Chemistry p-Block Elements Study Flashcards

Group 15 Elements: The Pnicogens

• Occurrence and Sources

  • Nitrogen makes up $78\%$ of the atmosphere by volume.
  • In the earth's crust, it exists as sodium nitrate ($NaNO_3$, Chile saltpetre) and potassium nitrate ($KNO_3$, Indian saltpetre).
  • Phosphorus occurs in minerals of the apatite family, specifically $Ca_9(PO_4)_6 \cdot CaX_2$, where $X = F, Cl, \text{ or } OH$. An example is fluorapatite: $Ca_9(PO_4)_6 \cdot CaF_2$.
  • Arsenic ($As$), Antimony ($Sb$), and Bismuth ($Bi$) are primarily found as sulphide minerals.

• Atomic and Physical Properties

  • Electronic Configuration: The general valence shell configuration is $ns^2 np^3$.
  • Atomic and Ionic Radii: Radii increase down the group from $N$ to $Bi$ due to the addition of extra principal shells.
  • Ionisation Enthalpy: Decreases down the group as atomic size increases. Group 15 elements have higher ionisation enthalpies than Group 14 elements in the same period because of their smaller size and extra stable half-filled $p$ orbitals ($np^3$). The order of successive enthalpies is     $\Delta H_1 < \Delta H_2 < \Delta H_3$
  • Electronegativity: Decreases down the group from $N$ to $Bi$ as atomic size increases and nuclear attraction for electrons weakens.
  • Metallic Character: Increases down the group. Nitrogen and phosphorus are non-metals; arsenic and antimony are metalloids; bismuth is a metal.
  • States of Matter: Nitrogen is a diatomic gas ($N_2$) with a triple bond, short bond length, and high bond enthalpy ($944\,kJ\,mol^{-1}$). It condenses to a liquid at $77\,K$. Other elements are solids.
  • Boiling and Melting Points: Boiling points generally increase from top to bottom. Melting points increase up to arsenic and then decrease to bismuth.

• Chemical Properties and Oxidation States

  • Common oxidation states are $-3$, $+3$, and $+5$.
  • The stability of the $+3$ state increases down the group, while the stability of the $+5$ state decreases (inert pair effect). The only well-characterised $Bi(V)$ compound is $BiF_5$.
  • The tendency to show $-3$ oxidation state decreases down the group due to increased size and metallic character.
  • Nitrogen Specifics: At room temperature, nitrogen reacts only with lithium to form the nitride $Li_3N$. In acid solution, oxidation states from $+1$ to $+4$ tend to disproportionate:     $3HNO_2 \rightarrow HNO_3 + H_2O + 2NO$
  • Phosphorus Specifics: More reactive than nitrogen. Reacts with metals to form phosphides and ignites in air to form oxides. Intermediate oxidation states disproportionate into $+5$ and $-3$ in both acid and alkali.

Allotropic Forms and Anomalous Behavior of Nitrogen

• Anomalous Properties of Nitrogen

  • Nitrogen is a gas at ordinary temperatures; others are solids.
  • Nitrogen forms diatomic molecules ($N \equiv N$) using $p\pi - p\pi$ multiple bonds. Others form tetra-atomic molecules ($P_4, As_4, Sb_4$) with single bonds.
  • The heat of dissociation for $N_2$ is $225.8\,kcal\,mol^{-1}$, making it extremely unreactive.
  • It is the only member capable of forming hydrogen bonds due to high electronegativity.
  • It cannot expand its octet; its maximum covalency is limited to $4$ (e.g., $NH_4^+$) because it lacks $d$-orbitals.
  • Nitrogen forms the ionic nitride ion ($N^{3-}$), whereas phosphides and arsenides are not typically ionic.

• Phosphorus Allotropes

  • White Phosphorus: Translucent white waxy solid. Highly reactive with a low ignition temperature ($30^\circ C$). Stored under water. Exists as tetrahedral $P_4$ molecules with $60^\circ$ bond angles.
  • Red Phosphorus: Prepared by heating white phosphorus at $250^\circ C$ in the absence of air. Polymeric structure consisting of chains of $P_4$ tetrahedra. It is non-poisonous, stable, and insoluble in $CS_2$.
  • Black Phosphorus: Has two forms. $\alpha$-black is formed by heating red phosphorus at $803\,K$ in a sealed tube (opaque monoclinic/rhombohedral crystals). $\beta$-black is formed by heating white phosphorus at $473\,K$ under high pressure. It is a good conductor of electricity.

