Chapter 7: Elements of Groups 16, 17 and 18
For most of the nineteenth century, one element defeated every chemist who went after it. Fluorine could be located easily enough in its compounds — in the mineral fluorspar, in the enamel of teeth — but no one could pull the free element out of them. The attempts were dangerous. Several chemists were badly poisoned by the hydrogen fluoride they were handling, and some died. Later writers called them the martyrs of fluorine.
The difficulty was chemical, not technical. The moment fluorine was set free, it attacked whatever was holding it — glass, metal, water, the apparatus itself.
In 1886, Henri Moissan, working in Paris, found a way through. Instead of relying on a chemical reagent he used electrolysis: potassium hydrogen fluoride dissolved in anhydrous hydrogen fluoride, chilled to about −50 °C, in a platinum vessel closed with stoppers cut from fluorspar. A pale yellow gas collected at one electrode, and it set alight almost everything brought near it. Two decades later, Moissan received the Nobel Prize for this work.
Fluorine belongs to group 17. One column further along, in group 18, sit elements that for more than sixty years afterwards refused to form even a single compound. Same block of the periodic table, opposite extremes of behaviour — and the whole difference lies in how many electrons short of a filled shell each atom is. Fig. 7.0 places the three families of this chapter side by side, from the most electron-hungry to the completely satisfied.
Diagram
Fig. 7.0 – Three vertical strips of the p-block placed side by side, labelled Group 16 (ns²np⁴, “two short”), Group 17 (ns²np⁵, “one short”) and Group 18 (ns²np⁶, “complete”), with a left-to-right arrow beneath reading “reactivity falls as the octet fills”; show O, S, Se, Te, Po / F, Cl, Br, I, At / He, Ne, Ar, Kr, Xe, Rn in the strips.
Keep these questions in mind as you read; the chapter will give you everything you need to reason out each answer.
7.1 Introduction
In Std. XI you learnt that the elements of the periodic table are sorted into blocks according to which orbital receives the last electron added to the atom. This last-filling electron is called the differentiating electron, because it is what sets an element apart from the element immediately before it.
In the p-block elements, the differentiating electron enters a p-orbital of the outermost shell. Every shell from the second one onwards contains three p orbitals, and each orbital can hold at most two electrons. A p-subshell can therefore accommodate a maximum of six electrons.
That number six is exactly why the p-block is six columns wide. As the p-subshell is filled one electron at a time, it gives rise to six groups, numbered 13 to 18 in the modern periodic table.
In this chapter we shall study the properties of the elements of groups 16, 17 and 18 — the last three of these six columns. Their outermost p-subshells are, in order, two electrons short of full, one electron short of full, and exactly full. Keep that simple difference in mind, because it will explain almost everything that follows.
7.2 Occurrence
Group 16 — The Oxygen Family
The elements oxygen (\(_{8}\mathrm{O}\)), sulfur (\(_{16}\mathrm{S}\)), selenium (\(_{34}\mathrm{Se}\)), tellurium (\(_{52}\mathrm{Te}\)) and polonium (\(_{84}\mathrm{Po}\)) together constitute Group 16. The group is called the oxygen family, after its first member.
A large number of metal ores are oxides or sulfides, that is, compounds in which a metal is joined to oxygen or to sulfur. Because these elements are the ones that build ores, the group 16 elements are also called chalcogens, or ore-forming elements.
Oxygen is the most abundant of all the elements on earth. It forms 20.95 % by volume of air and about 46.6 % by mass of the earth’s crust, so it is plentiful both in the free state and in combination.
Sulfur, in contrast, forms only 0.034 % by mass of the earth’s crust. It occurs mainly in combined forms — as sulfates such as gypsum (\(\mathrm{CaSO_4{\cdot}2H_2O}\)), epsom salt (\(\mathrm{MgSO_4{\cdot}7H_2O}\)) and baryte (\(\mathrm{BaSO_4}\)).
