Exercises — Metals, non-metals, metalloids — properties

Use position + grip on electrons.

• K — far bottom-left, one loose outer electron, tiny price tag → metal.
• S — top-right region (period 3, group 16), grabs electrons → non-metal.
• Si — sits on the staircase → metalloid (semiconductor).
• Cu — a transition element, clearly a shiny conductor → metal.
• Ne — noble gas; full shell, does not want to lose or gain → conventionally non-metal (we exclude noble gases from the metallic-character trend).

Dull + brittle + insulator = electrons locked in fixed bonds, none free to roam. Acidic oxide is the chemical fingerprint of a non-metal (see Acidic and Basic Oxides). X is a non-metal.

Left → right across period 3 the order is Na, Mg, Al.

• Across a period the effective nuclear charge — the net positive pull the outer electron actually feels from the nucleus after the inner electrons partly cancel it — rises, and the radius shrinks → the outer electron is held tighter → IE₁ rises (Na 496 < Mg 738 < Al 578... careful: here Na < Mg holds, and although Al dips, all three still keep Na cheapest).
• Using real values IE₁: Na 496 < Al 578 < Mg 738 (kJ/mol). Metallic character ∝ 1/IE₁, so most metallic = lowest IE₁.
• Ranking by lowest IE₁: Na, then Al, then Mg.
• So metallic character: Na > Al > Mg.

(Note: the naïve "Na > Mg > Al" you may have written assumes IE₁ rises perfectly smoothly. It does not — Al's lone 3p electron is cheaper to remove than Mg's paired 3s electrons. Level 5 explains this dip in full.)

• Na is a metal → Na₂O is a metal oxide → with water gives NaOH, a base.
• P is a non-metal → P₄O₁₀ is a non-metal oxide → with water gives H₃PO₄, an acid. Root cause: metals lose electrons to form positive-ion oxides that release OH⁻; non-metals form covalent oxides that release H⁺. See Acidic and Basic Oxides.

Going down group 1, radius grows and inner electrons shield the outer one more, so the effective nuclear charge felt by the outer electron drops and the price tag (IE₁) falls. Lower IE₁ → more metallic. Cs is more metallic than Li. See Periodic Trends — Atomic Radius.

• Metalloid (semiconductor): electrons need a small energy nudge to become free. Heat supplies that nudge → more free carriers → conductivity rises. So A is the metalloid.
• Metal: already has a full sea of free electrons. Heat makes the ions vibrate more, which scatters the drifting electrons → conductivity falls. So B is the metal. The two obey opposite temperature rules — a definitive fingerprint. See Semiconductors and Doping.

Reason from are the electrons free?, not from the label. Each carbon in graphite uses only 3 of its 4 outer electrons in bonds; the 4th is delocalised between the flat layers — a partial "electron sea." Those free electrons drift under a voltage, so graphite conducts along its layers. The general rule ("non-metals insulate") assumes all valence electrons are locked in bonds; graphite's structure breaks that assumption. Structure can override the label.

The safe method: metallic character ∝ 1/IE₁, so most metallic = lowest IE₁. Get the trend intuition, then confirm with numbers.

Trend setup:

• Al vs Ga (same group 13, Ga below Al): down a group IE₁ falls → Ga should be more metallic than Al.
• Mg vs Al (same period 3): the naïve trend says Mg (further left) is more metallic — but period-3 has an anomaly right here, so we must check the data before trusting it.

Check the data (first IE, kJ/mol): $\text{Mg}=738,\text{Al}=578,\text{Ga}=579.$

• Al (578) < Ga (579): Al's IE₁ is just below Ga's, so Al is very slightly more metallic than Ga — the naïve "down-a-group" expectation is nearly a tie and actually reverses here (a d-block/relativistic quirk in Ga).
• Both Al and Ga (≈578–579) < Mg (738): so Al and Ga are both more metallic than Mg.

Ranking by lowest IE₁ (578 < 579 < 738): $\text{Al}>\text{Ga}>\text{Mg}.$ Lesson: the simple "Mg beats Al because it's further left" is the anomaly-trap — Al's lone 3p electron is cheaper to remove than Mg's paired 3s pair, so Al outranks Mg. Always sanity-check period-3 comparisons against IE₁ values. (See Electronic Configuration for why the 3s pair clings harder.)

Chain the whole page together:

1. Si is a metalloid on the staircase → its electrons are almost free.
2. "Almost free" means conduction is controllable: a tiny energy nudge (heat, light) or adding impurity atoms (doping) switches conduction on.
3. Copper (metal) is always on — you cannot turn its sea off, so it can't act as a switch.
4. Sulfur (non-metal) is always off — an insulator, no free carriers to switch on.
5. A microchip needs billions of tiny on/off switches; only the in-between metalloid gives you both states. Hence silicon. See Semiconductors and Doping.

Metallic bonding controls conductivity, lustre, malleability, not hardness. Hardness depends on how strongly the ion lattice resists deformation — a separate variable.

• Mercury has a mobile electron sea (it conducts, it's lustrous) but weak ion–ion binding at room temperature → it's a liquid metal. Still every inch a metal.
• Sodium loses its electron so easily (very low IE₁) that its lattice is soft — yet it conducts and is shiny. Softness confirms loose bonding, it doesn't disqualify it. Verdict: "hard" ≠ "metal." Test metallic character by ease of first-electron loss / conductivity / lustre, never by hardness.

Metallic character ∝ 1/IE₁, so most metallic = lowest IE₁. Sort IE₁ ascending: $496<578<738<786<1000<1012<1251$ maps to Na < Al < Mg < Si < S < P < Cl in IE₁, so metallic character (most → least): $\text{Na}>\text{Al}>\text{Mg}>\text{Si}>\text{S}>\text{P}>\text{Cl}.$ The anomaly: the smooth "rises left → right" trend breaks twice —

• Al (578) < Mg (738): removing Al's lone 3p electron is cheaper than breaking Mg's stable filled 3s pair.
• S (1000) < P (1012): P has a stable half-filled 3p³; S must pair an electron, which is slightly easier to remove. So metallic character mostly falls left → right but has two small bumps. (See Electronic Configuration for why half-filled and filled subshells are extra stable.)