Chemistry 2e · Representative Metals, Metalloids, and Nonmetals
Periodicity
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Periodicity Repeating patterns in element properties down the table Full entry → is the recurring pattern of element properties that Mendeleev exploited to build the periodic table: elements in the same column behave alike because they share the same valence-electron configuration. This topic bridges the atomic theory of earlier chapters and the descriptive chemistry of the rest of Chapter 18. Here you will learn the physical basis of the trends — Atomic radius Distance from nucleus to the outermost electron Full entry →, Ionization energy Energy to remove one electron from a gaseous atom Full entry →, Electron affinity Energy change when a gaseous atom gains an electron, Electronegativity How strongly an atom attracts shared bonding electrons Full entry →, and metallic character — and use them to organize the Representative elements Main-group (s- and p-block) elements Full entry → (the s- and p-block main group) into metals, metalloids, and nonmetals.
Why this matters
Every prediction in the coming topics — why sodium reacts violently with water while gold does not, why silicon is a semiconductor while carbon and lead are not — follows from periodicity. In practical terms, the trends decide which elements form ionic versus covalent compounds, which metals corrode, which ions are stable in biology (Na⁺, K⁺, Ca²⁺, Cl⁻), and how semiconductors are doped. Mastering the trends and their physical cause (effective nuclear charge) lets you predict behavior for elements you have never studied individually — the same skill a chemist uses to design materials.
The college version
Core Concepts
The physical engine: effective nuclear charge
Valence electrons are pulled toward the nucleus by a charge that is less than the full nuclear charge, because inner (core) electrons shield them. The effective nuclear charge is
Zeff = Z - S
where Z is the atomic number and S is the shielding constant — roughly the number of core electrons for a valence electron. Across a period, the core stays the same size while Z climbs, so Zeff increases steadily; down a group, each step adds a whole new shell of core electrons, so shielding grows and the valence electron sits farther from the nucleus with a similar Zeff. This one idea explains nearly every trend.
Atomic radius and ionization energy
Atomic radius decreases across a period (higher Zeff pulls the same shell tighter) and increases down a group (new shells add distance). Ionization energy — the energy to remove an electron — is the mirror image: it increases across a period and decreases down a group, because a valence electron held more tightly is harder to remove. The first ionization energies of the third period illustrate the general climb: Na 496, Mg 738, Al 578, Si 786, P 1012, S 1000, and Cl 1251 kJ/mol (commonly taught values). Notice the two small dips (Al below Mg, S below P): the 3p electron of Al is shielded by the 3s² pair, and removing an electron from the half-filled 3p³ of P costs extra stability. These bumps are the exceptions exam questions love.
Electron affinity and electronegativity
Electron affinity is the energy change when a neutral atom gains an electron. Across a period it generally becomes more negative as Zeff rises; down a group it becomes less negative, with notable exceptions — chlorine has a more negative electron affinity than fluorine, because fluorine's small size crowds the added electron. Electronegativity (Pauling's scale) measures how strongly an atom attracts bonding electrons: it increases across a period and decreases down a group, so fluorine (3.98) is the most electronegative element. The difference between two bonded atoms' electronegativities predicts bond polarity and whether the compound leans ionic or covalent.
Metals, metalloids, and nonmetals
The representative elements divide into three bands. Metals (left and lower regions): few valence electrons, low ionization energies, lose electrons to form cations, shiny, malleable, good conductors. Nonmetals (right and upper regions): high ionization energies and electronegativities, gain or share electrons, poor conductors; hydrogen is the odd one out — placed above the alkali metals but behaving as a nonmetal. Metalloids sit along the staircase between them (boron, silicon, germanium, arsenic, antimony, tellurium, sometimes astatine) with intermediate properties, notably semiconducting behavior — the basis of modern electronics. Within a group, metallic character increases downward: carbon a nonmetal, silicon a Metalloid Element with intermediate metal/nonmetal properties Full entry →, lead a metal.
