Group 1 Elements: Alkali Metals
Group 1 elements, also known as alkali metals, are found on the far left side of the periodic table. This group includes lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). Their electron configurations end in s¹, while Group 2 elements end in s². Although hydrogen (H) is often mentioned in connection with them, it doesn’t count as an alkali metal because it behaves more like a non-metal.
Oxidation States and Physical Properties
Oxidation States
All Group 1 elements (Li, Na, K, Rb, Cs, Fr) exhibit a +1 oxidation state exclusively in their compounds. This occurs because:
- They have a single valence electron (ns1 configuration)
- Losing this electron achieves a stable noble gas configuration
- The second ionization energy is extremely high (removing an electron from the core)
Trends in Physical Properties
Key trends observed down Group 1:
1. Atomic Radius
- Increases down the group (Li → Cs)
- Due to the addition of electron shells
- Li: 152 pm → Cs: 262 pm
Figure: Atomic radii trend in Group 1 elements
2. Ionization Energy
- Decreases down the group
- The outer electron is farther from the nucleus and better shielded
- Li: 520 kJ/mol → Cs: 376 kJ/mol
Figure: First ionization energy trend
3. Electronegativity
- Decreases down the group
- Atoms become less able to attract bonding electrons
- Li: 1.0 → Cs: 0.8 (Pauling scale)
This graph shows the electronegativity of Group 1 elements, their tendency to attract electrons. They all have very low values, meaning they’re not great at pulling electrons from other atoms. Additionally, electronegativity decreases as you move down the group.
4. Melting and Boiling Points
- Decrease down the group
- Due to weaker metallic bonds as the atomic size increases
- Li: 180.5°C → Cs: 28.7°C (melting points)
Figure: Melting and boiling point trends
| Element | Atomic Radius (pm) | 1st IE (kJ/mol) | Electronegativity | Melting Point (°C) |
|---|---|---|---|---|
| Lithium (Li) | 152 | 520 | 1.0 | 180.5 |
| Sodium (Na) | 186 | 496 | 0.9 | 97.8 |
| Potassium (K) | 227 | 419 | 0.8 | 63.7 |
| Rubidium (Rb) | 248 | 403 | 0.8 | 38.9 |
| Caesium (Cs) | 262 | 376 | 0.7 | 28.7 |
Chemical Reactions of Group 1 Elements
1. Reactions with Water
All alkali metals react vigorously with water, producing metal hydroxides and hydrogen gas:
2M + 2H2O → 2MOH + H2
- Reactivity increases down the group (Li → Cs)
- Lithium reacts steadily, sodium melts and moves on the water surface, potassium ignites, rubidium and caesium explode
- The reactions are exothermic, with enthalpy becoming more negative down the group
2. Reactions with Oxygen
Alkali metals form different types of oxides depending on the element:
- Lithium: Forms normal oxide (Li2O)
- 4Li + O2 → 2Li2O
- Sodium: Forms peroxide (Na2O2)
- 2Na + O2 → Na2O2
- Potassium, Rubidium, Caesium: Form superoxides (MO2)
- M + O2 → MO2
3. Reactions with Chlorine
All alkali metals react vigorously with chlorine gas to form ionic chlorides:
2M + Cl2 → 2MCl
- Reactions become more vigorous down the group
- All produce white crystalline salts (except CsCl, which has a different structure)
- Chlorides are all soluble in water and conduct electricity when molten
Solubility Trends of Compounds
1. Hydroxides (MOH)
- All are soluble in water
- Solubility increases down the group (LiOH is least soluble, CsOH is most soluble)
- All form strongly alkaline solutions
2. Sulfates (M2SO4)
- Solubility decreases down the group
- Li2SO4 is very soluble, Cs2SO4 is least soluble
- Due to the increasing lattice energy with larger cations
3. Carbonates (M2CO3)
- All are soluble in water
- Solubility increases down the group
- Li2CO3 has relatively low solubility (unlike other Group 1 carbonates)
| Compound | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Hydroxide | 12.7 | 109 | 112 | 180 | 197 |
| Sulfate | 34.9 | 19.5 | 11.1 | 6.2 | 2.4 |
| Carbonate | 1.3 | 21.5 | 112 | 450 | Very soluble |
Thermal Stability of Compounds
1. Nitrates (MNO3)
- Decompose to nitrites and oxygen (except LiNO3)
- 2MNO3 → 2MNO2 + O2
- Lithium nitrate decomposes differently (similar to Group 2):
- 4LiNO3 → 2Li2O + 4NO2 + O2
- Stability increases down the group (LiNO3 least stable)
2. Carbonates (M2CO3)
- Generally stable to heat (except Li2CO3)
- Lithium carbonate decomposes:
- Li2CO3 → Li2O + CO2
- Stability increases down the group due to the decreasing polarizing power of cations
Key Concept: Polarizing Power
Smaller cations (like Li+) have higher charge density and can distort the electron cloud of anions (like CO32-), making the compound less stable to heat. Larger cations (like Cs+) have lower polarizing power, resulting in more stable compounds.
| Carbonate | Formula | Decomposition Temperature (°C) | Decomposition Products | Notes |
|---|---|---|---|---|
| Lithium carbonate | Li2CO3 | ~700 | Li2O + CO2 | Only Group 1 carbonate that decomposes readily |
| Sodium carbonate | Na2CO3 | ~851 (melts) | No decomposition | Stable up to the melting point |
| Potassium carbonate | K2CO3 | ~891 (melts) | No decomposition | Extremely thermally stable |
| Rubidium carbonate | Rb2CO3 | ~900 (melts) | No decomposition | Most stable Group 1 carbonate |
| Caesium carbonate | Cs2CO3 | ~610 (melts) | No decomposition | Lower melting point but still thermally stable |
Key Observations:
- Lithium carbonate is the only Group 1 carbonate that decomposes at reasonable temperatures (similar to Group 2 carbonates)
- All other Group 1 carbonates melt without decomposition
- Thermal stability increases down the group (Li2CO3 least stable, Rb2CO3 most stable)
- This trend is due to the decreasing polarizing power of the cations (Li+ can distort CO32- more than larger ions)
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