S-Block Elements flash cards

1. 1.Which elements make up the s-block of the periodic table?

   Group 1 (alkali metals: $Li, Na, K, Rb, Cs, Fr$) and Group 2 (alkaline earth metals: $Be, Mg, Ca, Sr, Ba, Ra$). Their last electron enters the $s$-orbital.

   Hint: First two groups

2. 2.What is the general valence-shell electronic configuration of alkali metals and alkaline earth metals?

   Alkali metals: $ns^1$; alkaline earth metals: $ns^2$.

   Hint: Groups 1 and 2

3. 3.Why are alkali metals the most electropositive elements in their periods?

   They have a single loosely held $ns^1$ electron and low ionisation enthalpy, so they readily lose it to form $M^+$ ions. Electropositive character increases down the group.

   Hint: Ease of losing one electron

4. 4.How does atomic/ionic radius vary down Group 1 and Group 2?

   It increases down each group as new shells are added. For a given period, Group 1 atoms are larger than Group 2 atoms (fewer protons, weaker nuclear pull).

   Hint: More shells downward

5. 5.Compare the atomic radius of an alkali metal with the alkaline earth metal of the same period.

   The alkali metal is larger. E.g. $Na > Mg$, because Mg has one more proton pulling the same shell inward, and $M^{2+}$ formation increases effective nuclear charge.

   Hint: Na vs Mg

6. 6.How does ionisation enthalpy change down Group 1 and across to Group 2?

   IE decreases down each group (size increases). For a period, Group 2 has higher first IE than Group 1, but alkali metals have very high second IE (removing from noble-gas core).

   Hint: Down: decreases

7. 7.Why do alkaline earth metals have higher first ionisation enthalpy than alkali metals but lower second ionisation enthalpy?

   Higher $IE_1$: smaller size, higher nuclear charge. Lower $IE_2$: for Group 2 the second electron is still a valence electron, whereas for Group 1 it comes from the stable noble-gas core.

   Hint: Second IE core vs valence

8. 8.Why are alkali and alkaline earth metals strong reducing agents?

   They have low ionisation enthalpies and highly negative standard reduction potentials, so they readily lose electrons (get oxidised) and reduce other species.

   Hint: Low IE, negative E°

9. 9.Why does Li have an anomalously high (most negative) standard electrode potential despite high ionisation enthalpy?

   Because of its very high hydration enthalpy: the tiny $Li^+$ ion is strongly hydrated, releasing large energy, making $E°(Li^+/Li) = -3.04\,V$, the most negative among alkali metals.

   Hint: Hydration of small Li+

10. 10.How does hydration enthalpy of Group 1 and Group 2 ions vary down the group?

    It decreases down the group as ionic size increases (charge density falls). Small ions like $Li^+$ and $Be^{2+}$ are most heavily hydrated.

    Hint: Smaller ion, more hydration

11. 11.Why are Group 2 ions more heavily hydrated than Group 1 ions?

    Group 2 ions carry a $+2$ charge in a smaller size, giving much higher charge density and stronger ion–dipole attraction with water. So $M^{2+}$ salts are often hydrated.

    Hint: Higher charge, smaller size

12. 12.How does the melting point / hardness trend differ between Group 1 and Group 2 metals?

    Both decrease down the group, but Group 2 metals have higher m.p. and are harder than Group 1, because two valence electrons per atom give stronger metallic bonding.

    Hint: 2 bonding electrons vs 1

13. 13.Why are alkali metals soft and can be cut with a knife?

    They have weak metallic bonding (only one valence electron, large atomic size), giving low cohesive energy, low melting points, and softness.

    Hint: Weak metallic bond

14. 14.How are alkali metals stored and why?

    Under kerosene or paraffin oil (Li in wax/paraffin due to low density), because they react vigorously with air moisture and oxygen. Na and K are highly reactive.

    Hint: Kept away from air/water

15. 15.What flame colours do the alkali metals impart, and why?

    $Li$ crimson-red, $Na$ golden-yellow, $K$ lilac/violet, $Rb$ red-violet, $Cs$ blue. Low IE lets electrons get excited by flame energy and emit visible light on returning.

    Hint: Na = yellow

16. 16.What flame colours do Ca, Sr and Ba give? Why doesn't Be or Mg show one?

    $Ca$ brick-red, $Sr$ crimson, $Ba$ apple-green. Be and Mg have very high IE, so electrons cannot be excited by the flame's low energy — no characteristic colour.