Sulfur also occurs widely as sulfides, among them galena (\(\mathrm{PbS}\)), zinc blende (\(\mathrm{ZnS}\)) and copper pyrites (\(\mathrm{CuFeS_2}\)). These three are the principal ores from which lead, zinc and copper are obtained, which is a good illustration of why the family earned the name chalcogen.
Selenium and tellurium are found in the same company. They occur as metal selenides and tellurides present within sulfide ores, and are recovered as by-products when those ores are processed.
Polonium is radioactive and has no ore of its own. It is a decay product of thorium and uranium, and is found only in traces wherever those two elements occur.
Group 17 — The Halogens
Fluorine (\(_{9}\mathrm{F}\)), chlorine (\(_{17}\mathrm{Cl}\)), bromine (\(_{35}\mathrm{Br}\)), iodine (\(_{53}\mathrm{I}\)) and astatine (\(_{85}\mathrm{At}\)) constitute Group 17. They are known collectively as the halogens, that is, salt-producing elements.
Halogens have very high electronegativities, so each atom pulls electrons strongly towards itself and reacts readily with almost anything that will part with an electron. As a direct result, the halogens are not found in the free state in nature; they occur only in the form of compounds.
Fluorine occurs mainly as insoluble fluorides: fluorspar (\(\mathrm{CaF_2}\)), cryolite (\(\mathrm{Na_3AlF_6}\)) and fluorapatite (\(\mathrm{3Ca_3(PO_4)_2{\cdot}CaF_2}\)). Cryolite is industrially valuable because molten cryolite is used as the medium in which alumina is dissolved during the electrolytic extraction of aluminium.
Small quantities of fluorine are also present in soil, in fresh water plants, and in the bones and teeth of animals — the same fluoride that hardens tooth enamel.
Sea water contains chlorides, bromides and iodides of Na, K, Mg and Ca. However, it mainly contains sodium chloride, \(\mathrm{NaCl}\), at about 2.5 % by mass.
Where old sea beds have dried up, the deposits left behind contain sodium chloride together with carnallite, \(\mathrm{KCl{\cdot}MgCl_2{\cdot}6H_2O}\).
Marine life also carries iodine within its systems. Sea weed, for example, contains up to 0.5 % iodine, and chile saltpetre contains up to 0.2 % of sodium iodate.
Astatine, the last member of the halogen family, is radioactive and has a half life of 8.1 hours. A half life that short means half of any sample decays away in a little over eight hours, so astatine can never be collected and stored in weighable amounts.
Group 18 — The Noble Gases
The elements helium (\(_{2}\mathrm{He}\)), neon (\(_{10}\mathrm{Ne}\)), argon (\(_{18}\mathrm{Ar}\)), krypton (\(_{36}\mathrm{Kr}\)), xenon (\(_{54}\mathrm{Xe}\)) and radon (\(_{86}\mathrm{Rn}\)) constitute Group 18.
All the noble gases except radon occur in the atmosphere. Their abundance in dry air is about 1 % by volume, with argon as the major constituent of that one per cent.
The main commercial source of helium is not air but natural gas, from which it is separated. Helium and neon are also found in minerals of radioactive origin, for example pitchblende, monazite and cleveite, where they have accumulated as the products of radioactive decay.
Xenon and radon are the rarest elements of the group. Radon is a decay product of \(^{226}\mathrm{Ra}\).
7.3 Electronic Configuration of Elements of Group 16, 17 and 18
Within each of these three groups every element carries the same kind of outermost arrangement, and only the shell number changes as we move down. The general electronic configuration of the group 16 elements is \(\mathrm{ns^2np^4}\), that of the group 17 elements is \(\mathrm{ns^2np^5}\), and the group 18 elements are shown by the \(\mathrm{ns^2np^6}\) configuration, where \(n\) is the number of the outermost shell.
Read the three together and the pattern is plain: moving from group 16 to group 17 to group 18 adds one p-electron at each step, until the p-subshell is completely filled.
Compare these with the nearest noble gas and a second pattern appears. The elements of groups 16 and 17 respectively have two and one electrons less than the stable electronic configuration of the nearest noble gas, while the group 18 elements already possess it.
