Trends in Alkali Metals
This topic covers the predictable patterns (trends) in the physical and chemical properties of the alkali metals in Group 1 of the periodic table. Understanding these trends allows you to predict the behaviour of unfamiliar alkali metals and explain their reactivity.
Part of the ESAT Chemistry syllabus — revision for the Engineering and Science Admissions Test (ESAT), the UAT-UK admissions test for Cambridge, Imperial, Oxford and UCL.
Key points
- Alkali metals (Li, Na, K) all have one electron in their outer shell, which they readily lose to form a +1 ion.
- Reactivity increases as you go down Group 1. This is because the outer electron is further from the nucleus and experiences more shielding, making it easier to remove (lower first ionisation energy).
- Melting and boiling points decrease down the group. The atoms get larger, so the metallic bonds between the positive ions and delocalised electrons become weaker.
- Density generally increases down the group (with the exception of potassium being slightly less dense than sodium).
- Alkali metals react vigorously with water to produce a metal hydroxide and hydrogen gas. The vigour of this reaction increases down the group, with potassium producing a lilac flame.
Diagram
Formulae
2M(s) + 2H2O(l) → 2MOH(aq) + H2(g) This is the general balanced equation for the reaction of an alkali metal (M) with water. Note the 2:1 mole ratio between the metal and hydrogen gas.
Definitions
- Alkali Metals
- The elements in Group 1 of the periodic table, which are highly reactive metals that form alkaline solutions (hydroxides) when they react with water.
- First Ionisation Energy
- The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous +1 ions.
- Metallic Bonding
- The electrostatic attraction between a lattice of positive metal ions and a 'sea' of delocalised electrons.
Worked example
A 0.92 g sample of pure sodium is completely reacted with excess water. What volume of hydrogen gas, in cm3, is produced at room temperature and pressure (RTP)? (Molar mass of Na = 23 g/mol; Molar volume of a gas at RTP = 24,000 cm3/mol)
- 1
Write the balanced chemical equation:
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g).
- 2
Calculate the moles of sodium used:
moles = mass / molar mass = 0.92 g / 23 g/molThis is 92/23 × 10-2 = 4 × 10-2 = 0.04 mol.
- 3
Use the stoichiometry from the balanced equation.
The ratio of Na to H2 is 2:1.
Therefore, moles of H2 produced = 0.04 mol / 2 = 0.02 mol.
- 4
Calculate the volume of hydrogen gas:
volume = moles × molar volume = 0.02 mol × 24,000 cm3/mol - 5
Simplify the calculation:
2 × 10-2 × 24,000 = 2 × 240 = 480 cm3.
Answer: 480 cm3
Common mistakes
- ×Confusing the reactivity trend of Group 1 (increases down the group) with that of Group 17 (decreases down the group).
- ×Making arithmetic mistakes in stoichiometry calculations, especially when dealing with the 2:1 mole ratio in reactions with water.
- ×Forgetting the conditions of a reaction, such as the lilac flame produced when potassium reacts with water, which can be a required observation.
No-calculator tips
- ✓For mass-to-moles calculations, look for simple multiples. For example, 0.92 g of Na (Mr=23) is (92/23) / 100 = 4/100 = 0.04 moles. Practice recognising multiples of key atomic masses.
- ✓Always write the balanced equation first in stoichiometry questions. The mole ratios are crucial and easily forgotten if you try to do it all in your head.
- ✓Remember that molar volume at RTP is 24 dm3 or 24,000 cm3. Pay close attention to the units requested in the answer to avoid being off by a factor of 1000.