p-Block Elements-II: One Mark Questions, Book Back Answers and Notes

About p-Block Elements-II

Hard ~150 min study

This unit completes the p-block with groups 15, 16, 17 and 18. Most of the elements here are non-metals, and several of them, such as nitrogen, oxygen, sulphur and chlorine, are used in industry on a very large scale. The unit is therefore a mix of group trends and specific compounds you need to know well.

Structure explains a lot of what you will read. Nitrogen is an inert N₂ molecule with a triple bond, while white phosphorus is a reactive P₄ tetrahedron. Sulphur and phosphorus form many oxoacids, and their structures tell you their basicity and strength. Halogens combine with each other to form interhalogen compounds, and xenon, once thought to be completely inert, forms fluorides and oxides.

In the exam, be ready to write the conditions for the Haber process and the contact process, draw structures of oxoacids and xenon fluorides, and balance the redox reactions of nitric acid and sulphuric acid. Questions on these topics come as both MCQs and longer answers.

What you'll learn

• Describe the physical and chemical properties of group 15 elements and their compounds.
• Write the conditions and equations for making ammonia and nitric acid.
• Compare the allotropes of phosphorus and sulphur from their molecular structures.
• Explain the contact process and the properties of sulphuric acid and other oxoacids of sulphur.
• Predict the shapes and reactivity of interhalogen compounds.
• Describe how xenon fluorides and oxides are formed and state their structures.

Before you start

• Periodic trends from groups 13 and 14.
• Oxidation numbers and electronic configurations.
• Chemical bonding and VSEPR theory for predicting molecular shapes.

p-Block Elements-II explained

Groups 15 to 18

Nitrogen, Phosphorus and Their Compounds

Group 15 begins with nitrogen, obtained industrially by fractional distillation of liquid air. Ammonia is made by the Haber process: nitrogen and hydrogen are passed over an iron catalyst at about 750 K and 200 atm, with a promoter (molybdenum, or small amounts of K₂O and Al₂O₃) to make the catalyst work better. Nitric acid is made from ammonia by the Ostwald process, in which ammonia is oxidised over a platinum gauze. Nitrogen forms oxides in oxidation states from +1 to +5. Phosphorus shows allotropy: white phosphorus is made of discrete \(P_{4}\) tetrahedra and is very reactive, while red phosphorus is polymeric and much more stable.

Oxygen and Sulphur

Group 16 covers oxygen and its allotrope ozone, which absorbs ultraviolet light in the upper atmosphere and is a strong oxidising agent. Sulphuric acid is made by the contact process. Sulphur dioxide is oxidised to sulphur trioxide over vanadium pentoxide at about 720 K, SO₃ is absorbed in concentrated sulphuric acid to give oleum, and oleum is diluted to the required strength. Because the oxidation of SO₂ is exothermic, temperature has to be controlled. Sulphur forms many oxoacids, such as sulphurous acid, sulphuric acid and peroxodisulphuric acid, which differ in the number of oxygen atoms and S–S or O–O links.

Halogens and Interhalogen Compounds

Group 17 focuses on chlorine: its preparation, its bleaching action and the preparation of hydrochloric acid. Halogens are very reactive because they have high electronegativity and a strong tendency to gain one electron to complete their octet. Interhalogen compounds contain two different halogens, such as ClF, BrF₃ and IF₇. The A–X bond is polar and weaker than the X–X bond in the parent halogen, so interhalogens are generally more reactive than the halogens they come from, fluorine excepted.

Noble Gases and Xenon Compounds

Group 18 was long treated as completely inert. In 1962 Neil Bartlett made the first noble gas compound, a xenon hexafluoroplatinate, using the powerful oxidant PtF₆. Xenon reacts directly with fluorine when heated in a nickel vessel, and the ratio of the gases decides whether XeF₂, XeF₄ or XeF₆ forms. Hydrolysis of the fluorides gives xenon oxides such as XeO₃. The shapes of these compounds are predicted using VSEPR theory.

