Identify trends in groups, given information about the elements

Cambridge IGCSE Chemistry 0620 – Periodic Table: Arrangement of Elements

1. Structure of the Periodic Table

• Periods – horizontal rows (1 → 8). The period number equals the number of electron shells occupied by the ground‑state atoms in that row.
• Groups – vertical columns (1 → 18). Elements in the same group have the same number of valence electrons; consequently they show similar chemical behaviour. The group number (for the main‑group elements) predicts the charge of the ion formed (e.g. Group 1 → +1, Group 17 → –1).
• Blocks –
  • s‑block: Groups 1‑2 and He
  • p‑block: Groups 13‑18
  • d‑block: Transition elements, Groups 3‑12
  • f‑block: Lanthanides & Actinides (shown separately)

2. Main‑Group Elements: Group Number, Valence‑Electron Configuration & Typical Ionic Charge

3. Trends Down a Group (Vertical)

4. Trends Across a Period (Horizontal)

5. Representative Groups – Key Features & Example Data

Group 1 – Alkali Metals (Li, Na, K, Rb, Cs, Fr)

• Valence configuration $ns^{1}$ → form $M^{+}$ ions.
• Atomic radius (pm): Li 152, Na 186, K 227, Rb 248, Cs 265.
• First IE (kJ mol⁻¹): Li 520, Na 496, K 419, Rb 403, Cs 376 (decreases down the group).
• Melting point: decreases down the group (Li 180 °C → Cs 28 °C).
• Density: decreases down the group (Li 0.53 g cm⁻³ → Cs 1.93 g cm⁻³).
• Reactivity with water: 2 M + 2 H₂O → 2 MOH + H₂↑. Vigor increases down the group (Li slow, Na moderate, K vigorous, Cs explosive).

Group 2 – Alkaline‑Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

• Valence configuration $ns^{2}$ → form $M^{2+}$ ions.
• Atomic radius (pm): Be 112, Mg 160, Ca 197, Sr 215, Ba 222.
• First IE (kJ mol⁻¹): Be 900, Mg 738, Ca 590, Sr 550, Ba 503.
• Melting point: decreases down the group (Be 1287 °C → Ba 727 °C).
• Density: decreases down the group (Be 1.85 g cm⁻³ → Ba 3.62 g cm⁻³).
• Reactivity with water: lower than Group 1 but increases down the group (Ca slow, Sr moderate, Ba vigorous).

Group 13 – Boron Group (B, Al, Ga, In, Tl)

• Valence $ns^{2}np^{1}$ → typically form $M^{3+}$ ions (B forms covalent compounds).
• Trend: metallic character increases down the group; electronegativity decreases overall (B 2.04, Al 1.61, Ga 1.81, In 1.78, Tl 1.62).
• Melting points and densities fall down the group (e.g., Al 660 °C → Tl 304 °C; Al 2.70 g cm⁻³ → Tl 11.85 g cm⁻³).

Group 14 – Carbon Group (C, Si, Ge, Sn, Pb)

• Valence $ns^{2}np^{2}$ → can exhibit $+4$, $+2$, $-4$ oxidation states.
• Gradual shift from non‑metal (C) to metal (Pb).
• Typical ion/oxidation states: C$^{4-}$, Si$^{4+}$, Ge$^{2+}$, Sn$^{2+}$/Sn$^{4+}$, Pb$^{2+}$/Pb$^{4+}$.

Group 15 – Nitrogen Group (N, P, As, Sb, Bi)

• Valence $ns^{2}np^{3}$ → typical ion $X^{3-}$.
• Electronegativity falls down the group (N 3.04 → Bi 2.02).
• Physical state: gases (N₂) → non‑metals (P) → metalloids (As, Sb) → metal (Bi).

Group 16 – Chalcogens (O, S, Se, Te, Po)

• Valence $ns^{2}np^{4}$ → typical ion $X^{2-}$.
• Reactivity as non‑metals decreases down the group (O most reactive, Po least).
• Physical state: gases (O₂) → solids (S, Se, Te, Po).

