Types of Elements — Metals, Non-metals, Metalloids & Noble Gases

Four Kinds of Elements, One Periodic Table

So far in this chapter we have classified elements by their position — by period, group, and block. Now we change the lens and classify them by the kind of chemistry they do: whether they behave like metals, non-metals, metalloids or noble gases. This is the classification you will use most often in everyday reasoning — in reactions, in Lewis structures, in predicting acidic/basic oxides, and in every problem that asks "which of these is more electropositive?"

The four families at a glance

• Metals (about 78 % of the 118 known elements) — shiny, malleable, good conductors, low ionisation enthalpy, form positive ions.
• Non-metals (only about 17 elements, mostly in the top-right corner plus hydrogen) — typically gases or brittle solids, poor conductors, high ionisation enthalpy, form negative ions or share electrons covalently.
• Metalloids (just 7 elements, sitting exactly on the staircase line) — mixed properties, most important as semi-conductors (Si, Ge).
• Noble gases (Group 18, 7 elements) — a special class of non-metals with completely-filled valence shells; nearly chemically inert.

Notice the asymmetry: metals vastly outnumber non-metals, yet the non-metals drive most of the interesting molecular chemistry (water, DNA, every organic compound, every acid).

Where the four families sit on the table

The visual rule is simple: a staircase line cuts through the p-block. Everything to the left and below the line is a metal; everything to the right and above the line is a non-metal; the elements sitting right on the line are metalloids; and the rightmost column (Group 18) is the noble gases.

[Board Level] Memorise the 7 metalloids — B, Si, Ge, As, Sb, Te, Po — because the staircase is rarely drawn in exam booklets. If you can name these seven, you can identify the staircase. (Some textbooks add At, making 8; NCERT's standard list is the 7 given here.)

Why this classification works

Each of the four families has a single electronic signature that predicts its chemistry:

• Metals: low ionisation enthalpy, positive ions, large atoms at a given period.
• Non-metals: high ionisation enthalpy, high electron affinity, small atoms at a given period.
• Metalloids: intermediate values of both — hence the semi-conductor behaviour that turned Si into the basis of the entire electronics industry.
• Noble gases: closed-shell valence configuration — hence the near-zero reactivity.

[JEE Tip] Whenever a problem gives you an element by atomic number or configuration and asks for its "type", the shortcut is: (1) find its position, (2) check its location relative to the staircase. No need to compute the full electronic configuration if you can read the table mentally.

Metals — The Dominant Family

Metals make up three out of every four elements on the periodic table. There is a good electronic reason for that: across the majority of the table, atoms have only a few valence electrons loosely held in diffuse outer shells, and the energetically favourable move is to lose them. That is the essence of metallic character.

Physical properties you already know (but now understand why)

• Lustre. Free conduction-band electrons absorb and re-emit visible light across a broad frequency range, producing the characteristic metallic sheen.
• Malleability and ductility. Metallic bonding is non-directional (unlike covalent bonds), so layers of atoms can slide past each other without breaking the bond network.
• High electrical and thermal conductivity. The "sea of electrons" model: delocalised valence electrons respond instantly to an applied field or thermal gradient, carrying charge and energy.
• High melting and boiling points. Strong metallic bonding (except in the alkali metals, which are surprisingly soft and low-melting because each atom contributes only one valence electron).

Chemical behaviour

• Form positive ions by losing valence electrons. Na → Na⁺, Ca → Ca²⁺, Al → Al³⁺. The charge is usually the group number (for Groups 1, 2, 13) or the group number minus 10 (d-block Groups 3–12, with variable oxidation states).
• Reducing agents. Because they give up electrons readily, metals reduce other species. Group 1 and Group 2 metals are powerful reducing agents; iron and zinc are moderate ones.
• Basic oxides. $\text{Na}_{2}\text{O}, \text{CaO}, \text{MgO}$ all dissolve in water to produce alkaline solutions. As we move across a period, the oxides become progressively less basic (MgO is strongly basic, $\text{Al}_{2}\text{O}_{3}$ is amphoteric, and by the time we reach $\text{SO}_{3}$ the oxide is strongly acidic).
• React with acids to liberate hydrogen. $\text{Zn} + 2\,\text{HCl} \to \text{ZnCl}_{2} + \text{H}_{2}\uparrow$.

