S- and P- Block Elements

The s-block (Groups 1–2) and p-block (Groups 13–18) together make up the "main-group" elements. On the MCAT this material is tested almost entirely through periodic trends — how atomic and ionic radius, ionization energy, electron affinity, electronegativity, and metallic character vary across a period and down a group, and the effective-nuclear-charge reasoning behind each. Descriptive group-by-group reactions are low-yield, so we cover only the few patterns the MCAT actually rewards.

MCAT focus (Chem/Phys, Content Category 5E). The high-yield core is periodic trends and the reasoning behind them (Z_eff, shielding, shell number). Know how to predict and compare radius, ionization energy, and electronegativity for any two main-group elements.

S/P/D/F Block Elements (overview)

The block to which an element belongs is decided by the orbital being filled by the last electron in its ground-state configuration.

s-block — last electron enters an s orbital. Groups I (alkali metals) and II (alkaline-earth metals); also H and He. General configuration ns¹ or ns². Soft, low density, low IE, highly reactive metals (except H/He).
p-block — last electron enters a p orbital. Groups III to VIII / 13–18. Configuration ns²np^1–6. Includes metals (Al, Sn, Pb), metalloids (B, Si, Ge, As), and all the non-metals plus noble gases.
d-block (transition elements) — last electron enters a (n−1)d orbital. Groups 3–12. Configuration (n−1)d^1–10ns^0–2. Hard, high melting points, variable oxidation states, often colored compounds.
f-block (inner transition) — last electron enters an (n−2)f orbital. Lanthanides (4f, Ce–Lu) and actinides (5f, Th–Lr). Many actinides are radioactive.

Periodic Trends (the high-yield core)

Nearly every trend follows from two competing factors acting on the outermost (valence) electrons:

Effective nuclear charge, Z_eff ≈ Z − S — the net positive pull the valence electron feels after inner electrons "shield" (screen) part of the nuclear charge Z. Higher Z_eff pulls electrons in tighter.
Principal quantum number, n — higher shells are physically larger and farther from the nucleus.

Across a period, n is constant but Z_eff rises (each added proton is only weakly shielded by same-shell electrons) → atoms contract and hold electrons more tightly. Down a group, n increases and inner-shell shielding grows → atoms expand and hold electrons more loosely. Keep these two levers in mind and you can derive every trend below.

Atomic radius

Across a period (L → R): decreases — rising Z_eff in the same shell pulls electrons closer.
Down a group (T → B): increases — each new period adds a shell (larger n), and inner shells shield the valence electrons.

Ionic radius

Cations are smaller than their parent atoms (lose the outer shell and/or reduce electron–electron repulsion, so remaining electrons are pulled in). Anions are larger than their parent atoms (added electrons increase repulsion). For an isoelectronic series (same electron count, e.g. N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺), radius decreases as nuclear charge increases — more protons pulling on the same number of electrons.

Ionization energy (IE)

Energy to remove the most loosely bound electron from a gaseous atom (X(g) → X⁺(g) + e⁻). Opposite trend to atomic radius: increases across a period, decreases down a group.
Successive IEs always rise (IE₁ < IE₂ < IE₃…) because you are pulling an electron off an increasingly positive ion. A large jump appears when you start removing a core (noble-gas) electron — e.g. the jump after IE₂ for Mg reveals its two valence electrons.
Two classic anomalies: Be > B (removing from the higher-energy 2p of B is easier than from the filled 2s of Be) and N > O (O must pair an electron in a 2p orbital, and pairing repulsion lowers its IE).

Electron affinity (EA)

Energy released when a gaseous atom gains an electron. Generally becomes more exothermic (more negative) across a period and less so down a group. Halogens have the most favorable EA. Note the small-atom anomaly: Cl has a more negative EA than F because F's tiny 2p shell forces strong electron–electron repulsion on the incoming electron.

Electronegativity (EN)

Tendency of a bonded atom to attract the shared electron pair. Tracks IE and EA: increases across, decreases down. F is the most electronegative element (Pauling 4.0). EN differences drive bond polarity: ΔEN ≈ 0 nonpolar covalent, intermediate polar covalent, large (>~1.7) ionic.

Metallic character & acid/base of oxides

Metallic (electron-donating) character decreases across a period and increases down a group — the mirror image of ionization energy. As a corollary, main-group oxides trend from basic (metals, left) to amphoteric (e.g. Al₂O₃) to acidic (nonmetals, right). Most reactive metal → toward Cs/Fr; most reactive nonmetal → F.

High-Yield Group Highlights

The MCAT rarely asks you to recall specific descriptive reactions. It wants the trends and the reasoning. These are the few group facts worth knowing, framed as consequences of the periodic trends above.

