P-Block flash cards

1. 1.Which groups of the periodic table make up the p-block?

   Groups 13 to 18, in which the last electron enters the $p$-orbital. The general valence configuration is $ns^2\,np^{1-6}$.

   Hint: Right side of periodic table.

2. 2.Write the general valence-shell electronic configuration of p-block elements.

   $ns^2\,np^{1-6}$ (where $n=2$ to $6$).

   Hint: s filled, p filling.

3. 3.How does atomic/ionic radius vary down a p-block group and across a period?

   Down a group radius increases (new shells added). Across a period (left to right) radius decreases due to increasing effective nuclear charge.

   Hint: Shells vs Z_eff.

4. 4.Why is there only a small increase in size from Al to Ga in group 13?

   Poor shielding by intervening $3d$ electrons in Ga increases effective nuclear charge, so Ga is nearly the same size as (even slightly smaller than) Al.

   Hint: d-block contraction effect.

5. 5.How does ionization enthalpy generally vary across a period and down a group in the p-block?

   It increases across a period (rising $Z_{eff}$) and decreases down a group (increasing size and shielding).

   Hint: Opposite of atomic radius.

6. 6.What is the inert pair effect?

   The reluctance of the $ns^2$ electron pair to participate in bonding on descending a group, making the lower oxidation state (group valence minus 2) more stable for heavier elements.

   Hint: ns^2 stays inert.

7. 7.Give an example of the inert pair effect in group 14.

   For carbon and silicon $+4$ is stable, but for lead the $+2$ state is more stable than $+4$ (e.g. $\text{PbCl}_2$ is more stable than $\text{PbCl}_4$).

   Hint: Pb prefers +2.

8. 8.What are the common oxidation states shown by group 15 elements?

   $-3$, $+3$ and $+5$. The $+5$ state becomes less stable and $+3$ more stable down the group (inert pair effect); Bi is stable mainly in $+3$.

   Hint: N shows -3 to +5.

9. 9.Why do the first elements of each p-block group show anomalous behaviour?

   Because of their small size, high electronegativity, high charge/radius ratio, and absence of $d$-orbitals in the valence shell.

   Hint: No d-orbitals, small size.

10. 10.How does the absence of d-orbitals affect the covalency of first-row p-block elements?

    They cannot expand their octet, so their maximum covalency is 4 (e.g. N, O, F), whereas heavier congeners can show higher covalency using $d$-orbitals (e.g. $\text{PCl}_5$, $\text{SF}_6$).

    Hint: N can't form NCl5.

11. 11.Why does nitrogen form $p\pi$-$p\pi$ multiple bonds while heavier group 15 elements do not?

    Small N atoms allow effective sideways $2p$-$2p$ overlap giving strong $\pi$ bonds ($\text{N}\equiv\text{N}$); larger atoms (P, As) have poor $p\pi$-$p\pi$ overlap and prefer single bonds/catenation.

    Hint: Size and orbital overlap.

12. 12.QUESTION: Why is nitrogen gaseous and diatomic ($\text{N}_2$) while phosphorus is a solid ($\text{P}_4$)?

    N forms a strong triple bond $\text{N}\equiv\text{N}$ giving small, stable $\text{N}_2$ molecules with only weak van der Waals forces between them (gas). P cannot form stable $p\pi$-$p\pi$ bonds, so it forms single-bonded $\text{P}_4$ tetrahedra, a solid.

    Hint: pπ-pπ bonding difference.

13. 13.Where is boron placed and what is its most common oxidation state?

    Boron is the first element of group 13; it is a non-metal/metalloid and shows the $+3$ oxidation state, forming only covalent compounds.

    Hint: Group 13 head, +3.

14. 14.Why is boron always covalent while other group 13 elements can be ionic?

    The sum of its first three ionization enthalpies is very high, so forming $\text{B}^{3+}$ is energetically impossible; boron shares electrons (covalent) instead.

    Hint: Too high IE for B3+.

15. 15.Write the formula and structure of borax.

    Borax is sodium tetraborate decahydrate, $\text{Na}_2\text{B}_4\text{O}_7\cdot 10\text{H}_2\text{O}$; it actually contains the ion $[\text{B}_4\text{O}_5(\text{OH})_4]^{2-}$ with two 3-coordinate and two 4-coordinate boron atoms.

    Hint: Tetraborate decahydrate.

16. 16.Why is an aqueous solution of borax alkaline (basic)?

