Periodic Table

Periodic Table

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The periodic table arranges all known elements in order of increasing atomic number (Z), which equals the number of protons in the nucleus. Horizontal rows are called periods; vertical columns are called groups. The periodic law states that the properties of elements repeat at regular intervals, which is why elements in the same group behave alike.

Two must-know relationships govern nuclear composition:

• Z = p = e in a neutral atom (proton count equals electron count).
• A = Z + n, where A is mass number, Z is atomic number, and n is neutron count.

For NABTEB Chemistry, high-yield points are: (1) valency of main-group elements equals their group number for Groups 1–2 and (10 − group number) for Groups 13–17; (2) ionization energy increases across a period and decreases down a group; (3) electronegativity follows the same pattern — highest at fluorine, lowest at francium/cesium.

🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Structure of the Modern Table

The modern periodic table is built on Moseley’s law, which uses X-ray frequencies to show that atomic number, not atomic mass, is the fundamental ordering principle. Mendeleev’s earlier table had problems (e.g., Co/Ni and Ar/K inversions) that Moseley resolved by reordering.

• Periods (1–7) correspond to the filling of successive electron shells with principal quantum number n.
• Groups (1–18) contain elements with the same number of valence electrons, which is why they share similar chemistry.
• The table is divided into four blocks: s-block (Groups 1–2), p-block (Groups 13–18), d-block (transition metals, Groups 3–12), and f-block (lanthanides and actinides placed separately below).

Trends Across a Period (e.g. Period 3: Na → Ar)

Property	Trend	Reason
Atomic radius	Decreases	Increasing nuclear charge pulls electrons closer
Ionization energy	Generally increases	Tighter electron binding
Electronegativity	Increases	Stronger pull on bonding electrons
Metallic character	Decreases	Fewer delocalised electrons, higher IE

Trends Down a Group (e.g. Group 1: Li → Cs)

• Atomic radius increases (new electron shell added).
• Ionization energy decreases (outer electron farther from nucleus and shielded).
• Reactivity of alkali metals increases down the group.
• For Group 17 halogens, reactivity decreases down the group because electron affinity drops.

Key Families

• Group 1 — Alkali metals: Li, Na, K, Rb, Cs, Fr. Soft, low-density metals forming +1 ions. Stored under oil to prevent reaction with air/water.
• Group 2 — Alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra. Form +2 ions, harder than Group 1.
• Group 17 — Halogens: F, Cl, Br, I, At. Reactive non-metals forming −1 ions.
• Group 18 — Noble gases: He, Ne, Ar, Kr, Xe, Rn. Have complete octets (He has a duplet) and are largely inert.

Worked Relationship

For $^{23}_{11}\text{Na}$: protons = 11, electrons = 11, neutrons = 23 − 11 = 12.

Isotopes

Isotopes are atoms of the same element with the same Z but different A (different neutron counts). Example: $^{12}{6}\text{C}$, $^{13}{6}\text{C}$, $^{14}_{6}\text{C}$ are isotopes of carbon.

NABTEB Question Patterns

• Identifying the group/period of an element from its electron configuration.
• Predicting valency and formula of a compound (e.g., oxide of Group 2 element X is XO).
• Comparing properties such as atomic radius, IE, or reactivity between two elements.
• Distinguishing isotopes from isobars or atoms of different elements.

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Why Trends Sometimes Break

Ionization energy does not increase smoothly across a period. The IE jumps between Group 2 → Group 13 (Be to B: B’s outer 2p electron is higher in energy and easier to remove than Be’s filled 2s sub-shell) and between Group 15 → Group 16 (N to O: O has a paired electron in 2p that repels its partner, lowering the removal energy). These anomalies are favourite NABTEB/WAEC assertion-reason traps.

Transition Metals and Variable Valency

d-block elements show variable oxidation states because the (n−1)d and ns electrons have similar energies. Examples:

• Fe: Fe²⁺ and Fe³⁺
• Cu: Cu⁺ and Cu²⁺
• Mn: Mn²⁺, Mn⁴⁺, Mn⁶⁺, Mn⁷⁺

They also form coloured compounds (due to d–d electronic transitions) and act as catalysts (Fe in Haber process, Ni in hydrogenation, V₂O₅ in Contact process).

Metallic vs Non-metallic Character

• The metallic–non-metallic divide runs as a staircase from B to At on the right of the table.
• Oxides of metals are generally basic (Na₂O, MgO); oxides of non-metals are acidic (SO₃, CO₂, P₄O₁₀).
• Aluminium and a few others near the staircase are amphoteric (Al₂O₃, ZnO react with both acids and bases).

Periodicity in Bonding

Across a period, bonding shifts from metallic (Na, Mg, Al) to giant covalent (Si) to simple molecular (P₄, S₈, Cl₂) to monoatomic (Ar). This explains the change in electrical conductivity, melting point and structure across Period 3 — a common extended-response question.

Common Mistakes

• Confusing atomic number (Z) with mass number (A) — Z is proton count, A is proton + neutron count.
• Assuming atomic mass determines order (it does not — Moseley’s work proved Z does).
• Treating noble gases as completely non-reactive — Xe and Kr do form compounds (e.g., XeF₂) under forcing conditions.
• Forgetting that hydrogen has no fixed group; it can lose 1 electron (like alkali metals) or gain 1 (like halogens).
• Writing “IE always increases across a period” — the trend has the two exceptions above.

Practice Prompts

1. Explain, in terms of atomic structure, why the second ionization energy of sodium is much larger than its first, while the second IE of magnesium is only modestly larger.
2. Element X is in Period 3, Group 16. Element Y is in Period 2, Group 1. Predict and justify which element has (a) the larger atomic radius, (b) the higher electronegativity, (c) the more metallic character.

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