Where they live on the table

Metals fill the entire left half of the table and much of the p-block:

• All of Groups 1 and 2 (s-block metals) — alkali and alkaline-earth metals.
• All of Groups 3–12 (d-block) — the 40 transition metals.
• All 28 elements of the f-block — lanthanoids and actinoids.
• Most of Groups 13–16 on the lower-left side of the staircase — e.g. Al, Ga, Sn, Pb, Bi.

One memorable exception

Mercury (Hg) is the only metal that is liquid at room temperature (melting point −38.8 °C). The reason is genuinely quantum-mechanical: relativistic contraction of the $6s$ orbital in Hg leaves its filled $5d^{10}\,6s^{2}$ "pseudo-noble-gas" configuration unusually stable, so metallic bonds between Hg atoms are weak. Caesium and gallium come close (mp 28.5 °C and 29.8 °C respectively — they melt in your palm).

[JEE Tip] The liquid-at-room-temperature trap is a favourite of both JEE and NEET. Remember: Hg is the only liquid metal, while Br is the only liquid non-metal (besides water-based systems). No other elemental liquids exist at STP.

Non-metals — A Small but Chemically Rich Minority

Only about 17 elements are non-metals — but they include every gas we breathe, every atom that makes up water, every building block of organic molecules, and every oxidising agent in inorganic chemistry. Non-metals punch far above their weight.

Who's in, who's out?

The non-metals are: H, C, N, O, P, S, Se, F, Cl, Br, I, At (and some argue Ts), plus the Group-18 noble gases (He, Ne, Ar, Kr, Xe, Rn, Og), which we treat as a separate sub-family.

They all sit in the top-right corner of the periodic table, above the staircase line (except hydrogen, which is a non-metal perched at the top of Group 1 for electronic-configuration reasons).

Physical properties

• State at 25 °C: a mixture. The halogens alone cover gas ($\text{F}_{2}, \text{Cl}_{2}$), liquid ($\text{Br}_{2}$) and solid ($\text{I}_{2}$). Solid non-metals are brittle.
• Dull or coloured (S is yellow, $\text{Br}_{2}$ reddish-brown, $\text{I}_{2}$ violet), not shiny.
• Poor conductors of heat and electricity. Two famous exceptions — graphite (electrically conducting because of its delocalised π-electron system) and diamond (high thermal conductivity despite being an electrical insulator).
• Low density, low melting point in general — a sulfur or phosphorus crystal crumbles when you strike it.

Chemical behaviour

• Form negative ions by gaining electrons. Cl → Cl⁻, O → O²⁻, N → N³⁻. The charge equals (group number − 18) for Groups 14–17.
• Oxidising agents. Because they pull electrons, they oxidise other species. F₂ is the strongest common oxidiser; O₂ and Cl₂ are everyday oxidisers.
• Acidic (or neutral) oxides. $\text{CO}_{2} + \text{H}_{2}\text{O} \to \text{H}_{2}\text{CO}_{3}$, $\text{SO}_{3} + \text{H}_{2}\text{O} \to \text{H}_{2}\text{SO}_{4}$, $\text{P}_{4}\text{O}_{10} + 6\,\text{H}_{2}\text{O} \to 4\,\text{H}_{3}\text{PO}_{4}$.
• Form covalent compounds with each other (e.g. $\text{CO}_{2}, \text{H}_{2}\text{O}, \text{SO}_{2}$), and ionic compounds with metals (e.g. $\text{NaCl}, \text{MgO}, \text{AlCl}_{3}$).

Why non-metals look so different from each other

The answer is electronic-shell size. A small atom like F or O holds its valence electrons tightly, so gaining electrons is strongly exothermic — producing intensely oxidising behaviour. A larger atom like I holds them more loosely, so $\text{I}_{2}$ is a mild oxidiser (and the molecular solid is easy to sublime). This single variable — size, controlled by period — tracks most of the trends you will see in Sections 6–10.

[NEET Important] Hydrogen is usually called a "non-metal" even though its configuration $1s^{1}$ matches the alkali metals. In practice, H behaves like a non-metal: it is a gas, it has a high ionisation enthalpy (1312 kJ/mol, comparable to S), and it prefers covalent bonds in almost every compound. The single well-known exception is its ionic hydrides (e.g. NaH), where H forms H⁻.

Metalloids — The Staircase Class

Metalloids are the least famous but most technologically important elements of the lot. There are exactly seven of them — B, Si, Ge, As, Sb, Te, Po — and they all sit on or adjacent to the staircase line in the p-block. Some textbooks also count astatine (At), giving 8.

