P-Block Elements (halogen & Noble Gases) flash cards

1. 1.Which elements make up Group 17 and what is their group name?

   Group 17 = halogens: F, Cl, Br, I, At (and Ts). Name means salt producers. General electronic configuration $ns^2\,np^5$.

   Hint: 'Salt formers'

2. 2.What is the general valence-shell electronic configuration of halogens, and why are they highly reactive?

   $ns^2\,np^5$ — one electron short of the noble-gas octet, so they have a strong tendency to gain one electron, giving very high reactivity and high electron affinity.

   Hint: One short of octet

3. 3.Give the trend in atomic/ionic radius down Group 17.

   Radius increases down the group ($F < Cl < Br < I$) as new shells are added. Halogens have the smallest radii in their respective periods due to maximum effective nuclear charge.

   Hint: Extra shells win

4. 4.How does ionisation enthalpy vary down Group 17?

   It decreases down the group (F highest, I lowest) because atomic size increases and the outer electron is more shielded/less tightly held. Halogens have very high IE overall.

   Hint: Bigger atom, easier to remove

5. 5.Why does fluorine have a lower electron gain enthalpy (less negative) than chlorine?

   F is very small, so the incoming electron enters a compact $2p$ subshell with strong inter-electronic repulsion. Thus $\Delta_{eg}H$: $Cl > F > Br > I$ (Cl most negative).

   Hint: Small size = repulsion

6. 6.Arrange the halogens in order of electron gain enthalpy (most negative first).

   $Cl > F > Br > I$. Chlorine has the most negative $\Delta_{eg}H$ (least = I).

   Hint: Cl beats F

7. 7.Which is the most electronegative element, and how does electronegativity trend down Group 17?

   Fluorine is the most electronegative element. Electronegativity decreases down the group: $F > Cl > Br > I$.

   Hint: F tops the periodic table

8. 8.Why are halogens coloured?

   They absorb visible light to promote electrons to higher levels; the absorbed wavelength increases (energy gap decreases) down the group. $F_2$ pale yellow, $Cl_2$ greenish-yellow, $Br_2$ red-brown, $I_2$ violet.

   Hint: Absorb visible, show complement

9. 9.State the physical states of the four common halogens at room temperature.

   $F_2$ gas (pale yellow), $Cl_2$ gas (greenish-yellow), $Br_2$ liquid (red-brown), $I_2$ solid (dark violet, sublimes).

   Hint: Gas, gas, liquid, solid

10. 10.What oxidation states does chlorine commonly exhibit? Why can't fluorine show positive states?

    Cl shows $-1, +1, +3, +5, +7$. Fluorine shows only $-1$ (and 0) because it is the most electronegative and has no $d$-orbitals, so it cannot show positive oxidation states.

    Hint: F is only -1

11. 11.Why does only fluorine among halogens lack positive oxidation states and expanded octet?

    F has no low-lying vacant $d$-orbitals and highest electronegativity. Other halogens (Cl, Br, I) have available $d$-orbitals, allowing $+3,+5,+7$ and expanded octets.

    Hint: No d-orbitals in 2nd period

12. 12.Compare bond dissociation enthalpies of $F_2$ and $Cl_2$. Why is $F–F$ anomalously weak?

    $Cl_2 > F_2$ (order overall: $Cl_2 > Br_2 > F_2 > I_2$). $F–F$ is weak due to strong repulsion between non-bonding lone pairs on the two small, close F atoms.

    Hint: Small atoms, lone-pair repulsion

13. 13.Which halogen is the strongest oxidising agent and why (in aqueous solution)?

    Fluorine — highest $E^\circ$. Reasons: low bond dissociation enthalpy of $F_2$, high hydration enthalpy of $F^-$, and high electron gain. Oxidising power: $F_2 > Cl_2 > Br_2 > I_2$.

    Hint: Low bond energy + high hydration

14. 14.QUESTION: A halogen $X_2$ displaces $Br^-$ and $I^-$ from solution but not $Cl^-$. Identify $X_2$.

    $X_2 = Cl_2$. A halogen displaces those below it in oxidising power: $Cl_2$ oxidises $Br^-$ and $I^-$ but cannot displace $Cl^-$ (nothing) — it lies below $F_2$ only.

    Hint: Displaces lower halogens

15. 15.Why does fluorine oxidise water while iodine does not?

    $F_2$ has very high $E^\circ$ and oxidises water: $2F_2 + 2H_2O \to 4HF + O_2$. $I_2$ is a weak oxidiser (low $E^\circ$); instead water/O$_2$ can oxidise $I^-$ is not favourable — iodine is even liberated by many oxidants.

    Hint: F2 strong enough for water

16. 16.Write the reaction of chlorine with cold and hot NaOH.

    Cold, dilute: $Cl_2 + 2NaOH \to NaCl + NaOCl + H_2O$. Hot, conc.: $3Cl_2 + 6NaOH \to 5NaCl + NaClO_3 + 3H_2O$. Both are disproportionation reactions.

    Hint: Disproportionation, temp-dependent

17. 17.What is disproportionation? Give a halogen example.

    A single species is simultaneously oxidised and reduced. Example: $Cl_2$ in NaOH — Cl goes from $0$ to $-1$ (in $Cl^-$) and $+1$ (in $OCl^-$).

    Hint: Same element up and down

18. 18.How is $Cl_2$ prepared in the laboratory from $MnO_2$?

    $MnO_2 + 4HCl \to MnCl_2 + Cl_2 + 2H_2O$. Alternatively $2KMnO_4 + 16HCl \to 2KCl + 2MnCl_2 + 5Cl_2 + 8H_2O$ (no heating needed with $KMnO_4$).

    Hint: Oxidise conc. HCl

19. 19.How is fluorine prepared industrially (why not by chemical oxidation)?

    By electrolysis of anhydrous $KHF_2$ (in liquid HF), since no chemical oxidant is strong enough to oxidise $F^-$ to $F_2$ (F is the strongest oxidiser).

    Hint: Electrolysis of KHF2

20. 20.What is Deacon's process?

    Industrial $Cl_2$ manufacture: $4HCl + O_2 \xrightarrow{CuCl_2,\,723K} 2Cl_2 + 2H_2O$. $CuCl_2$ is the catalyst.

    Hint: HCl + O2, Cu catalyst

21. 21.How does chlorine act as a bleaching agent?

    In presence of moisture it forms nascent oxygen: $Cl_2 + H_2O \to 2HCl + [O]$. Nascent $[O]$ oxidises coloured matter. Bleaching is permanent (oxidative), unlike $SO_2$.

    Hint: Nascent oxygen, oxidative

22. 22.Arrange HX acid strength and explain the trend.

    $HF < HCl < HBr < HI$. Down the group $H–X$ bond enthalpy decreases (bond length increases), so it dissociates more easily — HI is the strongest acid, HF the weakest.

    Hint: Weaker bond = stronger acid

23. 23.Why is HF a weak acid despite F being most electronegative?

    $H–F$ bond is very strong (short bond) and strong H-bonding stabilises undissociated HF. Its high bond dissociation enthalpy makes it dissociate least, so it is the weakest of the hydrohalic acids.

    Hint: Strong short bond

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