Periodic Properties and Group Elements

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your exam.

The Periodic Table: Elements arranged in order of increasing atomic number. In the modern periodic table, elements in the same group have the same number of outer (valence) electrons and similar chemical properties.

Periodic Trends Across Period 3 (Na → Ar):

Property	Trend across Period 3
Atomic radius	Decreases Na → Ar
Ionisation energy	Generally increases (with minor drops at Group 3 and 6)
Electronegativity	Increases Na → Ar
Metallic character	Decreases Na → Si; non-metallic increases Si → Ar
Melting/boiling point	Peaks at Si (macromolecular) then drops

Atomic Radius: Distance from the nucleus to the outermost electron shell. Increases DOWN a group (more shells) and decreases ACROSS a period (increasing nuclear charge pulls electrons closer).

First Ionisation Energy ($IE_1$): Energy required to remove one electron from a neutral gaseous atom. $$Na_{(g)} \rightarrow Na^+_{(g)} + e^- \quad \Delta H = +496 \text{ kJ mol}^{-1}$$

⚡ WAEC Exam Tip: Ionisation energy always has a positive $\Delta H$ (endothermic) because energy is required to overcome electrostatic attraction between the nucleus and the electron. WAEC questions ask for the correct sign — do not write a negative value.

Electronegativity: Measure of the attraction of a bonded atom for the pair of electrons in a covalent bond. Fluorine (EN = 4.0) is the most electronegative element. Increases across a period and decreases down a group.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for students with a few days to months.

Explanation of Periodic Trends:

Atomic Radius: Across a period, protons increase (+1 each step) while electrons are added to the same shell. The increasing nuclear charge pulls electrons inward, decreasing atomic radius. Down a group, electrons are added to new shells (distance increases) and inner electrons shield the nuclear charge, so atomic radius increases.

Ionisation Energy: This generally increases across a period due to decreasing atomic radius and increasing nuclear attraction. The drops at Group 3 (from Al to Ga: $3p^1 \rightarrow 3p^2$, extra electron repulsion in $p$-orbitals reduces ionisation energy) and Group 6 (from N to O: electron-electron repulsion in the doubly-occupied $p$-orbital reduces ionisation energy) are important exceptions WAEC examiners test.

Electronegativity: Increases across a period as effective nuclear charge increases. Decreases down a group because the outer electrons are further from the nucleus and shielded by inner shells.

Group 1 — Alkali Metals (Li to Fr):

• Electronic configuration: $ns^1$ (one electron in outer shell)

• Very low $IE_1$ and electronegativity — most reactive metals

• React with water: $2Na + 2H_2O \rightarrow 2NaOH + H_2$

• React with oxygen: $4Li + O_2 \rightarrow 2Li_2O$ (forms oxide, no peroxide for Li) $$Na + O_2 \rightarrow Na_2O_2 \text{ (peroxide)}$$ $$K + O_2 \rightarrow KO_2 \text{ (superoxide)}$$

• React with chlorine: $2Na + Cl_2 \rightarrow 2NaCl$

Group 7 — Halogens (F to At):

• Electronic configuration: $ns^2np^5$ (seven electrons in outer shell)
• Diatomic molecules: $Cl_2$, $Br_2$, $I_2$ (exist as $Cl–Cl$, $Br–Br$, $I–I$)
• High electronegativity — strong oxidising agents
• React with metals: $2Fe + 3Cl_2 \rightarrow 2FeCl_3$ (produces iron(III) chloride, brownish solution)
• React with hydrogen: $H_2 + Cl_2 \xrightarrow{h\nu} 2HCl$ (UV light required)
• Displacement reactions: $Cl_2 + 2NaBr \rightarrow 2NaCl + Br_2$ (chlorine displaces bromine because it is more electronegative)
• Oxidising power decreases down the group: $F_2 > Cl_2 > Br_2 > I_2$
• WAEC favourite: $Cl_2$ bleaches by oxidation ($Cl_2 + H_2O \rightarrow HCl + HOCl$; $HOCl$ liberates oxygen which bleaches dye)

⚡ WAEC Exam Tip: In displacement reactions, only a more reactive halogen can displace a less reactive one. Fluorine is the most reactive but does not displace any halogen because no element is less reactive. WAEC questions often test: “chlorine does NOT displace fluoride” — state the reason (fluorine is more electronegative and more reactive than chlorine).

Group 0 — Noble Gases (He to Rn):

• Electronic configuration: $ns^2np^6$ (full outer shell — stable)
• Very low reactivity due to complete outer shell
• Uses: Helium in balloons (non-flammable alternative to hydrogen); Argon in light bulbs; Neon in advertising signs

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for students on a longer study timeline.