Why they are intermediate

Metalloids occupy the transition zone in two ways:

• Electronic: their ionisation enthalpies and electronegativities sit between the typical metal range (< 1000 kJ/mol, χ < 2.0) and the typical non-metal range (> 1100 kJ/mol, χ > 2.5).
• Structural: they form crystal structures that blend metallic and covalent bonding. Silicon, for example, crystallises in the diamond structure — a purely covalent framework — but behaves electrically like a poor metal.

The semi-conductor insight

The signature property of metalloids is semiconduction — they conduct electricity somewhat, and the conduction increases sharply with temperature (unlike true metals, whose conduction decreases).

This behaviour is explained by band theory:

• In a metal, the conduction band is partially filled — electrons are always available to carry current.
• In an insulator (diamond, NaCl), the gap between filled and empty bands is large (> 4 eV), so no electrons can jump into the conduction band at ordinary temperatures.
• In a metalloid, the gap is small — about 1.1 eV for silicon, 0.67 eV for germanium. At room temperature, a few thermally excited electrons make it across; at higher temperatures many more do.

This is the physical basis of the $\text{Si}$-based transistor. The entire electronics industry — computer chips, solar cells, LEDs — rests on the ability to dope Si or Ge with tiny amounts of other elements to precisely control their semi-conduction.

Chemical behaviour

Metalloids are chemically flexible:

• Boron behaves like a non-metal in most compounds ($\text{H}_{3}\text{BO}_{3}$ is a weak acid; $\text{BF}_{3}$ is a covalent Lewis acid), but $\text{B}_{2}\text{O}_{3}$ dissolves in molten bases to give borates.
• Silicon forms covalent $\text{SiO}_{2}$ (quartz), polymeric silicones, and covalent halides like $\text{SiCl}_{4}$ that hydrolyse in water.
• Arsenic shows both +3 and +5 oxidation states; $\text{As}_{2}\text{O}_{3}$ is amphoteric (dissolves in both acids and bases).

The staircase rule (memory aid)

Read the staircase from the top-left of the p-block going down-right, and the element that sits on the "flat" part of each step is a metalloid: B (between Groups 13/14), Si (between 14/15 on Period 3), Ge, As (Period 4), Sb, Te (Period 5), Po (Period 6). This is the canonical zig-zag.

[Board Level] If an exam question gives you an element on the staircase and asks whether it is a metal, non-metal or metalloid, pick metalloid unless you have a chemical reason to override it (e.g. Al is below the line despite being adjacent to B). The visual rule almost always wins.

The Four Block Families — A Chemistry Cheat-Sheet

This is the reference view we will use through the rest of the chapter: each of the four blocks (s, p, d, f) has a collective personality you can lean on. Learning one elementary member of each family unlocks the behaviour of the dozens of elements it contains.

s-block — alkali and alkaline-earth metals

• Valence config: $ns^{1-2}$
• Groups: 1 and 2 (14 elements in total, ignoring H)
• Headline traits: soft, silvery, low-melting, extremely reactive metals with low ionisation enthalpies. Lose the outer $ns$ electron(s) to form $\text{M}^{+}$ or $\text{M}^{2+}$ ions with noble-gas-like electronic cores.
• Iconic reactions: $2\,\text{Na} + 2\,\text{H}_{2}\text{O} \to 2\,\text{NaOH} + \text{H}_{2}\uparrow$ (violent), $\text{Mg} + \frac{1}{2}\text{O}_{2} \to \text{MgO}$ (white ash from dazzling flame).
• Oxides: strongly basic (alkalis, hence the name).
• Flame colours (diagnostic test): Li crimson, Na golden-yellow, K lilac, Rb red, Cs blue.

p-block — representative + noble gases

• Valence config: $ns^{2}\,np^{1-6}$ (except He, which is $1s^{2}$)
• Groups: 13 through 18 (36 elements across Periods 2–7)
• Headline traits: the chemically most diverse block — contains metals (Al, Sn, Pb, Bi), metalloids (B, Si, Ge, As, Sb, Te), non-metals (C, N, O, P, S, Se, halogens), and noble gases all in one block. Oxidation-state behaviour is very flexible — the inert-pair effect causes heavy p-block elements to prefer oxidation states two units below the group maximum (e.g. Tl⁺ more stable than Tl³⁺; Pb²⁺ more stable than Pb⁴⁺).
• Group names: 13 = boron family, 14 = carbon family, 15 = pnictogens, 16 = chalcogens, 17 = halogens, 18 = noble gases.

