+, Cu+/Cu2+ etc.) due to similar energies of (n−1)d and ns; (3) coloured ions due to d-d electron transitions; (4) paramagnetism from unpaired d electrons; (5) excellent catalytic activity (Fe in Haber, Pt in catalytic converters); (6) ability to form complex ions ([Cu(NH3)4]2+); and (7) formation of alloys easily.

Potassium (K), which lies just below sodium in Group 1. Both have a single ns1 valence electron, so both readily lose this electron to form +1 cations (Na+, K+), react vigorously with water to form alkalis (NaOH, KOH) and hydrogen gas, and form halides with similar properties (NaCl, KCl). Lithium also belongs to the same group but lithium's small size makes it slightly different (more covalent character). For the closest match, K is the answer; rubidium and cesium are even more similar but less commonly studied at school level.

Sulfur is a yellow, brittle, non-metallic solid at room temperature. It has several allotropes — rhombic (most stable, density 2.07 g/cm3), monoclinic (above 96 °C) and plastic (formed when molten S is poured into cold water). Sulfur is insoluble in water but soluble in carbon disulfide. Its melting point is 115 °C and boiling point is 445 °C. It is a poor conductor of heat and electricity, has a faint smell when burning, and the molecule consists of S8 rings.

Neon (Ne, Z = 10) is a noble gas belonging to Group 18. Its outermost shell is completely filled (1s2 2s2 2p6), making it extremely stable and chemically inert — it forms no stable compounds at all under ordinary conditions. Neon is colourless, odourless, tasteless and monoatomic. It has very low melting (−248 °C) and boiling (−246 °C) points. When an electric discharge passes through neon at low pressure, it emits a brilliant orange-red glow — used in the famous neon advertising signs.

Ernest Rutherford studied the radiations emitted by radioactive elements and showed that they consist of three distinct types: alpha (α) particles — positively charged helium nuclei (2p + 2n), heavy and easily stopped by paper; beta (β) particles — fast electrons, negatively charged, stopped by aluminium foil; and gamma (γ) rays — high-energy electromagnetic waves, no charge, requiring thick lead to stop. He also discovered that radioactive decay is a random, exponential process and proposed the concept of half-life. His work earned him the 1908 Nobel Prize in Chemistry. Penetrating power: α < β < γ ; Ionising power: α > β > γ