That shortfall is the most useful single fact about these families. It decides how many electrons an atom must gain or share before it is satisfied, and therefore what charge its ion carries and how many covalent bonds it forms.
Reading a Condensed Configuration
Writing out a full configuration for a heavy element such as polonium would take a long line of symbols. In the condensed form, the symbol of the preceding noble gas is written in square brackets to stand for all the inner electrons, and only the electrons added after it are written out.
Sulfur, for instance, has 16 electrons; the first ten of them match neon exactly, so sulfur is written as \(\mathrm{[Ne]\,3s^2 3p^4}\) instead of \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^4}\). The part inside the brackets is the inert core; the part outside it is the valence shell that does the chemistry. Table 7.1 shows the condensed electronic configuration of the elements of group 16, 17 and 18.
Table 7.1 : Condensed electronic configuration of elements of group 16, 17 and 18
| Group 16 (Oxygen family) | Condensed Electronic Configuration | Group 17 (Halogen family) | Condensed Electronic Configuration | Group 18 (Noble gases) | Condensed Electronic Configuration |
|---|---|---|---|---|---|
| — | — | — | — | \(_{2}\mathrm{He}\) | \(\mathrm{1s^2}\) |
| \(_{8}\mathrm{O}\) | \(\mathrm{[He]\,2s^2 2p^4}\) | \(_{9}\mathrm{F}\) | \(\mathrm{[He]\,2s^2 2p^5}\) | \(_{10}\mathrm{Ne}\) | \(\mathrm{[He]\,2s^2 2p^6}\) |
| \(_{16}\mathrm{S}\) | \(\mathrm{[Ne]\,3s^2 3p^4}\) | \(_{17}\mathrm{Cl}\) | \(\mathrm{[Ne]\,3s^2 3p^5}\) | \(_{18}\mathrm{Ar}\) | \(\mathrm{[Ne]\,3s^2 3p^6}\) |
| \(_{34}\mathrm{Se}\) | \(\mathrm{[Ar]\,3d^{10}4s^2 4p^4}\) | \(_{35}\mathrm{Br}\) | \(\mathrm{[Ar]\,3d^{10}4s^2 4p^5}\) | \(_{36}\mathrm{Kr}\) | \(\mathrm{[Ar]\,3d^{10}4s^2 4p^6}\) |
| \(_{52}\mathrm{Te}\) | \(\mathrm{[Kr]\,4d^{10}5s^2 5p^4}\) | \(_{53}\mathrm{I}\) | \(\mathrm{[Kr]\,4d^{10}5s^2 5p^5}\) | \(_{54}\mathrm{Xe}\) | \(\mathrm{[Kr]\,4d^{10}5s^2 5p^6}\) |
| \(_{84}\mathrm{Po}\) | \(\mathrm{[Xe]\,4f^{14}5d^{10}6s^2 6p^4}\) | \(_{85}\mathrm{At}\) | \(\mathrm{[Xe]\,4f^{14}5d^{10}6s^2 6p^5}\) | \(_{86}\mathrm{Rn}\) | \(\mathrm{[Xe]\,4f^{14}5d^{10}6s^2 6p^6}\) |
Notice in the table that from selenium onwards a completely filled \(d\) subshell — and, in the last row, a filled \(f\) subshell as well — sits inside the valence shell. These inner subshells are full and take no part in ordinary bonding, but they shield the outer electrons from the nucleus rather poorly, and that will matter when we come to atomic size and ionisation enthalpy in the next section.
7.4 Atomic and Physical Properties of Elements of Group 16, 17 and 18
7.4.1 Atomic Properties of Group 16, 17 and 18 Elements
These properties are given in Tables 7.2, 7.3 and 7.4. Before reading the trends, look at the numbers in the tables themselves — every statement that follows is a summary of what those columns actually show.
i. Atomic and Ionic Radii
In groups 16, 17 and 18, atomic and ionic radii increase down the group. The reason is that each successive element has one more quantum shell than the element above it, so the outermost electrons sit further from the nucleus.
