P block elements

What is the general valence shell electronic configuration of p-block elements, and what influences their properties?

The general valence shell electronic configuration is $ns^2np^{1-6}$ (except He). Their properties are greatly influenced by atomic sizes, ionisation enthalpy, electron gain enthalpy, and electronegativity. The absence/presence of d/f orbitals also has significant effects.

List the elements of Group 15 and describe the trend in their non-metallic to metallic character.

Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), and Bismuth (Bi). As you go down the group, there is a shift from non-metallic to metallic character, with arsenic and antimony being metalloids and bismuth a typical metal.

Explain why Group 15 elements have extra stable electronic configurations.

The valence shell electronic configuration of these elements is $ns^2np^3$. The s orbital is completely filled, and the p orbitals are exactly half-filled, making their electronic configuration extra stable.

Describe the trends in atomic and ionic radii down Group 15.

Covalent and ionic radii increase down the group. There is a considerable increase from N to P, but only a small increase from As to Bi. This is due to the presence of completely filled d and/or f orbitals in the heavier members.

How does ionisation enthalpy trend down Group 15, and how does it compare to Group 14 elements?

Ionisation enthalpy decreases down the group due to the gradual increase in atomic size. It is much greater than that of Group 14 elements in corresponding periods because of the extra stable half-filled p orbitals electronic configuration and smaller size.

What are the common oxidation states exhibited by Group 15 elements, and how do their stabilities vary down the group?

The common oxidation states are -3, +3, and +5. The tendency to exhibit -3 oxidation state decreases down the group (bismuth hardly forms -3 compounds). The stability of the +5 oxidation state decreases, and that of the +3 state increases down the group (due to the inert pair effect).

Why is nitrogen restricted to a maximum covalency of 4?

Nitrogen is restricted to a maximum covalency of 4 because only four (one s and three p) orbitals are available for bonding in its valence shell. It lacks d-orbitals.

List the anomalous properties of nitrogen and explain the reasons for this behaviour.

Nitrogen differs due to its small size, high electronegativity, high ionisation enthalpy, and non-availability of d orbitals. It forms pπ-pπ multiple bonds with itself (N≡N) and other small, highly electronegative elements (C, O). The N-N single bond is weaker than the P-P bond due to high interelectronic repulsion, leading to weaker catenation. It cannot form dπ-pπ bonds, unlike heavier elements (e.g., $R_3P=O$).

Describe the trends in stability, reducing character, and basicity of Group 15 hydrides ($EH_3$).

Stability: Decreases from $NH_3$ to $BiH_3$ (due to decreasing bond dissociation enthalpy). Reducing Character: Increases from $NH_3$ (mild) to $BiH_3$ (strongest). Basicity: Decreases in the order $NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3$.

Why does nitrogen not form pentahalides like $PX_5$, while phosphorus does?

Nitrogen does not form pentahalides because it lacks d orbitals in its valence shell, thus it cannot expand its covalence beyond four. Phosphorus, being a heavier element, has vacant d orbitals to expand its covalency.

How is pure dinitrogen obtained in the laboratory? Write the chemical equation.

Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide. $Ba(N_3)_2 \rightarrow Ba + 3N_2$. Alternatively: $2NaN_3 \rightarrow 2Na + 3N_2$.

Why is dinitrogen ($N_2$) rather inert at room temperature?

Dinitrogen is inert at room temperature because of the very high bond enthalpy of the N≡N triple bond. Reactivity increases rapidly with rising temperature.

Describe the industrial manufacture of ammonia by Haber's process, including optimum conditions and catalysts.

Ammonia is manufactured on a large scale by Haber's process: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$; $ \Delta_f H^\circ = -46.1 \text{ kJ mol}^{-1}$. Optimum conditions: High pressure (approx. $200 \times 10^5$ Pa or 200 atm), Temperature of ~700 K, Catalyst: Iron oxide with small amounts of $K_2O$ and $Al_2O_3$.

Explain why ammonia ($NH_3$) acts as a Lewis base and mention one application of this property.