d-block — transition metals

• Valence config: $(n-1)d^{1-10}\,ns^{0-2}$
• Groups: 3–12, Periods 4–7 (40 elements)
• Headline traits: hard, high-melting, high-density metals with variable oxidation states (e.g. Mn: +2, +3, +4, +6, +7). Their partially-filled $d$-orbitals give rise to four signature properties:

1. Variable oxidation states — because removing successive electrons from $d$ and $s$ is energetically similar.
2. Coloured ions — $d{-}d$ electron transitions fall in visible frequencies ($\text{Cu(H}_{2}\text{O)}_{6}^{2+}$ blue, $\text{MnO}_{4}^{-}$ deep purple, $\text{Cr}_{2}\text{O}_{7}^{2-}$ orange).
3. Complex-ion formation — empty $d$-orbitals accept lone pairs from ligands; the entire field of coordination chemistry lives here.
4. Catalysis — Fe (Haber), $\text{V}_{2}\text{O}_{5}$ (contact process), Ni (hydrogenation), Pt/Pd (hydrogenation, catalytic converters).

• Exception-heavy: Cr, Cu, Mo, Ag, Pd, Au all have anomalous configurations (Section 4).

f-block — inner transition metals (lanthanoids + actinoids)

• Valence config: $(n-2)f^{1-14}\,(n-1)d^{0-1}\,ns^{2}$
• Members: 14 lanthanoids (Ce–Lu, Period 6) and 14 actinoids (Th–Lr, Period 7) — 28 elements shown as two separate rows beneath the main table.
• Headline traits:
• Lanthanoids: predominantly $+3$ oxidation state; almost indistinguishable chemistry within the row (lanthanoid contraction → small step-wise decrease in ionic radius from Ce³⁺ = 101 pm to Lu³⁺ = 85 pm).
• Actinoids: all radioactive; elements beyond U (Np onwards) are man-made; show a wider range of oxidation states (+3 to +7 for early members) because $5f$ electrons are less tightly held than $4f$.
• Applications: Nd magnets (speakers, electric motors), Eu/Tb phosphors (TV screens), U and Pu (nuclear fuel).

The four-block rule of thumb

If a question asks any chemistry question about an unnamed element, identifying its block narrows the answer to one of four stereotypes — and one of those stereotypes will almost always be the correct answer.

Noble Gases and a Few Exam Traps

We have mentioned noble gases in passing — now we give them their own spotlight and collect the edge-cases that examiners love to test.

Noble gases in one paragraph

Noble gases (He, Ne, Ar, Kr, Xe, Rn, Og — seven elements in Group 18) all share a completely-filled valence shell: $ns^{2}\,np^{6}$ for periods 2–7, and $1s^{2}$ for helium. This closed-shell configuration gives them the highest ionisation enthalpies in their respective periods, positive (endothermic) electron-gain enthalpies, and essentially no reactivity under ordinary conditions. They exist as monatomic gases at room temperature — no He₂, no Ne₂ (well, diatomic He₂ exists as an unbelievably weakly-bound van der Waals molecule at ultra-low temperatures, but it is irrelevant for Class 11).

The historical surprise came in 1962, when Neil Bartlett reported $\text{XePtF}_{6}$, proving that noble-gas chemistry was possible, just very difficult. Today, compounds of Xe (like $\text{XeF}_{2}, \text{XeF}_{4}, \text{XeO}_{3}$), some Kr compounds, and even a couple of Rn compounds are known.

Classic exam traps

Trap 1 — Liquid elements. Which elements are liquid at room temperature? Only two: mercury (Hg), a metal, and bromine (Br₂), a non-metal. Hydrogen, nitrogen and oxygen are gases; iodine and sulfur are solids. Do not answer "caesium" or "gallium" — those are liquid at 30 °C, not room temperature (25 °C).

Trap 2 — Graphite and diamond. Both are carbon (a non-metal), yet graphite conducts electricity (because of delocalised π-electrons in layers) and diamond does not. Diamond, meanwhile, is the best thermal conductor of any natural substance. Knowing these two counter-examples to the "non-metals don't conduct" rule is worth easy marks.