Diagonal relationships
Adjacent elements in different groups and periods sometimes resemble each other more than their own group members — lithium resembles magnesium, beryllium resembles aluminum, and boron resembles silicon. The cause is a near match in charge-to-size ratio: one step down a group (bigger radius) and one step right (higher charge) keep the ionic charge density similar. These diagonal relationships explain real chemistry, such as the likeness of lithium and magnesium compounds.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Atomic radius trend | Ionic radius trend | Cations are smaller than their atoms; anions are larger — the radius rules differ for ions |
| Ionization energy | Electron affinity | IE is energy to remove an electron; EA is energy change on gaining one |
| "First" ionization energy | Total ionization | Removing a second electron costs far more energy (big jump after a noble-gas core) |
| Hydrogen's group | Hydrogen's behavior | H sits above Li but is a nonmetal — it forms H⁺, H⁻, and covalent bonds |
| Metalloid | Metal or nonmetal | Metalloids have intermediate properties; their semiconductivity is what matters |
| Trend across a period | Trend down a group | Across: radii shrink, IE rises; down: radii grow, IE falls — easy to invert on exams |
Eli explains
The same idea, in plain words
Explain it like I’m 10
The periodic table is like a playground slide made of atom properties. Going across a row, the atoms' centers pull harder on their outer electrons, so the atoms shrink and cling tighter — harder to pull an electron off. Going down a column, atoms add new layers, so they get bigger and hold their outer electrons loosely. That's why sodium is a wild metal that fizzes in water, while chlorine is a nasty gas that grabs an electron.
Worked example
Worked Example 1 — Effective nuclear charge across a period. Estimate Zeff for a valence electron in sodium (Z = 11) and chlorine (Z = 17). Both third-period atoms have 10 core electrons (1s²2s²2p⁶), so S ≈ 10:
Na: Zeff = 11 - 10 = +1 Cl: Zeff = 17 - 10 = +7
(Refined shielding models, such as Slater's rules, give about +2.2 and +6.1.) Chlorine's valence electron feels roughly seven times the pull of sodium's — which is why its atom is smaller, its ionization energy about 2.5 times larger, and it grabs electrons rather than donating them.
Worked Example 2 — Coulomb's law quantifies the pull. The electrostatic force on a valence electron is F = k Zeff e2r2, so F ∝ Zeffr2. Using approximate radii of 186 pm (Na) and 99 pm (Cl):
FClFNa = 7/(99)21/(186)2 = 79801 × 34,5961 ≈ 25
The valence electron of chlorine is pulled about 25 times harder than sodium's — explaining the third period's story: smaller atoms, higher ionization energies, and a shift from metallic (Na, Mg, Al) through metalloid (Si) to nonmetallic (P, S, Cl) behavior.
Key takeaways
- Zeff = Z - S: across a period it rises (same core, more protons); down a group it is roughly constant (new shells).
- Atomic radius: decreases across, increases down. Ionization energy: increases across, decreases down.
- First-ionization dips: Al < Mg, S < P (3p shielding; half-filled-subshell stability).
- Electron affinity generally becomes more negative across a period; Cl beats F because F's small size crowds the added electron.
- Electronegativity (Pauling): increases across, decreases down; F ≈ 3.98 is the most electronegative.
- Representative elements = s- and p-block; metals lose electrons, nonmetals gain or share, metalloids (B, Si, Ge, As, Sb, Te) are in between and semiconducting.
- Metallic character increases down a group (C → Si → Ge → Sn → Pb).
- Diagonal relationships (Li–Mg, Be–Al, B–Si) come from similar charge density.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Explain in one sentence why atomic radius decreases across a period.
Show answer
Across a period the core stays constant while the nuclear charge grows, so Zeff rises and pulls the shell tighter.
Rank Na, Mg, and Cl by increasing first ionization energy, and give the approximate trend.
Show answer
Na (496) < Mg (738) < Cl (1251) kJ/mol — IE increases across a period.
Why is aluminum's first ionization energy lower than magnesium's?
Show answer
Aluminum's valence electron is in the 3p subshell, shielded by the 3s² pair, so it is slightly easier to remove than magnesium's 3s electron.
Which is more electronegative, F or Cl, and why does Cl have the more negative electron affinity?
Show answer
Fluorine is more electronegative (about 3.98 vs. 3.16). But chlorine has the more negative electron affinity, because fluorine's tiny atom crowds the added electron with electron–electron repulsion.
Identify the metalloids in the third period and explain what makes them special.
Show answer
Silicon. Metalloids are semiconductors — the foundation of electronics.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- Periodicity
- Repeating patterns in element properties down the table
- Effective nuclear charge Zeff
- Net pull on a valence electron after shielding
- Atomic radius
- Distance from nucleus to the outermost electron
- Ionization energy
- Energy to remove one electron from a gaseous atom
- Electron affinity
- Energy change when a gaseous atom gains an electron
- Electronegativity
- How strongly an atom attracts shared bonding electrons
- Metalloid
- Element with intermediate metal/nonmetal properties
- Representative elements
- Main-group (s- and p-block) elements
- Effective nuclear charge (Zeff)
- Net nuclear charge felt by a valence electron after shielding
Sources & references
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