Group 1 — alkali metals (ns¹)

Lowest ionization energies in each period → lose one electron readily → strongest metallic (reducing) character; always +1. Reactivity increases down the group (Cs > … > Li) because IE falls. React with water to give MOH (a strong base) + H₂.

Group 2 — alkaline-earth metals (ns²)

Higher IE than Group 1 (smaller, higher Z_eff) but still reactive metals; always +2. Reactivity again increases down the group. Biologically important: Ca²⁺ (bone, signaling, muscle) and Mg²⁺ (enzyme cofactor, ATP, chlorophyll).

Groups 13–18 — the p-block

Spans metals (Al), metalloids (B, Si), nonmetals, and the noble gases. Character shifts from metallic (left) to nonmetallic (right). Highlights: carbon's catenation (strong, stable C–C bonds) underpins all of organic chemistry; the halogens (Group 17) are the most reactive nonmetals with the most favorable electron affinities; the noble gases (Group 18) are inert because of their full valence shells and very high IE.

Common trap. Atomic radius and ionization energy are inversely related, but do not confuse an atom with its ion. A cation is always smaller than its parent atom and an anion always larger — and within an isoelectronic series, more protons means a smaller radius (compare Na⁺ < Ne < F⁻).

One rule to anchor them all. Ionization energy, electron affinity magnitude, and electronegativity all trend the same way (up and to the right, peaking near F); atomic and metallic character trend the opposite way. If you remember the radius trend and that "small + high Z_eff = holds electrons tightly," you can reconstruct the rest.

Worked MCQs

Five MCQs that capture the high-yield testing patterns for this chapter. Read the explanation even when you get the answer right — it's where the deeper concept lives.

Q1. Across a period from left to right, atomic radius generally:

Decreases
Increases
Remains constant
First decreases then increases

Across a period, electrons are added to the same shell while nuclear charge increases. Effective nuclear charge rises and pulls outer electrons closer, so atomic radius decreases.

Q2. The order of reactivity of alkali metals with water is:

Li > Na > K > Rb > Cs
Cs > Rb > K > Na > Li
Na > K > Li > Rb > Cs
K > Na > Cs > Li > Rb

Going down Group 1, atomic radius increases and ionization energy decreases — the single ns¹ electron is lost more easily. Therefore reactivity (a reducing/electron-donating property) rises Cs > Rb > K > Na > Li.

Q3. The species N³⁻, O²⁻, F⁻, Na⁺, and Mg²⁺ are isoelectronic (10 electrons each). Which has the smallest ionic radius?

N³⁻
O²⁻
F⁻
Mg²⁺

Within an isoelectronic series every species has the same number of electrons, so radius is set by nuclear charge. Mg has the most protons (Z = 12) pulling on those 10 electrons, giving the strongest inward pull and the smallest radius. Radius order: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺.

Q4. Oxygen has a lower first ionization energy than nitrogen, breaking the general left-to-right increase. The best explanation is that:

Oxygen has a larger atomic radius than nitrogen
Nitrogen has a higher nuclear charge than oxygen
Removing an electron from oxygen relieves electron–electron repulsion in a doubly occupied 2p orbital
Nitrogen has a completely filled 2p subshell

Nitrogen is 2p³ (each 2p orbital singly occupied, a stable half-filled set). Oxygen is 2p⁴, so one 2p orbital holds a pair; the repulsion between those paired electrons makes oxygen's fourth 2p electron easier to remove, lowering its IE below nitrogen's. (The analogous Be > B anomaly comes from the filled 2s of beryllium.)

Q5. An element with electronic configuration [Ar] 3d¹⁰ 4s² 4p³ belongs to which block?

s-block
p-block
d-block
f-block

The last electron entered the 4p orbital, so the element is in the p-block (it is As, Group V). Block assignment depends on which orbital receives the differentiating electron, not on which orbitals are filled overall.

Quick Recap

Every trend traces back to two levers: effective nuclear charge (Z_eff) and shell number (n).
Atomic radius: ↓ across, ↑ down. Ionization energy, electron affinity magnitude, and electronegativity: ↑ across, ↓ down (all peak near F).
Cations < parent atom; anions > parent atom. Isoelectronic series: more protons → smaller radius.
Successive IEs rise; a large jump signals entering a noble-gas core. Anomalies: Be > B, N > O.
Metallic character ↓ across, ↑ down; oxides shift basic → amphoteric → acidic across a period.
Group 1 = strongest reducers; halogens = most reactive nonmetals; noble gases inert. Carbon's catenation underlies organic chemistry.

Test yourself. Take a timed S/P-Block quiz or browse all Chemistry MCQs to lock these concepts in.