Dinitrogen and Ammonia

• Dinitrogen ($N_2$) Preparation

  • Laboratory: Heating aqueous $NH_4Cl$ and $NaNO_2$:     $NH_4Cl(aq) + NaNO_2(aq) \rightarrow N_2(g) + 2H_2O(l) + NaCl(aq)$
  • Thermal Decomposition: Heating ammonium dichromate:     $(NH_4)_2Cr_2O_7 \rightarrow N_2 + 4H_2O + Cr_2O_3$
  • Very Pure Nitrogen: Decomposition of sodium or barium azides:     $Ba(N_3)_2 \rightarrow Ba + 3N_2$
  • Commercial: Fractional distillation of liquid air. $N_2$ (b.p. $77.2\,K$) distills before $O_2$ (b.p. $90\,K$).

• Ammonia ($NH_3$) Manufacture: Haber's Process

  • Reaction: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$
  • Enthalpy: $\Delta H = -46.1\,kJ\,mol^{-1}$
  • Optimum Conditions: Pressure of $200 \times 10^5\,Pa$ ($200\,atm$), Temperature of $700\,K$, and catalysts (Iron oxide with small amounts of $K_2O$ and $Al_2O_3$).
  • Structure: $sp^3$ hybridised nitrogen, pyramidal shape, $H-N-H$ angle of $107.5^\circ$, and bond length of $101.7\,pm$.

• Ammonia Properties

  • Basicity: Forms $OH^-$ in water: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$.
  • Reaction with salts: Precipitates metal hydroxides:     $2FeCl_3(aq) + 3NH_4OH(aq) \rightarrow Fe_2O_3 \cdot xH_2O(s) + 3NH_4Cl(aq)$ (brown precipitate).
  • Lewis Base: The lone pair on $N$ allows it to form complexes with metals like $Cu^{2+}$ (deep blue complex $[Cu(NH_3)_4]^{2+}$) and $Ag^+$ (soluble complex $[Ag(NH_3)_2]Cl$).

Oxides and Oxoacids of Nitrogen

• Nitrogen Oxides Data

  • Nitrous Oxide ($N_2O$): $+1$, colourless neutral gas. Prepared by heating $NH_4NO_3$.
  • Nitric Oxide ($NO$): $+2$, colourless neutral gas. Prepared by action of dil. $HNO_3$ on $Cu$ or catalytic oxidation of $NH_3$.
  • Dinitrogen Trioxide ($N_2O_3$): $+3$, blue acidic solid. Prepared from $NO$ and $NO_2$ at $250\,K$.
  • Nitrogen Dioxide ($NO_2$): $+4$, brown acidic gas. Prepared by heating heavy metal nitrates like $Pb(NO_3)_2$.
  • Dinitrogen Tetraoxide ($N_2O_4$): $+4$, colourless acidic liquid/solid. Dimer of $NO_2$.
  • Dinitrogen Pentoxide ($N_2O_5$): $+5$, colourless acidic solid. Prepared by dehydrating $HNO_3$ with $P_4O_{10}$. Structure exists as planar units.

• Nitric Acid ($HNO_3$)

  • Ostwald's Process (Commercial):
    1. Catalytic oxidation of $NH_3$: $4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O$ ($Pt/Rh$ catalyst, $500\,K$, $9\,bar$).
    2. Oxidation of $NO$: $2NO + O_2 \rightleftharpoons 2NO_2$.
    3. Hydration: $3NO_2 + H_2O \rightarrow 2HNO_3 + NO$.
  • Oxidising Properties: Concentrated $HNO_3$ oxidises non-metals to their highest states (e.g., $I_2 \rightarrow HIO_3$, $P_4 \rightarrow H_3PO_4$, $S \rightarrow H_2SO_4$, $C \rightarrow CO_2$).
  • Reaction with Metals:
    • Copper: With dilute $HNO_3 \rightarrow NO$; with conc. $HNO_3 \rightarrow NO_2$.
    • Tin: With dilute $HNO_3 \rightarrow NH_4NO_4$; with conc. $HNO_3 \rightarrow H_2SnO_3$ (metastannic acid).
    • Iron: Passivated by conc. $HNO_3$ due to oxide film formation.
    • Magnesium/Manganese: React with very dilute $HNO_3$ to liberate $H_2$ gas.
  • Aqua-Regia: Mixture of $1$ part conc. $HNO_3$ and $3$ parts conc. $HCl$. Dissolves gold and platinum by forming nascent chlorine.

