Questions & explanations
1. Which of the following orders of hydration enthalpy is correct for the s-block ions?
- Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺
- Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺
- Li⁺ = Na⁺ = K⁺ = Rb⁺ = Cs⁺
- Be²⁺ < Mg²⁺ < Ca²⁺ < Sr²⁺ < Ba²⁺
Answer: Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺
Hydration enthalpy is the energy released when one mole of gaseous ions is hydrated. It depends inversely on ionic size: smaller ions have stronger hydration due to higher charge density. Among alkali metal ions, Li⁺ is the smallest and thus has the highest hydration enthalpy (most negative), while Cs⁺ is the largest and has the lowest. The correct decreasing order is Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺. For alkaline earth metals, Be²⁺ has the highest hydration enthalpy, so the order Ba²⁺ < Sr²⁺ < Ca²⁺ < Mg²⁺ < Be²⁺ would be correct, making option D incorrect.
2. The anomalous behaviour of lithium as compared to other alkali metals is primarily due to its:
- Small size and high polarising power
- Large size and low ionisation enthalpy
- High reactivity and high melting point
- Low ionisation enthalpy and low electronegativity
Answer: Small size and high polarising power
Lithium exhibits anomalous properties due to its exceptionally small atomic and ionic size, which leads to high polarising power. This enhances covalent character in its compounds, making it behave differently from the rest of the alkali metals. High polarising power also explains its ability to form stable nitrides, decompose carbonates and hydroxides upon heating, and show a diagonal relationship with magnesium.
3. Which pair of elements shows a diagonal relationship in the periodic table?
- Lithium and Magnesium
- Sodium and Calcium
- Beryllium and Boron
- Lithium and Beryllium
Answer: Lithium and Magnesium
Diagonal relationships occur between elements of the second period and those of the third period that are diagonally adjacent, due to similar charge/size ratios. The classic examples are Li–Mg and Be–Al. Among the options, only Li–Mg is correct.
4. The first ionization enthalpy of Group 2 elements is generally ________ than that of Group 1 elements because:
- higher; Group 2 atoms have a larger nuclear charge and more tightly held outer electrons
- lower; Group 2 atoms have larger atomic radii
- higher; Group 2 atoms have lower atomic radii
- lower; Group 2 atoms have fully filled s-orbitals
Answer: higher; Group 2 atoms have a larger nuclear charge and more tightly held outer electrons
In the same period, Group 2 elements have a higher nuclear charge and smaller atomic radius than Group 1 elements, making their valence electrons more strongly attracted to the nucleus, so more energy is required to remove an electron.
5. Alkali metals impart characteristic colours to flames because:
- their valence electrons are excited to higher energy levels and emit light on returning to ground state
- they react with oxygen to form coloured oxides
- they sublime easily and their vapours absorb specific wavelengths
- they form peroxides that emit light
Answer: their valence electrons are excited to higher energy levels and emit light on returning to ground state
In a flame, the heat energy excites the outermost s-electron to a higher energy level. When the electron falls back to the ground state, it emits light of a specific wavelength, producing characteristic colours.
6. Compared to Group 2 metals, Group 1 metals are:
- more reactive due to lower ionization enthalpy and larger atomic size
- less reactive due to higher ionization enthalpy
- equally reactive
- more reactive only with water but less with oxygen
Answer: more reactive due to lower ionization enthalpy and larger atomic size
Group 1 metals have lower ionization enthalpies and larger atomic radii than Group 2 metals in the same period, making it easier for them to lose electrons and hence they are more reactive overall.
7. Which of the following correctly describes the trend in atomic radii down Group 1 and Group 2?
- Atomic radii increase down both groups due to increasing number of shells
- Atomic radii decrease down both groups due to increasing nuclear charge
- Atomic radii increase down Group 1 but decrease down Group 2
- Atomic radii remain constant down a group
Answer: Atomic radii increase down both groups due to increasing number of shells
Down a group, the principal quantum number increases, adding new electron shells, which outweighs the effect of increased nuclear charge, so atomic radius increases in both Group 1 and Group 2.
8. Alkali metals are strong reducing agents because:
- they have low ionization enthalpy and can easily lose electrons
- they have high electronegativity
- they have high ionization enthalpy
- they have small atomic radii
Answer: they have low ionization enthalpy and can easily lose electrons
A strong reducing agent readily donates electrons. Alkali metals have low ionization enthalpy, meaning they can lose their valence electron easily, making them powerful reducing agents.
9. Which of the following groups of elements constitute the s-block of the periodic table?
- Group 1 and Group 2
- Group 13 and Group 14
- Group 1 and Group 17
- Group 2 and Group 18
Answer: Group 1 and Group 2
The s-block of the periodic table consists of Group 1 (alkali metals) and Group 2 (alkaline earth metals), because in these elements the last electron enters the s-orbital.
10. The general valence electron configuration of s-block elements is:
- ns¹ for Group 1 and ns² for Group 2
- ns² for Group 1 and ns¹ for Group 2
- ns²np¹ for Group 1 and ns²np² for Group 2
- ns¹ for both groups
Answer: ns¹ for Group 1 and ns² for Group 2
s-block elements have their valence electrons in the s-orbital: Group 1 elements have one s-electron (ns¹) and Group 2 elements have two s-electrons (ns²).