Trap 3 — Hydrogen placement. Hydrogen's configuration $1s^{1}$ matches Group 1 (alkali metals) — but H is a gas, a non-metal, and gives a diatomic molecule $\text{H}_{2}$. Some periodic tables place it above F at the top of Group 17 to reflect this, but NCERT follows the IUPAC convention: H at the top of Group 1, with a footnote noting its anomalous character. Accept either placement on Board exams; answer "Group 1" on JEE/NEET unless asked specifically why it could be Group 17.

Trap 4 — Helium's block. Helium ($1s^{2}$) is an s-block element by configuration but is placed in Group 18 because it is chemically a noble gas. Questions like "which element is an s-block element but sits in a p-block group?" or "which Group 18 element does not have the $ns^{2}np^{6}$ configuration?" have one answer: helium.

Trap 5 — Metalloid count. NCERT lists 7 metalloids: B, Si, Ge, As, Sb, Te, Po. Some competitive syllabi add At, giving 8. For CBSE Boards, memorise the NCERT list of 7.

Trap 6 — Alkaline-earth family name. Group 2 metals are called alkaline-earth metals — the full name, not "alkaline" or "earth metals" alone. Alkaline refers to the basic ("alkali"-like) hydroxides they form; "earth" is a historical name for oxides of these metals. Do not confuse with Group 1 (alkali metals).

Trap 7 — Group 18 former name. Mendeleev's original periodic table labelled the noble-gas column Group 0, because these elements were thought to have "zero" valency. IUPAC retired this label in 1988 after xenon and krypton compounds had been synthesised. The modern label is Group 18. If a question gives you "Group 0" as an option, it's testing whether you know the name is obsolete.

Trap 8 — Metals with low melting points. Caesium (28 °C), gallium (30 °C), and mercury (−39 °C) all melt near or below room temperature. Gallium melts in your hand. Caesium is solid in a sealed tube but liquefies on a hot day. Mercury is always liquid above −39 °C. Do not confuse "low-melting metal" with "non-metal" — all three are metallic in character.

[JEE Tip] The "match the column" type of question asking you to match elements to types/blocks/families is a free 4-marker if you have internalised this Section. Spend 30 seconds listing the 7 metalloids and the liquid elements before leaving the table — 95 % of classification questions fall into those categories.

Solved Examples

Example 1: Classifying an element from its configuration

An element has the ground-state configuration $[\text{Ne}]\,3s^{2}\,3p^{4}$. State whether it is a metal, non-metal, metalloid or noble gas, and name it.

Solution:

Let's decode the configuration:

• Largest $n = 3$ → Period 3
• Last-entering electron is in $3p$ → p-block
• Group = 10 + (2 + 4) = Group 16 (chalcogens)

Total electrons = 10 + 2 + 4 = 16, so the element is sulfur (S).

Sulfur sits above the staircase line in Group 16, so it is a non-metal. It is a yellow, brittle solid at room temperature — every physical property fits the non-metal profile.

Example 2: Identifying a metalloid

Element X has atomic number 32. Is X a metal, non-metal or metalloid? Justify using its position in the periodic table.

Solution:

$Z = 32$ → configuration $[\text{Ar}]\,3d^{10}\,4s^{2}\,4p^{2}$.

• Period 4, Group 14 (carbon family).
• Position: Group 14, Period 4 — this is germanium (Ge).
• Ge sits exactly on the staircase line — hence a metalloid.

Chemically, Ge behaves as expected: it forms a diamond-like covalent crystal, it is a semi-conductor (band gap 0.67 eV), and it dissolves in both acids and hot alkalis (like a non-metal).

Example 3: Block classification

Classify the following elements into s, p, d or f blocks: (a) Potassium (Z = 19) (b) Iron (Z = 26) (c) Arsenic (Z = 33) (d) Gadolinium (Z = 64)

Solution:

Write each ground-state Aufbau configuration and read off the classification.

Element	Configuration	Classification
(a) K	$[\text{Ar}]\,4s^{1}$	s-block
(b) Fe	$[\text{Ar}]\,3d^{6}\,4s^{2}$	d-block
(c) As	$[\text{Ar}]\,3d^{10}\,4s^{2}\,4p^{3}$	p-block
(d) Gd	$[\text{Xe}]\,4f^{7}\,5d^{1}\,6s^{2}$	f-block (lanthanoid by standard periodic-table convention)

Gadolinium has an exceptional configuration with a $5d^{1}$ electron, but it is still treated as a lanthanoid and placed in the f-block in the periodic table.

