Predict the properties of other elements in Group I, given information about the elements

Cambridge IGCSE Chemistry 0620

Topic: Predicting the Properties of Group I (Alkali‑metal) Elements

Learning objective

By the end of this lesson you will be able to predict the physical and chemical properties of any element in Group I by using the trends shown by the known alkali metals (Li, Na, K, Rb, Cs, Fr) and the concepts required in syllabus section 8.2.

1. What is Group I?

• Group I (the alkali metals) occupy the first column of the periodic table.
• All have a single valence electron: ns¹ where n = period number.
• Because they lose only one electron, they almost exclusively form a +1 oxidation state in their compounds.
• Typical compounds:
  • Halides – ionic salts of the form M⁺X⁻ (e.g. NaCl). All alkali‑metal halides are soluble in water.
  • Hydroxides – strong bases of the form M⁺OH⁻ (e.g. KOH). All are highly soluble and give alkaline solutions.

2. Key physical & chemical data for the known alkali metals

3. Down‑group trends required by the syllabus (8.2)

4. Why these trends occur (link to electronic structure)

1. Atomic radius: Adding a new electron shell (higher n) enlarges the atom.
2. Shielding & effective nuclear charge: Inner‑shell electrons shield the nucleus; as the number of shells grows, Z_eff on the valence electron falls, lowering ionisation energy.
3. Reactivity: Lower ionisation energy → easier loss of the single valence electron → more vigorous reactions, especially with water.
4. Metallic bonding: Larger cations produce weaker electrostatic attraction between ions in the metal lattice, giving lower melting points.
5. Density: Mass increases faster than volume, so density rises down the group.
6. Electrode potential: The ease of oxidation (loss of e⁻) makes the reduction potential more negative.
7. Thermal conductivity: A larger sea of delocalised electrons in the larger metals enhances heat transfer.

5. Characteristic reactions of Group I metals

5.1 Reaction with water

Overall molecular equation:

\(\displaystyle \text{M(s)} + \text{H}_2\text{O(l)} \;\longrightarrow\; \text{MOH(aq)} + \tfrac12\text{H}_2(g)\)

Ionic (half‑)equations:

• Oxidation (metal): \(\displaystyle \text{M(s)} \;\longrightarrow\; \text{M}^{+}(aq) + e^{-}\)
• Reduction (water): \(\displaystyle 2\text{H}_2\text{O(l)} + 2e^{-} \;\longrightarrow\; \text{H}_2(g) + 2\text{OH}^{-}(aq)\)

The reaction becomes increasingly vigorous down the group (Li < Na < K < Rb ≈ Cs < Fr).

5.2 Reaction with halogens

General molecular equation (formation of ionic halides):

\(\displaystyle 2\text{M(s)} + \text{X}_2(g) \;\longrightarrow\; 2\text{MX(s)}\)

All alkali‑metal halides are white, crystalline, and **soluble in water**.

5.3 Oxidation‑reduction behaviour (standard electrode potentials)

Standard reduction half‑reaction (vs the standard hydrogen electrode, SHE):

\(\displaystyle \text{M}^{+}(aq) + e^{-} \;\longrightarrow\; \text{M(s)}\)

The E° values in the table are for this reduction; the more negative the value, the easier the metal is oxidised.

6. Predicting the properties of an unknown Group I element

When you are told that an element “X” belongs to Group I but its period is not given, follow these steps:

1. Determine the period (n) – the valence‑electron configuration will be ns¹ with n equal to the period number.
2. Atomic radius – larger than the element directly above it, smaller than the element directly below it.
3. Ionisation energy – lower than the element above, higher than the element below.
4. Reactivity with water – more vigorous than the element above, less vigorous than the element below.
5. Melting point & density – melting point will be lower, density higher than the element above.
6. Standard electrode potential – more negative than the element above, less negative than the element below.
7. Thermal conductivity – higher than the element above.

Worked example

“X” is a Group I element in period 6.

• Period 6 → valence configuration 6s¹. The element is **Cesium (Cs)**.
• Atomic radius: larger than rubidium, smaller than francium.
• Ionisation energy: a little lower than Rb (≈ 4.0 kJ mol⁻¹) but higher than Fr (≈ 3.9 kJ mol⁻¹).
• Reactivity with water: more violent than Rb (explosive) but slightly less than Fr (theoretical).
• Melting point: about 28 °C (lower than Rb’s 39 °C).
• Density: ≈ 1.93 g cm⁻³ (higher than Rb’s 1.53 g cm⁻³).
• E° (reduction): ≈ –2.91 V (more negative than Rb’s –2.92 V).
• Thermal conductivity: higher than Rb (≈ 36 W m⁻¹ K⁻¹).

7. Optional extension – Element 119 (Ununennium, Uue)

Uue would be the first element of the eighth period and would sit in Group I.

• Electron configuration: \([Rn]\,5f^{14}\,6d^{10}\,7s^{2}\,8s^{1}\)
• Atomic radius: larger than francium.
• Ionisation energy: expected to be even lower than that of francium, making oxidation extremely easy.
• Reactivity with water: predicted to be violently exothermic, possibly detonating on contact.
• Standard electrode potential: likely more negative than –2.90 V.
• Physical state: soft, silvery metal, similar in appearance to the lighter alkali metals.

8. Practice questions (Cambridge‑style)

1. Arrange the following alkali metals in order of increasing ionisation energy: K, Li, Rb, Na.
2. Which reaction is more exothermic and why?
   
   • Li + H₂O → LiOH + H₂
   • Cs + H₂O → CsOH + H₂
3. An unknown metal “M” reacts with chlorine to give a white solid and has a standard electrode potential of –2.80 V. Identify the most probable group and period of “M”.
4. Predict the melting‑point trend for the series Li → Na → K → Rb → Cs and give a brief explanation.

9. Summary

• All Group I elements have the configuration ns¹ and form a +1 ion.
• Physical trends down the group: atomic radius ↑, density ↑, melting point ↓, thermal conductivity ↑.
• Chemical trends down the group: ionisation energy ↓, standard reduction potential becomes more negative, reactivity (especially with water) ↑.
• Typical compounds (halides, hydroxides) are all soluble; hydroxides are strong bases.
• Understanding the electronic‑structure reasons (shielding, effective nuclear charge) allows you to predict the properties of any known or as‑yet‑undiscovered Group I element.