The nitrogen atom in $NH_3$ has one lone pair of electrons which is available for donation. Therefore, it acts as a Lewis base. This property allows it to form linkage with metal ions, finding application in the detection of metal ions like $Cu^{2+}$ and $Ag^+$ by forming complex compounds (e.g., $[Cu(NH_3)_4]^{2+}$ (deep blue) from $Cu^{2+}$).

Why does $NO_2$ dimerise to form $N_2O_4$?

$NO_2$ contains an odd number of valence electrons, making it a typical odd molecule. Upon dimerisation, it converts to the stable $N_2O_4$ molecule, which has an even number of electrons.

Describe the industrial preparation of nitric acid ($HNO_3$) by Ostwald's process.

Ostwald's process involves three main steps: 1. Catalytic oxidation of $NH_3$ by atmospheric oxygen: $4NH_3(g) + 5O_2(g) \rightarrow 4NO(g) + 6H_2O(g)$ (Pt/Rh gauge catalyst, 500K, 9 bar). 2. Nitric oxide combines with oxygen to form nitrogen dioxide: $2NO(g) + O_2(g) \rightleftharpoons 2NO_2(g)$. 3. Nitrogen dioxide dissolves in water to give $HNO_3$: $3NO_2(g) + H_2O(l) \rightarrow 2HNO_3(aq) + NO(g)$ (NO is recycled).

How does concentrated nitric acid react with copper and zinc, illustrating its strong oxidising nature?

Concentrated nitric acid is a strong oxidising agent. With Copper: $Cu + 4HNO_3(\text{conc.}) \rightarrow Cu(NO_3)_2 + 2NO_2 + 2H_2O$. With Zinc: $Zn + 4HNO_3(\text{conc.}) \rightarrow Zn(NO_3)_2 + 2H_2O + 2NO_2$.

What is the Brown Ring Test, and what does it detect?

The Brown Ring Test is used to detect the presence of nitrate ions ($NO_3^−$). It relies on $Fe^{2+}$ reducing nitrates to nitric oxide (NO), which then reacts with $Fe^{2+}$ to form a brown-coloured complex, $[Fe(H_2O)_5(NO)]^{2+}$.

Compare the stability and reactivity of white phosphorus and red phosphorus.

White Phosphorus: Less stable and more reactive due to angular strain in the tetrahedral $P_4$ molecule (60° bond angles). It readily catches fire in air and is poisonous. Red Phosphorus: Much less reactive, does not glow in the dark, and is non-poisonous. It is formed by heating white phosphorus at 573K in an inert atmosphere.

Describe the general structure of phosphorus oxoacids and the role of P-H bonds.

In oxoacids, phosphorus is tetrahedrally surrounded by other atoms. All these acids contain at least one P=O bond and one P-OH bond. Acids with phosphorus in lower oxidation states also contain P-P or P-H bonds. P-H bonds impart strong reducing properties (e.g., $H_3PO_2$).

How do you determine the basicity of a phosphorus oxoacid? What are the basicities of $H_3PO_3$ and $H_3PO_4$?

Only hydrogen atoms attached to oxygen in P-OH bonds are ionisable. $H_3PO_3$ (orthophosphorous acid) has two P-OH bonds, so it is dibasic. $H_3PO_4$ (orthophosphoric acid) has three P-OH bonds, so it is tribasic.

List the elements of Group 16 and explain the origin of the name "chalcogens."

Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te), and Polonium (Po). The name "chalcogens" means "ore forming," as many metal ores are oxides or sulfides.

Compare the first ionisation enthalpy values of Group 16 elements with those of Group 15 elements in the corresponding periods. Provide a reason.

Group 16 elements have lower first ionisation enthalpies than Group 15 elements. This is because Group 15 elements have extra stable, half-filled p-orbital electronic configurations ($ns^2np^3$), requiring more energy to remove an electron.

Why does oxygen have a less negative electron gain enthalpy than sulphur?

Due to the exceptionally small size of the oxygen atom, strong interelectronic repulsions in its compact 2p orbitals make it less accommodating for an incoming electron compared to the larger sulphur atom.

Explain why oxygen is a gas, but sulphur is a solid at room temperature.