Group 17 – Halogens (F, Cl, Br, I, At)

• Valence $ns^{2}np^{5}$ → form $X^{-}$ ions.
• Atomic radius (pm): F 42, Cl 79, Br 94, I 133, At 150 (increases down the group).
• First IE (kJ mol⁻¹): F 1681, Cl 1251, Br 1140, I 1008 (decreases down the group).
• Physical state at 25 °C: gas (F, Cl) → liquid (Br) → solid (I, At).
• Colour: characteristic (F colourless, Cl green‑yellow gas, Br red‑brown liquid, I violet solid).
• Reactivity with metals: halogen + metal → metal halide; reactivity decreases down the group (F > Cl > Br > I).

Group 18 – Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

• Full valence shells; extremely low chemical reactivity.
• Uses: inert atmospheres (Ar), lighting (Ne signs, Xe flash lamps), cryogenics (He).

6. Transition Elements (Groups 3‑12 – d‑block)

• Electron configuration: $ns^{2}(n-1)d^{1-10}$; electrons are added to the inner $(n-1)d$ subshell.
• Key characteristics:
  • High densities and high melting/boiling points (e.g., Fe, Cu, W).
  • Coloured compounds due to d‑d electron transitions.
  • Variable oxidation states (e.g., Fe +2/+3, Cu +1/+2, Mn +2 to +7).
  • Common catalysts (e.g., V₂O₅ in the Contact process, Pt in hydrogenation).
• Trend in metallic character: remains high across the block; ionisation energies are relatively low but do not follow a simple monotonic pattern because of subshell filling.

7. Predicting the Properties of an Unknown Element

Example: An element with atomic number 35 is discovered.

1. Locate Z = 35 on the periodic table → Period 4, Group 17.
2. Group 17 ⇒ valence configuration $ns^{2}np^{5}$, typical ion charge = –1.
3. Period 4 ⇒ four electron shells → atomic radius larger than Cl (Period 3) but smaller than Br (Period 5).
4. Trend‑based predictions:
   • Non‑metal, halogen, diatomic (X₂) at standard temperature.
   • Reactivity with metals: high, but lower than chlorine.
   • Electronegativity ≈ 2.5 (between Cl 3.16 and Br 2.96).
   • Physical state: likely a reddish‑brown solid (as for bromine) that sublimes easily.

These steps mirror the syllabus requirement to predict physical state, typical ion, and relative reactivity from period and group information.

8. Sample Examination‑Style Questions

1. Given that the atomic radius of magnesium (Mg) is 160 pm and that of calcium (Ca) is 197 pm, estimate the atomic radius of strontium (Sr) and explain the trend you used.
2. Explain why sodium (Na) reacts more vigorously with water than lithium (Li), referring to ionisation energy and atomic radius.
3. Chlorine (Cl) and bromine (Br) are both halogens. Which has the higher electronegativity? Justify your answer using the trend down a group.
4. Predict the most likely oxidation state(s) of manganese (Mn) and give one example of a compound for each oxidation state you mention.
5. An unknown element is known to form a –1 ion and belongs to Period 5. Predict its group, likely physical state at room temperature, and one characteristic reaction.

9. Summary of Key Points

• Groups share the same number of valence electrons → similar chemical behaviour.
• Down a group: atomic radius ↑, ionisation energy ↓, electronegativity ↓, metallic character ↑, metal reactivity ↑, non‑metal reactivity ↓.
• Across a period (left → right): atomic radius ↓, ionisation energy ↑, electronegativity ↑, metallic character ↓.
• For main‑group elements the group number predicts the charge of the ion formed (e.g., Group 1 → +1, Group 17 → –1).
• Transition elements have variable oxidation states, high densities, coloured compounds and are important catalysts.
• Using an element’s period and group you can predict its physical state, type of ion, typical reactions and relative reactivity – a skill required in the IGCSE exam.