    Hint: Be, Mg colourless flame

17. 17.Why do heavier alkali/alkaline earth metals photo-emit electrons (photoelectric effect)?

    Their outer electrons are very loosely bound (low IE), so even ordinary light can eject them. Cs and K are used in photoelectric cells.

    Hint: Low work function

18. 18.How do alkali metals react with water?

    They react vigorously: $2M + 2H_2O \rightarrow 2MOH + H_2$, liberating $H_2$ and much heat. Reactivity increases down the group; K, Rb, Cs can ignite the hydrogen.

    Hint: Hydroxide + H2 gas

19. 19.Why is the reaction of Li with water less vigorous than that of Na or K, despite Li being the strongest reducing agent in solution?

    Li has high melting point and high lattice/activation energy; the reaction is kinetically slow. Its strong reducing power (negative E°) is a thermodynamic, not a kinetic, property.

    Hint: Thermodynamics vs kinetics

20. 20.What are the three oxide types alkali metals can form with excess oxygen?

    Monoxide $M_2O$ (Li), peroxide $M_2O_2$ (Na, contains $O_2^{2-}$), and superoxide $MO_2$ (K, Rb, Cs, contains $O_2^-$). Larger cations stabilise larger anions.

    Hint: Oxide, peroxide, superoxide

21. 21.Why does Li form only the normal oxide $Li_2O$, Na the peroxide, and K/Rb/Cs superoxides on burning in air?

    The small $Li^+$ stabilises only the small $O^{2-}$. Larger cations ($Na^+, K^+$) stabilise larger $O_2^{2-}/O_2^-$ anions through better lattice energy matching (like sizes pack stably).

    Hint: Cation–anion size match

22. 22.Why does potassium superoxide $KO_2$ find use in space capsules and submarines?

    It supplies $O_2$ and absorbs $CO_2$: $4KO_2 + 2CO_2 \rightarrow 2K_2CO_3 + 3O_2$; and $2KO_2 + 2H_2O \rightarrow 2KOH + H_2O_2 + O_2$.

    Hint: CO2 removal + O2 release

23. 23.How do the alkali metals react with hydrogen, and what is the nature of the product?

    $2M + H_2 \rightarrow 2MH$, forming ionic (saline) hydrides $M^+H^-$. Stability decreases down the group; $LiH$ is the most stable. The $H^-$ ion is the reducing/basic species.

    Hint: Ionic hydrides M+ H-

24. 24.How do Group 2 metals react with hydrogen? Which is an exception?

    $M + H_2 \rightarrow MH_2$ (ionic hydrides for Ca, Sr, Ba). $BeH_2$ and $MgH_2$ are covalent/polymeric; $BeH_2$ is electron-deficient with $Be-H-Be$ 3-centre bonds.

    Hint: BeH2 covalent

25. 25.Why are alkali metal halides highly ionic and water-soluble?

    Large, low-charge $M^+$ cations with low IE readily transfer electrons to halogens, giving high ionic character. High hydration energy of ions favours dissolution.

    Hint: Low charge density cations

26. 26.Which alkali halide is the most covalent, and why?

    $LiI$ — the small $Li^+$ has high polarising power and $I^-$ is large and highly polarisable, so by Fajans' rules covalent character is greatest.

    Hint: Small cation + large soft anion

27. 27.State Fajans' rules governing covalent character in an ionic bond.

    Covalency increases with: (i) small cation, (ii) large anion, (iii) high charge on either ion, and (iv) cation with pseudo/incomplete d-configuration. High polarising power + high polarisability = more covalent.

    Hint: Polarisation of anion

28. 28.Why is lithium the most anomalous alkali metal?

    Its very small size and high charge/radius ratio (high polarising power), high IE and hydration energy, and absence of d-orbitals make it behave differently and resemble Mg (diagonal relationship).

    Hint: Small size, high polarising power

29. 29.Give three ways in which lithium differs from the other alkali metals.

    (i) Forms only $Li_2O$ (no peroxide/superoxide); (ii) $LiOH, Li_2CO_3, LiF, Li_3PO_4$ are less soluble/less stable; (iii) $Li$ reacts with $N_2$ to give $Li_3N$; forms covalent, mostly hydrated salts.

    Hint: Li vs Na family

30. 30.Which alkali metal reacts directly with nitrogen, and what is the product?

    Only lithium: $6Li + N_2 \rightarrow 2Li_3N$ (lithium nitride). This resembles magnesium ($3Mg + N_2 \rightarrow Mg_3N_2$).

    Hint: Li3N, diagonal with Mg

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