    Borax hydrolyses to give boric acid (weak acid) and NaOH (strong base): $\text{Na}_2\text{B}_4\text{O}_7 + 7\text{H}_2\text{O} \rightarrow 4\text{H}_3\text{BO}_3 + 2\text{NaOH}$.

    Hint: Salt of weak acid + strong base.

17. 17.What is the borax bead test used for?

    Heated borax gives a transparent glassy bead of $\text{NaBO}_2$ and $\text{B}_2\text{O}_3$; with coloured metal salts it forms characteristically coloured metaborates, used to identify transition metals.

    Hint: Coloured metaborate beads.

18. 18.Is boric acid $\text{H}_3\text{BO}_3$ a protic acid? Explain.

    No. It is a weak monobasic Lewis acid: it does not donate $\text{H}^+$ but accepts $\text{OH}^-$ from water, $\text{B(OH)}_3 + \text{H}_2\text{O} \rightleftharpoons [\text{B(OH)}_4]^- + \text{H}^+$.

    Hint: Accepts OH-, not H+ donor.

19. 19.Describe the structure of solid boric acid.

    It has a layered structure in which planar $\text{BO}_3$ units are joined by hydrogen bonds; the layers are held by weak van der Waals forces, giving it a soapy/slippery feel.

    Hint: H-bonded 2D sheets.

20. 20.What is the formula of diborane and how many electrons does it have for bonding?

    Diborane is $\text{B}_2\text{H}_6$. It is electron-deficient: it has only 12 valence electrons but appears to need more for eight normal 2-centre-2-electron bonds.

    Hint: Electron-deficient hydride.

21. 21.Describe the bonding and structure of diborane.

    Four terminal B–H bonds are normal 2c-2e bonds; the two bridging H atoms form banana (3-centre-2-electron) B–H–B bonds. The four terminal H and two B are coplanar; bridging H lie above and below.

    Hint: Two 3c-2e bridge bonds.

22. 22.QUESTION: How many bridging and terminal hydrogen atoms are present in $\text{B}_2\text{H}_6$?

    There are 2 bridging (B–H–B) and 4 terminal hydrogen atoms.

    Hint: 6 total = 2 + 4.

23. 23.Why does aluminium exhibit a strong tendency to form complexes like $[\text{AlF}_6]^{3-}$?

    Because of its small size, high charge, and availability of vacant $d$-orbitals, Al can expand its coordination number to 6, unlike boron (max 4).

    Hint: Vacant d-orbitals, small ion.

24. 24.What does 'amphoteric' mean, illustrated by aluminium oxide?

    $\text{Al}_2\text{O}_3$ reacts with both acids and bases: with acid it gives $\text{Al}^{3+}$ salts, and with alkali it gives aluminates $[\text{Al(OH)}_4]^-$.

    Hint: Reacts both ways.

25. 25.Why does aluminium resist corrosion despite being reactive?

    A thin, tough, adherent oxide film ($\text{Al}_2\text{O}_3$) forms on the surface and protects the metal beneath from further oxidation (passivation).

    Hint: Protective oxide layer.

26. 26.Name the four important allotropes of carbon.

    Diamond, graphite, and fullerenes are the main allotropes (with amorphous forms like coke/charcoal). Fullerenes such as $\text{C}_{60}$ are the only pure/crystalline molecular allotrope.

    Hint: Diamond, graphite, fullerene.

27. 27.Compare the hybridisation of carbon in diamond and graphite.

    In diamond each carbon is $sp^3$ hybridised (3-D tetrahedral network). In graphite each carbon is $sp^2$ hybridised, forming planar hexagonal layers with delocalised $\pi$ electrons.

    Hint: sp3 vs sp2.

28. 28.Why is graphite a good conductor and lubricant but diamond is not?

    Graphite's delocalised $\pi$ electrons in each layer conduct electricity, and layers slide over one another (weak forces) giving lubrication. Diamond's electrons are all localised in rigid $sp^3$ bonds, so it is a hard insulator.

    Hint: Delocalised electrons + sliding layers.

29. 29.Describe the structure of fullerene $\text{C}_{60}$.

    $\text{C}_{60}$ (Buckminsterfullerene) is a football-shaped cage of 20 six-membered and 12 five-membered rings; each carbon is $sp^2$ hybridised.

    Hint: Buckyball, 20 hexagons + 12 pentagons.

30. 30.How are the two oxides of carbon, CO and $\text{CO}_2$, classified?

    CO is a neutral oxide and $\text{CO}_2$ is a weakly acidic oxide (it forms carbonic acid with water).

    Hint: Neutral vs acidic.

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