The nuclear charge also rises as we go down a group, and that pulls electrons inwards. But the addition of a whole new shell, together with the screening provided by all the inner electrons, outweighs it. The net result is a steady increase in size.
Across a period, atomic or ionic radii decrease with increasing atomic number. Here no new shell is added; electrons keep entering the same shell while protons keep being added to the nucleus. The effective nuclear charge (\(Z_{\text{eff}}\)) — the net pull actually felt by an outer electron after screening by the inner electrons — therefore increases, and the electron cloud is drawn in tighter.
Group 17 elements (halogens) have the smallest atomic radii in their respective periods, since they experience the highest effective nuclear charge among the elements that still belong to the same shell.
ii. Ionisation Enthalpy
The group 16, 17 and 18 elements have high ionisation enthalpy. Ionisation enthalpy is the energy needed to remove the most loosely held electron from an isolated gaseous atom, and these elements hold their electrons tightly because their nuclear pull on a small valence shell is strong.
The ionisation enthalpy decreases down the group due to increase in the atomic size. As the valence electron moves further from the nucleus, less energy is needed to pull it away.
Across a period ionisation enthalpy increases with increase of atomic number. This is due to addition of electrons in the same shell, which leaves the electron more and more firmly held as the nuclear charge builds up.
However, the elements of group 16 have lower ionisation enthalpy values compared to those of group 15 in the corresponding periods, owing to the extra stable half filled electronic configuration of p-orbitals in elements of group 15.
Diagram
Fig. 7.2 – Line graph of first ionisation enthalpy (kJ/mol, y-axis 900–1800) against atomic number for F(9,1680), Cl(17,1256), Br(35,1142), I(53,1008); label each point with the element symbol and mark the y-axis “ΔiH₁ / kJ mol⁻¹”.
iii. Electronegativity
In a group (16, 17 and 18) the electronegativity decreases down the group. Electronegativity measures how strongly an atom attracts a shared pair of electrons towards itself, and a larger atom, with its bonding pair further from the nucleus, attracts it less.
Oxygen has the highest electronegativity next to fluorine amongst all the elements.
Halogens have very high electronegativity. Fluorine is the most electronegative element in the periodic table.
iv. Electron Gain Enthalpy
In groups 16 and 17, electron gain enthalpy becomes less negative down the group. Electron gain enthalpy is the enthalpy change when a gaseous atom accepts an electron; the more negative it is, the more readily the atom takes that electron.
However, in group 16, oxygen has less negative electron gain enthalpy than sulfur due to its small atomic size. The incoming electron is forced into a compact 2p subshell that is already crowded, and the repulsion it meets there cancels part of the energy that would otherwise be released.
In group 17, fluorine has less negative electron gain enthalpy than that of chlorine. This is due to the small size of the fluorine atom, for exactly the same reason.
Group 18 elements (noble gases) have no tendency to accept electrons because of their stable electronic configuration \(ns^2np^6\), and thus have large positive electron gain enthalpy. An extra electron would have to start a new shell altogether, which costs energy rather than releasing it.
7.4.2 Physical Properties of Group 16, 17 and 18 Elements
a. Group 16 Elements (Oxygen Family or Chalcogens)
Oxygen is a gas while the other elements are solids at room temperature.
Oxygen and sulfur are nonmetals, selenium and tellurium are metalloids, while polonium is a metal. Metallic character therefore increases down the group, which is the usual pattern in the p-block: the larger, less electronegative atoms lower down hold their electrons loosely enough to behave like metals. Polonium is radioactive, with a half life of 13.8 days.
Melting and boiling points increase with increasing atomic number.
All the elements of group 16 exhibit allotropy, a property taken up in detail in section 7.7.
b. Group 17 Elements (Halogen Family)
Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid at room temperature. \(\mathrm{F_2}\) is yellow, \(\mathrm{Cl_2}\) greenish yellow, \(\mathrm{Br_2}\) red and \(\mathrm{I_2}\) is violet, in colour.
Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water and are soluble in various organic solvents such as chloroform, carbon disulfide, carbon tetrachloride and hydrocarbons, which give coloured solutions.
Bond dissociation enthalpies of halogen molecules follow the order:
\[\mathrm{Cl-Cl} > \mathrm{Br-Br} > \mathrm{F-F} > \mathrm{I-I}\]
Read that order carefully. It is not a smooth trend down the group — fluorine, the smallest and most electronegative halogen, sits out of place near the bottom. The reason is worked out in Solved Example 7.8.
Diagram
Fig. 7.1 – Two fluorine atoms bonded as F–F, each carrying three lone pairs drawn as shaded lobes; curved double-headed arrows between the facing lone pairs of the two atoms, labelled “lone pair – lone pair repulsion”.
Thus it requires a smaller amount of energy to break the F–F bond.
In the \(\mathrm{Cl_2}\) molecule the lone pairs on each Cl atom are at a larger distance and the repulsion is negligible. Thus the Cl–Cl bond is comparatively stronger. Therefore the bond dissociation enthalpy of \(\mathrm{F_2}\) is lower than that of \(\mathrm{Cl_2}\).
c. Group 18 Elements (Noble Gases)
Noble gases are monoatomic. They exist as single atoms, not as \(\mathrm{X_2}\) molecules, because each atom already has a complete valence shell and gains nothing by pairing up.
They are sparingly soluble in water.
Noble gases have very low melting and boiling points. Helium has the lowest boiling point (4.2 K) of any known substance.
7.5 Anomalous Behaviour
In every group of the p-block, the first member behaves differently from the rest of the family. It is much smaller, it holds electrons far more tightly, and — most importantly — its valence shell has no d-orbitals available. Oxygen in group 16 and fluorine in group 17 are the two clearest examples, and both are worth studying closely because examiners return to them again and again.
7.5.1 Anomalous Behaviour of Oxygen
Oxygen shows the following anomalous properties compared to the other members of group 16.
Atomicity
Oxygen is a diatomic molecule (\(\mathrm{O_2}\)) while the others are polyatomic molecules — for example \(\mathrm{P_4}\), \(\mathrm{S_8}\). Being small, two oxygen atoms can form a strong double bond between themselves; the larger atoms of sulfur and selenium overlap poorly in that way and prefer to link into rings and chains of single bonds instead.
Magnetic Property
Oxygen is paramagnetic while the others are diamagnetic. A paramagnetic substance contains unpaired electrons and is drawn into a magnetic field; a diamagnetic one has all its electrons paired and is weakly pushed out. The reason behind the paramagnetism of \(\mathrm{O_2}\) is taken up in section 7.9.
Oxidation State
Oxygen shows −2, −1 and +2 oxidation states, while the other elements of the group show −2, +2, +4 and +6 oxidation states.
Oxygen cannot exhibit a higher oxidation state due to the absence of vacant d orbitals. To reach +4 or +6, an atom has to unpair its valence electrons and place them in empty orbitals of the same shell; sulfur has empty 3d orbitals for this, but the second shell of oxygen has only 2s and 2p orbitals, and no 2d orbitals exist at all.
Diagram
Fig. 7.4 – Two orbital box diagrams side by side: oxygen 2s²2p⁴ showing only two unpaired electrons and no 2d boxes available, and sulfur 3s²3p⁴ with an empty 3d set, arrows showing promotion of paired 3s and 3p electrons into 3d to give six unpaired electrons; label them “O: covalency 2” and “S: covalency up to 6”.
Nature of Hydrides
The hydride of oxygen (\(\mathrm{H_2O}\)) is a liquid at room temperature while the hydrides of the other members of the group are gases. Oxygen is small and highly electronegative, so the O–H bond is strongly polar and each water molecule is held to its neighbours by hydrogen bonds. Breaking that network needs far more energy than separating \(\mathrm{H_2S}\) molecules, which are held only by weak van der Waals forces.