Oxygen exists as diatomic molecules ($O_2$) with weak intermolecular van der Waals forces. Sulphur exists as larger, polyatomic molecules ($S_8$) with stronger intermolecular forces, leading to its solid state.

Describe the anomalous behaviour of oxygen compared to other Group 16 elements.

Oxygen's behaviour is anomalous due to its small size and high electronegativity. It exhibits strong hydrogen bonding (in $H_2O$), and its covalency is limited as it lacks d-orbitals for valence shell expansion.

How do the acidic character and thermal stability of Group 16 hydrides ($H_2E$) change down the group?

Acidic Character: Increases from $H_2O$ to $H_2Te$ due to a decrease in the H-E bond dissociation enthalpy. Thermal Stability: Decreases from $H_2O$ to $H_2Po$ for the same reason (weaker H-E bonds).

What are the general methods for preparing dioxygen ($O_2$) in the laboratory?

1. Heating oxygen-containing salts like chlorates ($2KClO_3 \rightarrow 2KCl + 3O_2$). 2. Thermal decomposition of oxides of less reactive metals ($2Ag_2O \rightarrow 4Ag + O_2$). 3. Catalytic decomposition of hydrogen peroxide ($2H_2O_2 \rightarrow 2H_2O + O_2$).

What is ozone ($O_3$), and how is it formed from oxygen?

Ozone is an allotropic form of oxygen. In the lab, it's prepared by passing a slow, dry stream of oxygen through a silent electrical discharge, which provides energy to convert $O_2$ to $O_3$.

Why is ozone considered a powerful oxidising agent? Provide an example.

Ozone is a powerful oxidising agent because it easily decomposes to produce nascent oxygen ($O_3 \rightarrow O_2 + O$). For example, it oxidises lead sulfide to lead sulfate: $PbS(s) + 4O_3(g) \rightarrow PbSO_4(s) + 4O_2(g)$.

List two human-made threats to the ozone layer.

1. Nitrogen oxides (like NO) from supersonic jet exhausts. 2. Chlorofluorocarbons (CFCs or Freons) used as refrigerants and in aerosol sprays.

Describe the two most important allotropic forms of sulphur and their relationship.

The two main allotropes are rhombic (α-sulphur) and monoclinic (β-sulphur). Rhombic is stable below 369 K, while monoclinic is stable above it. 369 K is the transition temperature where both forms, which consist of $S_8$ puckered rings, can coexist.

How is sulphur dioxide ($SO_2$) formed industrially, and what is its chemical nature when moist?

Industrially, $SO_2$ is produced by roasting sulphide ores (e.g., $4FeS_2(s) + 11O_2(g) \rightarrow 2Fe_2O_3(s) + 8SO_2(g)$). When moist, it acts as a reducing agent, for example, decolourising acidified potassium permanganate solution.

Briefly outline the three key steps in the Contact Process for the manufacture of sulphuric acid ($H_2SO_4$).

1. Burning sulphur or sulphide ores to generate $SO_2$. 2. Catalytic oxidation of $SO_2$ to $SO_3$ using a $V_2O_5$ catalyst. 3. Absorption of $SO_3$ in concentrated $H_2SO_4$ to form oleum ($H_2S_2O_7$), which is then diluted with water.

State four important chemical characteristics of sulphuric acid.

1. Low volatility. 2. Strong acidic character. 3. Strong dehydrating agent. 4. Moderately strong oxidising agent.

List the elements of Group 17 and explain why they are called "halogens."

Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). They are called "halogens," which means "salt-producers," because they react with metals to form a wide range of salts.

Compare the electron gain enthalpy of fluorine with that of chlorine. Explain the observed trend.

Fluorine has a less negative electron gain enthalpy than chlorine. This is because the small size of the fluorine atom causes strong interelectronic repulsions in the 2p orbitals, making it harder to add an electron compared to the larger chlorine atom.

Why is the bond dissociation enthalpy of $F_2$ smaller than that of $Cl_2$?

The F-F bond is weaker than the Cl-Cl bond due to the large electron-electron repulsion among the lone pairs on the small fluorine atoms, which are very close to each other.

Despite having a less negative electron gain enthalpy, why is fluorine a stronger oxidising agent than chlorine?