Ionisation Energy — Detailed Explanation: The general increase in ionisation energy across a period can be explained by:

1. Increasing effective nuclear charge (more protons in nucleus)
2. Decreasing atomic radius (outer electrons closer to nucleus)
3. Same shielding effect (electrons added to same shell)

The specific drops are explained by:

• Group 3 anomaly (Al, Ga): Electron removed from a $p$-orbital which is higher in energy and more shielded than the preceding $s$-electron in the same principal quantum shell
• Group 6 anomaly (N, P): Electron removed from a doubly-occupied $p$-orbital — electron-electron repulsion between the two electrons in the same orbital makes removal easier

Electron Affinity: Energy change when an electron is added to a neutral gaseous atom. $$Cl_{(g)} + e^- \rightarrow Cl^-_{(g)} \quad \Delta H = -349 \text{ kJ mol}^{-1}$$ Most electron affinities are exothermic (negative $\Delta H$). Chlorine has a higher electron affinity than fluorine because fluorine has a very small atomic radius — added electron experiences strong inter-electron repulsion in the compact $2p$-subshell.

Physical Properties of Period 3 Elements:

Element	Type	Structure	Melting Point	Electrical Conductivity
Na	Metal	Giant metallic	371 K	High (delocalised electrons)
Mg	Metal	Giant metallic	924 K	High
Al	Metal	Giant metallic	933 K	Very high
Si	Metalloid	Giant covalent (diamond lattice)	1683 K	Poor solid, moderate liquid
P	Non-metal	Simple molecular ($P_4$)	317 K (white)	Very poor
S	Non-metal	Simple molecular ($S_8$)	388 K	Very poor
Cl	Non-metal	Diatomic ($Cl_2$)	172 K	Poor
Ar	Noble gas	Monatomic	87 K	None

Chemical Properties of Period 3 Oxides:

Oxide	Type	Reaction with Water	Reaction with Acid/Base
$Na_2O$	Basic	$Na_2O + H_2O \rightarrow 2NaOH$	Reacts with acid only
$MgO$	Basic	Slightly soluble	Reacts with acid
$Al_2O_3$	Amphoteric	Insoluble	Reacts with acid AND base ($NaOH$)
$SiO_2$	Acidic	Insoluble	Reacts with $NaOH$ (etching with HF)
$P_4O_{10}$	Acidic	$P_4O_{10} + 6H_2O \rightarrow 4H_3PO_4$	Reacts with base
$SO_2$	Acidic	$SO_2 + H_2O \rightarrow H_2SO_3$	Reacts with base

⚡ WAEC Exam Tip: $Al_2O_3$ is amphoteric — it dissolves in both acids and alkalis. The reaction with $NaOH$: $Al_2O_3 + 2NaOH + 3H_2O \rightarrow 2Na[Al(OH)_4]$. This is a favourite WAEC question for distinguishing amphoteric oxides from purely basic or acidic ones.

Diagonal Relationship — Lithium and Magnesium: Li and Mg show diagonal similarity (properties more similar to each other than to elements in their own groups):

• Both form nitrides: $6Li + N_2 \rightarrow 2Li_3N$; $3Mg + N_2 \rightarrow Mg_3N_2$
• Both form carbonates that decompose on heating: $Li_2CO_3$ and $MgCO_3$ both decompose to oxide + $CO_2$
• Both are less reactive than other Group 1/Group 2 metals

Comparative Study — First Ionisation Energy Across Period 2:

| Element | Li | Be | B | C | N | O | F | Ne | |---------|----|----|---|---|---|美---|----|----| | $IE_1$ (kJ mol⁻¹) | 520 | 899 | 800 | 1086 | 1402 | 1314 | 1681 | 2081 |

Notable: Be (900) > B (800) and N (1402) > O (1314) — these two drops are the most important exceptions to learn.

WAEC Past Question Patterns:

• Explaining the trend in atomic radius across a period (mention nuclear charge and shielding)
• Calculating ionisation energy from successive ionisation energies (identify which electron is being removed at each step)
• Writing equations for reactions of Group 1 and Group 7 elements
• Predicting products of displacement reactions between halogens
• Describing the amphoteric nature of $Al_2O_3$ with equations
• Explaining the variation in melting point across Period 3 using structure and bonding

⚡ WAEC Exam Tip: When asked to explain periodic trends, always give two reasons: (1) the underlying physical cause (e.g., increasing nuclear charge) and (2) why that causes the observed effect (e.g., stronger attraction pulls electrons closer, decreasing radius). One-line answers rarely score full marks in Paper 2 explanations.