Phosphorus Compounds: Phosphine and Halides

• Phosphine ($PH_3$)

  • Preparation: Hydrolysis of $Ca_3P_2$ or reaction of white $P_4$ with boiling $NaOH$:     $P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2$
  • Properties: Colorless, rotten fish smell, highly poisonous. Spontaneously flammable if $P_2H_4$ traces are present. Used in Holme's signals.

• Phosphorus Trichloride ($PCl_3$)

  • Structure: $sp^3$ hybridised, pyramidal shape.
  • Properties: Colorless oily liquid. Fumes in moisture:     $PCl_3 + 3H_2O \rightarrow H_3PO_3 + 3HCl$

• Phosphorus Pentachloride ($PCl_5$)

  • Structure: Liquid/Gas phase: Trigonal bipyramidal. Axial bonds ($240\,pm$) are longer than equatorial bonds ($202\,pm$) due to repulsion. Solid phase: Ionic $[PCl_4]^+[PCl_6]^-$ ($Tetrahedral/Octahedral$).
  • Preparation: White $P_4$ with excess dry $Cl_2$.
  • Properties: Yellowish white powder. Decomposes on heating: $PCl_5 \rightarrow PCl_3 + Cl_2$.

• Oxoacids of Phosphorus

  • Hypophosphorous ($H_3PO_2$): $+1$, monobasic. Strong reducing agent (contains two $P-H$ bonds).
  • Phosphorous ($H_3PO_3$): $+3$, dibasic (two $P-OH$ bonds). Disproportionates on heating:     $4H_3PO_3 \rightarrow 3H_3PO_4 + PH_3$
  • Orthophosphoric ($H_3PO_4$): $+5$, tribasic.
  • Pyrophosphoric ($H_4P_2O_7$): $+5$, tetrabasic.

Group 16 Elements: The Chalcogens

• General Properties

  • Elements: Oxygen ($O$), Sulphur ($S$), Selenium ($Se$), Tellurium ($Te$), Polonium ($Po$).
  • Electronic Configuration: $ns^2 np^4$.
  • Trends: Atomic radii and metallic character increase down the group. Ionisation enthalpy decreases down the group. Group 16 ionisation enthalpies are slightly lower than Group 15 due to the extra stability of half-filled configurations in Group 15.
  • Electron Gain Enthalpy: High negative values. Oxygen has a less negative value than sulphur due to its small size and interelectronic repulsions in the $2p$ shell.
  • Electronegativity: Oxygen is the second most electronegative element after fluorine.
  • Allotropy: All group 16 elements show allotropy. Oxygen exists as $O_2$ and $O_3$. Sulphur exists as Rhombic ($\alpha$) and Monoclinic ($\beta$).

• Hydrides ($H_2E$)

  • Stability: Decreases from $H_2O$ to $H_2Te$.
  • Acidity: Increases down the group ($H_2O < H_2S < H_2Se < H_2Te$) as bond strength decreases.
  • Boiling Point: $H_2O$ is anomalously high due to hydrogen bonding. Trend: $H_2S < H_2Se < H_2Te < H_2O$.

Ozone ($O_3$) and Oxides of Sulphur

• Ozone

  • Preparation: Silent electric discharge through dry $O_2$:     $3O_2 \rightarrow 2O_3 \quad \Delta H = +142\,kJ\,mol^{-1}$
  • Structure: Angular, bond angle $117^\circ$, bond length $128\,pm$ (resonance hybrid).
  • Properties: Powerful oxidising agent. Thermodynamically unstable ($\Delta G$ is negative).