Fluorine is a stronger oxidising agent due to its low F-F bond dissociation enthalpy and the very high hydration enthalpy of the small fluoride ion ($F^-$).

Why does fluorine exhibit only a -1 oxidation state, while other halogens can show positive oxidation states like +1, +3, +5, and +7?

As the most electronegative element, fluorine only exhibits a -1 oxidation state. Other halogens can show positive oxidation states because they have d-orbitals and can expand their octet when bonded to a more electronegative element (like F or O).

Describe the reactivity of fluorine and chlorine with water.

Fluorine is highly reactive and oxidises water to oxygen: $2F_2(g) + 2H_2O(l) \rightarrow 4HF(aq) + O_2(g)$. Chlorine reacts with water to form a mixture of hydrochloric acid and hypochlorous acid: $Cl_2(g) + H_2O(l) \rightleftharpoons HCl(aq) + HOCl(aq)$.

What are interhalogen compounds, and why are they generally more reactive than individual halogens (except $F_2$)?

Interhalogen compounds are formed when two different halogens react (e.g., $ClF_3$). They are generally more reactive because the bond between two different halogen atoms (X-X') is weaker than the bond between two identical atoms (X-X), except for the F-F bond.

How is chlorine gas manufactured industrially by Deacon's process?

In Deacon's process, hydrogen chloride gas is oxidised by atmospheric oxygen using a $CuCl_2$ catalyst at 723 K: $4HCl(g) + O_2(g) \rightleftharpoons 2Cl_2(g) + 2H_2O(g)$.

Describe the reactions of chlorine gas with (a) cold and dilute NaOH and (b) hot and concentrated NaOH.

(a) With cold, dilute NaOH, it forms sodium chloride and sodium hypochlorite: $2NaOH + Cl_2 \rightarrow NaCl + NaOCl + H_2O$. (b) With hot, concentrated NaOH, it forms sodium chloride and sodium chlorate: $6NaOH + 3Cl_2 \rightarrow 5NaCl + NaClO_3 + 3H_2O$.

Explain the bleaching action of chlorine.

Chlorine's bleaching action is due to oxidation. In the presence of moisture, chlorine forms hypochlorous acid ($HOCl$), which releases nascent oxygen. This nascent oxygen oxidises coloured substances, rendering them colourless.

What is aqua regia, and what is its main use?

Aqua regia is a 3:1 mixture of concentrated HCl and concentrated $HNO_3$. It is highly corrosive and is used to dissolve noble metals like gold and platinum.

List the elements of Group 18 and state why they are called "noble gases."

Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rn). They are called noble gases because their valence shells are completely filled, making them chemically inert and resistant to forming compounds.

Why do noble gases have very low melting and boiling points?

Noble gases are monoatomic and are held together only by weak van der Waals dispersion forces. Very little energy is needed to overcome these forces, resulting in very low melting and boiling points.

What inspired Neil Bartlett to synthesise the first noble gas compound?

Neil Bartlett observed that molecular oxygen ($O_2$) reacted with $PtF_6$ to form a compound, $O_2^+[PtF_6]^-$. He realised that the first ionisation enthalpy of Xenon was very similar to that of $O_2$, which inspired him to try reacting Xe with $PtF_6$, successfully creating the first noble gas compound.

Describe how xenon difluoride ($XeF_2$) is prepared and its reaction with water.

$XeF_2$ is prepared by the direct reaction of xenon (in excess) with fluorine gas at 673 K and 1 bar. It is readily hydrolysed by water: $2XeF_2(s) + 2H_2O(l) \rightarrow 2Xe(g) + 4HF(aq) + O_2(g)$.

What are the molecular structures of $XeF_2$, $XeF_4$, and $XeF_6$?

$XeF_2$ is linear. $XeF_4$ is square planar. $XeF_6$ is a distorted octahedron.

List two important uses for (a) Helium and (b) Argon.

(a) *Helium: Used to fill weather balloons (as it is light and non-flammable) and in diving gas mixtures to prevent "the bends." (b) Argon:* Used to provide an inert atmosphere for welding and in incandescent light bulbs to prevent the filament from oxidizing.