Oxoacids and Acid Strength

Oxoacids appear in every group of this unit. For oxoacids of the same element, the one with the central atom in the higher oxidation state, and so with more oxygen atoms not bonded to hydrogen, is usually the stronger acid. Those oxygen atoms pull electron density away from the O–H bond and stabilise the anion formed when H⁺ leaves. With this idea you can place HClO, HClO₂, HClO₃ and HClO₄ in order of increasing acid strength, and explain why H₃PO₃ is dibasic.

Common mistakes to avoid

• Getting the oxidation state of nitrogen wrong in its oxides. Use the rule that oxygen is -2 and the molecule is neutral: N2O is +1, NO is +2, N2O3 is +3, NO2 is +4 and N2O5 is +5.
• Writing that noble gases form no compounds at all. Xenon forms stable fluorides and oxides with highly electronegative elements such as fluorine and oxygen.
• Drawing red phosphorus as P4 tetrahedra only. White phosphorus has discrete P4 units, while red phosphorus is a polymeric chain of linked P4 units.
• Judging acid strength by the total number of oxygen atoms. Count the oxygen atoms not bonded to hydrogen, and in H3PO3 remember that one hydrogen is bonded directly to phosphorus and is not acidic.
• Mixing up the Haber and contact processes. The Haber process uses an iron catalyst to make ammonia; the contact process uses V2O5 to make sulphuric acid.

Test yourself on these with the practice test, then check the worked reasoning in the solved MCQs.

Frequently asked questions

Why is nitrogen less reactive than phosphorus?

Nitrogen exists as N2 with a triple bond whose bond enthalpy is very high, about 941 kJ mol-1. Breaking it needs a lot of energy, so nitrogen is inert at ordinary temperature. Phosphorus does not form stable multiple bonds and exists as P4 with weaker, strained P–P single bonds, so it reacts much more readily.

What conditions are used in the Haber process?

Nitrogen and hydrogen are passed over an iron catalyst at about 750 K and 200 atm, with a promoter that improves the catalyst. Four moles of gas give two moles of ammonia, so by Le Chatelier's principle high pressure shifts the equilibrium towards ammonia. The reaction is exothermic, so a moderate temperature is chosen as a balance between yield and rate.

Why is sulphuric acid called the king of chemicals?

It is used in so many industries: fertilisers, petroleum refining, detergents, dyes, storage batteries and the manufacture of other chemicals. It is a strong acid, a dehydrating agent and, when hot and concentrated, an oxidising agent. The amount a country consumes is often taken as a sign of its industrial activity.

How are interhalogen compounds formed and why are they reactive?

Two different halogens react directly to form them, for example chlorine and fluorine giving ClF or ClF3. The larger, less electronegative halogen is the central atom. The bond between two different halogens is polar and generally weaker than the X–X bond in the parent halogen, so interhalogens are more reactive than those halogens. Fluorine is the exception: the F–F bond is itself weak, and fluorine is more reactive than its interhalogens.

Can noble gases really form compounds?

Yes. The heavier noble gases, especially xenon, have lower ionisation enthalpies. Xenon reacts with fluorine when the mixture is heated in a sealed nickel vessel, giving XeF2, XeF4 or XeF6 depending on the ratio of gases. Hydrolysis of XeF6 gives XeO3.

How does acid strength change in the oxoacids of chlorine?

It increases in the order HClO < HClO2 < HClO3 < HClO4. Each extra oxygen atom draws electron density from the O–H bond and spreads the negative charge of the anion, so the proton is released more easily. HClO4 is the strongest of the four.

How are the allotropes of sulphur different?

The two main crystalline forms are rhombic (α) sulphur and monoclinic (β) sulphur. Rhombic sulphur is stable below 369 K and monoclinic sulphur above it, so 369 K is the transition temperature. Both are made of crown-shaped S8 molecules, but the molecules pack differently in the crystal, which gives them different shapes and densities.