• Sulphur Allotropes

  • Rhombic ($\alpha$-sulphur): Stable at room temperature. Yellow, octahedral crystals, m.p. $112.8^\circ C$, soluble in $CS_2$.
  • Monoclinic ($\beta$-sulphur): Stable above $369\,K$ (transition temperature). Needle-shaped crystals.
  • Structure: Exists as $S_8$ crown-shaped molecules. In $cyclo-S_6$, it adopts a chair form.

• Sulphur Dioxide ($SO_2$)

  • Properties: Colorless pungent gas, highly soluble in water forming sulphurous acid ($H_2SO_3$). Acts as a reducing agent (moist $SO_2$) and a bleaching agent (temporary reduction).
  • Structure: Angular ($119.5^\circ$), $sp^2$ hybridised.

• Sulphuric Acid ($H_2SO_4$): King of Chemicals

  • Contact Process:
    1. Produce $SO_2$ by burning sulphur/ores.
    2. Catalytic oxidation: $2SO_2 + O_2 \rightleftharpoons 2SO_3$ (Catalyst: $V_2O_5$).
    3. Absorption: $SO_3 + H_2SO_4 \rightarrow H_2S_2O_7$ (Oleum).
    4. Dilution: $H_2S_2O_7 + H_2O \rightarrow 2H_2SO_4$.
  • Properties: Strong dibasic acid, high dehydrating agent (chars sugar), and moderately strong oxidising agent.

Group 17 Elements: The Halogens

• Physical and Atomic Trends

  • Electronic Configuration: $ns^2 np^5$.
  • Radii: Smallest atomic radii in their respective periods.
  • Electron Gain Enthalpy: Chlorine has a more negative value ($-349\,kJ\,mol^{-1}$) than fluorine ($-333\,kJ\,mol^{-1}$) because of interelectronic repulsions in fluorine’s compact $2p$ orbitals.
  • Bond Dissociation Enthalpy: Order: $Cl_2 > Br_2 > F_2 > I_2$. $F_2$ is lower than $Cl_2$ due to lone pair-lone pair repulsions.
  • Colors: Fluorine (Pale yellow), Chlorine (Greenish yellow), Bromine (Reddish brown liquid), Iodine (Deep violet solid).

• Chemical Properties

  • Oxidation States: $-1$ (all). $Cl, Br, I$ also show $+1, +3, +5, +7$. Fluorine shows only $-1$.
  • Reactivity: Strong oxidising agents. $F_2$ is the strongest.
  • Interhalogen Compounds: Formed as $XX', XX'_3, XX'_5, XX'_7$:
    • More reactive than constituent halogens (except $F_2$) because $X-X'$ bond is weaker.
    • $IF_7$ is pentagonal bipyramidal.
    • $ClF_3$ is bent T-shaped.

Group 18 Elements: The Noble Gases

• General Features

  • Elements: Helium ($He$), Neon ($Ne$), Argon ($Ar$), Krypton ($Kr$), Xenon ($Xe$), Radon ($Rn$).
  • Configuration: $ns^2 np^6$ ($He: 1s^2$).
  • Properties: Monoatomic, very high ionisation enthalpy, almost zero/positive electron gain enthalpy. Very low b.p. and m.p. ($He$ b.p. is $4.2\,K$).

• Xenon Chemistry

  • Neil Bartlett prepared the first compound $Xe^+PtF_6^-$ after noticing the similar ionisation enthalpies of $Xe$ and $O_2$.
  • Fluorides:
    • $XeF_2$: Linear, $3$ lone pairs, $2$ bond pairs.
    • $XeF_4$: Square planar, $2$ lone pairs, $4$ bond pairs.
    • $XeF_6$: Distorted octahedral, $1$ lone pair, $6$ bond pairs.
  • Oxides:
    • $XeO_3$: Pyramidal, formed by hydrolysis of $XeF_4$ or $XeF_6$.
    • $XeOF_4$: Square pyramidal.

• Uses

  • Helium: Filling balloons, cryogenics, and deep-sea diving (diluent for oxygen).
  • Neon: Discharge tubes and advertising signs.
  • Argon: Providing inert atmosphere in metallurgy and filling incandescent bulbs.
  • Radon: Radiotherapy for cancer.