Diagram
Fig. 7.3 – Three water molecules drawn with the bent H–O–H shape, partial charges δ⁻ on O and δ⁺ on H marked, and dotted lines from the H of one molecule to the lone pair of O on the next labelled “hydrogen bond”.
Covalency
The common covalency of oxygen is 2, and in rare cases it is four. But for the other members of group 16 the covalency can exceed four.
The anomalous behaviour of oxygen is due to the following reasons: small atomic size, high electronegativity, and absence of inner d-orbitals.
7.5.2 Anomalous Behaviour of Fluorine
Fluorine, the first member of group 17, differs in properties from the other members of the group. The anomalous behaviour of fluorine is due to the following reasons: its small atomic size, its high electronegativity, the absence of d-orbitals in its valence shell, and its low F–F bond dissociation enthalpy.
Some anomalous properties of fluorine are listed below.
Ionisation enthalpy, electronegativity and electrode potential are all higher for fluorine than the expected trends shown by the other halogens.
Ionic and covalent radii, melting point, boiling point and electron gain enthalpy are quite lower than expected — the values in Table 7.3 bear this out.
Most of the reactions of fluorine are exothermic. This is due to the short and strong bond formed by it with other elements: little energy is spent breaking the weak F–F bond, while a great deal is released when the strong bonds to the new partner are formed.
It forms only one oxoacid (HOF), while the other halogens form a number of oxoacids. Without d orbitals fluorine cannot expand its octet to hold several oxygen atoms, and being the most electronegative element it cannot take a positive oxidation state in the way chlorine and bromine do.
Hydrogen fluoride is a liquid (b.p. 293 K) due to strong hydrogen bonding, while the other hydrogen halides are gases — the same hydrogen bonding argument that explained liquid water above.
7.6 Chemical Properties of Elements of Groups 16, 17 and 18
Having seen how these atoms are built and how big they are, we can now ask what they actually do. The chemistry of all three families follows from one question: how far is the valence shell from being complete, and can that shell be expanded?
7.6.1 Oxidation State
i. The Group 16 Elements
The group 16 elements have the valence shell electronic configuration \(\mathrm{ns^2np^4}\). They attain a noble gas configuration either by gaining two electrons, forming \(\mathrm{E^{2-}}\) ions, or by sharing two electrons, forming two covalent bonds. These elements, thus, show −2 and +2 oxidation states in their compounds.
Oxygen, being highly electronegative, shows a common oxidation state of −2 except in two cases. In the case of \(\mathrm{OF_2}\) its oxidation state is +2, and in peroxides it shows the oxidation state −1 (\(\mathrm{H_2O_2}\), \(\mathrm{Na_2O_2}\)).
The other elements of the group exhibit +2, +4 and +6 oxidation states, with +4 and +6 being more common. The stability of the higher (+6) oxidation state decreases down the group while the stability of the lower (+4) oxidation state increases down the group, due to the inert pair effect.
Bonding in the +4 and +6 oxidation states is primarily covalent. An atom cannot realistically lose four or six electrons, so these states are reached by sharing rather than by transfer.
Flowchart
FC 7.1 – Deciding the oxidation states available to an element of groups 16–18: start → count valence electrons from the group → is the element the first member of its group (no d orbitals in the valence shell)? → if yes: only the negative state and, for oxygen, the special cases OF₂ (+2) and peroxides (−1); if no: d orbitals available → octet can expand → positive states up to the group maximum (+6 for group 16, +7 for group 17, +8 for xenon) → check inert pair effect for the heaviest members → lower state (+4) becomes the more stable one.
| Element | O | O | S | F |
|---|---|---|---|---|
| Compound | \(\mathrm{H_2O}\) | \(\mathrm{OF_2}\) | \(\mathrm{H_2S}\) | \(\mathrm{HF}\) |
| Oxidation state | −2 | …… | …… | …… |
| Element | Se | Se | Te | Cl |
|---|---|---|---|---|
| Compound | \(\mathrm{SeO_2}\) | \(\mathrm{SeO_3}\) | \(\mathrm{TeF_6}\) | \(\mathrm{HOCl}\) |
| Oxidation state | …… | +6 | …… | …… |
ii. The Group 17 Elements
The group 17 elements are represented by their valence shell electronic configuration \(\mathrm{ns^2np^5}\). They attain a noble gas configuration either by gaining one electron, forming \(\mathrm{E^{-}}\) ions, or by sharing one electron, forming one covalent bond.