Example 4: Predicting oxide character

Predict whether the oxide of each of the following will be basic, acidic, or amphoteric: (a) $\text{Na}_{2}\text{O}$ (b) $\text{SO}_{3}$ (c) $\text{Al}_{2}\text{O}_{3}$ (d) $\text{CO}$ (e) $\text{MgO}$

Solution:

Use the rule: metal oxides are basic, non-metal oxides are acidic, and borderline-metal oxides are often amphoteric or neutral.

• (a) $\text{Na}_{2}\text{O}$: Na is a Group 1 metal → basic (dissolves in water to give NaOH).
• (b) $\text{SO}_{3}$: S is a non-metal → acidic (gives $\text{H}_{2}\text{SO}_{4}$ with water).
• (c) $\text{Al}_{2}\text{O}_{3}$: Al sits near the staircase line and behaves borderline → amphoteric (dissolves in both HCl and NaOH).
• (d) $\text{CO}$: one of the few neutral oxides (along with NO and $\text{N}_{2}\text{O}$) — does not react with water, acids, or bases under ordinary conditions.
• (e) $\text{MgO}$: Mg is a Group 2 metal → basic (but far less soluble than $\text{Na}_{2}\text{O}$).

[JEE Tip] Amphoteric oxides: $\text{BeO}, \text{Al}_{2}\text{O}_{3}, \text{ZnO}, \text{SnO}, \text{PbO}, \text{Ga}_{2}\text{O}_{3}$. Neutral oxides: CO, NO, $\text{N}_{2}\text{O}$.

Example 5: Why only Hg is liquid among metals

Explain in one sentence why mercury is the only metal that is liquid at room temperature, while its neighbours zinc and cadmium (same group, one and two periods up) are solid.

Solution:

In mercury, the $6s^{2}$ electrons are pulled in very close to the nucleus by relativistic contraction — an effect that grows strongly for very heavy elements. The contracted $6s^{2}$ pair becomes unusually stable, so Hg atoms contribute their outer electrons only weakly to the metallic bonding. The bond energy between Hg atoms is therefore small enough that the metal melts at −39 °C.

Zn and Cd, being lighter, have much weaker relativistic effects; their outer $s^{2}$ electrons remain available for metallic bonding, giving normal melting points of 420 °C and 321 °C.

Example 6: Metalloid count and identification

(a) How many elements does NCERT classify as metalloids, and what are they? (b) Which of the following is/are metalloid(s)? C, Si, As, Se, Ge

Solution:

(a) NCERT recognises 7 metalloids: B, Si, Ge, As, Sb, Te, Po. They sit on or adjacent to the zig-zag staircase in the p-block.

(b) Checking each against the list:

• C — carbon, non-metal (above the staircase). Not a metalloid.
• Si — silicon, metalloid ✓
• As — arsenic, metalloid ✓
• Se — selenium, non-metal. Not a metalloid.
• Ge — germanium, metalloid ✓

So the metalloids in the list are Si, As, Ge.

Example 7: Non-metal or metal — hydrogen's case

Justify, in 3–4 lines, whether hydrogen should be classified as a metal or a non-metal.

Solution:

Hydrogen is best classified as a non-metal, even though its electronic configuration ($1s^{1}$) matches the alkali metals. The evidence:

• Physical state: a gas at room temperature (unlike all other Group 1 elements, which are soft metals).
• Ionisation enthalpy: 1312 kJ/mol — much higher than Li (520 kJ/mol), and comparable to the halogens.
• Bonding: forms a diatomic molecule $\text{H}_{2}$ with a strong covalent bond — more like a halogen than a Group-1 metal.
• Chemistry: prefers to share its electron covalently in most compounds, only occasionally forming $\text{H}^{+}$ or $\text{H}^{-}$ ionically.

These observations outweigh the purely configurational alkali-metal similarity. NCERT places H at the top of Group 1 with a footnote noting the ambiguity.

Example 8: Semi-conductor band gap

Explain why silicon (band gap $E_{g} \approx 1.12\ \text{eV}$) is a semi-conductor, while diamond ($E_{g} \approx 5.5\ \text{eV}$) is an insulator, even though both have identical tetrahedral covalent structures.

Solution:

In both materials, valence electrons occupy a filled valence band separated from an empty conduction band by an energy gap $E_{g}$. To conduct, electrons must be promoted across this gap.

At room temperature, the thermal energy available is $k_{B}T \approx 0.026$ eV, so the number of thermally excited electrons depends exponentially on the size of $E_g$.

• Silicon: $E_{g}$ is moderate, so a small but non-zero number of electrons can cross into the conduction band. Hence silicon conducts weakly and behaves as a semi-conductor.
• Diamond: $E_{g}$ is very large, so essentially no electrons cross into the conduction band at room temperature. Hence diamond is an insulator.

Takeaway: The moderate band gap of metalloids like Si and Ge places them in the semi-conductor sweet spot.

Example 9: Transition-metal properties checklist

Among the following elements, identify which are transition metals, and for each, state one characteristic property. (a) Ca (b) Fe (c) Zn (d) Mn (e) Al

Solution:

Transition metals are d-block elements whose atoms or common ions have partially-filled $d$-orbitals (IUPAC definition).

• (a) Ca ($[\text{Ar}]\,4s^{2}$) — s-block, not a transition metal.
• (b) Fe ($[\text{Ar}]\,3d^{6}\,4s^{2}$) — transition metal. Property: shows +2 and +3 oxidation states.
• (c) Zn ($[\text{Ar}]\,3d^{10}\,4s^{2}$) — d-block but not technically a transition metal, because both the atom and $\text{Zn}^{2+}$ ion have full $d^{10}$ configuration.
• (d) Mn ($[\text{Ar}]\,3d^{5}\,4s^{2}$) — transition metal. Property: shows a wide range of oxidation states.
• (e) Al ($[\text{Ne}]\,3s^{2}\,3p^{1}$) — p-block, not a transition metal.

Correct answer: Fe and Mn are definitely transition metals; Zn is d-block but not a true transition metal. Ca and Al are not transition metals.

Example 10: Electron configuration of a noble gas

Write the electronic configuration of xenon (Z = 54) and use it to justify why xenon is classified as a noble gas.

Solution:

Xenon has the Aufbau configuration:

$$1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{2}\,3d^{10}\,4p^{6}\,5s^{2}\,4d^{10}\,5p^{6}$$

Sum of superscripts: $2+2+6+2+6+2+10+6+2+10+6 = 54$. ✓

In noble-gas shorthand: $[\text{Kr}]\,4d^{10}\,5s^{2}\,5p^{6}$.

The outermost shell has configuration $5s^{2}\,5p^{6}$ — a completely-filled valence octet. This is the definition of a noble gas. Consequences:

• Extremely high first ionisation enthalpy.
• Electron-gain enthalpy is positive (endothermic).
• Chemically nearly inert, though Xe can form compounds with fluorine and oxygen.

Xe therefore sits at the rightmost edge of Period 5 in Group 18.

Example 11: Which element does NOT fit the family?

From the following set — Na, Mg, K, Al, Ca — identify the element that does not belong to the same block as the rest, and state the block of each.

Solution:

Write configurations:

• Na ($[\text{Ne}]\,3s^{1}$) → s-block
• Mg ($[\text{Ne}]\,3s^{2}$) → s-block
• K ($[\text{Ar}]\,4s^{1}$) → s-block
• Al ($[\text{Ne}]\,3s^{2}\,3p^{1}$) → p-block
• Ca ($[\text{Ar}]\,4s^{2}$) → s-block

Four of the five (Na, Mg, K, Ca) are s-block metals; Al is p-block and is the odd one out.

Takeaway: even though Al is a metal physically and chemically, by block classification it is a p-block element.

Example 12: School exam-style 3-marker on classification

(a) State the position of an element in the modern periodic table whose electronic configuration is $[\text{Kr}]\,4d^{5}\,5s^{1}$. (b) Identify the element and state its block, period and group. (c) Is it a metal, non-metal or metalloid? Give one reason.

Solution:

(a) Configuration $[\text{Kr}]\,4d^{5}\,5s^{1}$. Total electrons = 36 + 5 + 1 = 42. This is molybdenum (Mo, Z = 42). Largest $n = 5$ → Period 5. Last-entering electron in $4d$ → d-block. Group for d-block = valence $s$ + $(n-1)d$ = 1 + 5 = Group 6.

(b) Mo is a d-block, Period 5, Group 6 element. It is one of the Aufbau anomalies — the $4d^{5}\,5s^{1}$ configuration (half-filled $d$) is lower in energy than the naive $4d^{4}\,5s^{2}$.

(c) Mo is a metal — specifically, a transition metal. Reason: it lies in the d-block to the left of the staircase line and shows variable oxidation states.