All halogens exhibit the −1 oxidation state. However Cl, Br and I exhibit +1, +3, +5 and +7 oxidation states as well. This is because they are less electronegative than F and possess empty d-orbitals in the valence shell, and therefore can expand the octet. The oxidation states +4 and +6 occur in the oxides and oxoacids of Cl and Br.
The fluorine atom has no d-orbitals in its valence shell and therefore cannot expand its octet. Thus fluorine, being the most electronegative element, exhibits mostly the −1 oxidation state.
iii. The Group 18 Elements
Group 18 elements (noble gases) have a stable valence shell electronic configuration \(\mathrm{ns^2np^6}\) with completely filled orbitals. Thus they have no tendency to gain or lose electrons, that is, they are zero valent and mostly exist as monoatomic gases.
However, xenon has a large atomic size and lower ionisation enthalpy compared to He, Ne, Ar and Kr. Hence xenon exhibits higher oxidation states.
Its outermost shell has d-orbitals. The paired electrons of the valence shell can be unpaired and promoted to the empty d-orbitals. The unpaired electrons are then shared with fluorine or oxygen atoms, and covalent compounds showing higher oxidation states — such as \(\mathrm{XeF_2}\) (+2), \(\mathrm{XeF_4}\) (+4), \(\mathrm{XeF_6}\) (+6), \(\mathrm{XeO_3}\) (+6) and \(\mathrm{XeOF_4}\) (+6) — are formed.
Notice that the partners are always fluorine or oxygen. Only the two most electronegative elements pull hard enough to draw electrons out of a noble gas atom at all.
7.6.2 Chemical Reactivity Towards Hydrogen
i. Group 16 Elements
The elements of group 16 react with hydrogen to form hydrides of the type \(\mathrm{H_2E}\), where E = O, S, Se, Te, Po — for example \(\mathrm{H_2O}\), \(\mathrm{H_2S}\), \(\mathrm{H_2Se}\), \(\mathrm{H_2Te}\) and \(\mathrm{H_2Po}\). Some properties of the hydrides of group 16 are given in Table 7.5.
\(\mathrm{H_2O}\) is a colourless, odourless liquid, while \(\mathrm{H_2S}\), \(\mathrm{H_2Se}\), \(\mathrm{H_2Te}\) and \(\mathrm{H_2Po}\) are colourless, bad smelling, poisonous gases at ambient conditions.
All the hydrides have angular structures, which involve \(\mathrm{sp^3}\) hybridisation of the central atom (E). Four electron pairs surround the central atom — two bond pairs and two lone pairs — and the two lone pairs push the bonds closer together, which is why the molecule is bent rather than straight, as drawn in Fig. 7.5.
Diagram
Fig. 7.5 – Bent H–E–H molecule with the central atom E carrying two lone pairs drawn above it; mark the bond angle and show the tetrahedral arrangement of the four electron pairs in faint outline.
Table 7.5 : Properties of hydrides of group 16 elements
| Property | \(\mathrm{H_2O}\) | \(\mathrm{H_2S}\) | \(\mathrm{H_2Se}\) | \(\mathrm{H_2Te}\) |
|---|---|---|---|---|
| m.p (K) | 273 | 188 | 208 | 222 |
| b.p (K) | 373 | 213 | 232 | 269 |
| H–E bond length (pm) | 96 | 134 | 146 | 169 |
| \(\Delta_{\text{diss}}H\) (H–E) kJ/mol | 463 | 347 | 276 | 238 |
| \(\Delta_f H\) kJ/mol | −286 | −20 | 73 | 100 |
| HEH angle (°) | 104 | 92 | 91 | 90 |
| \(\mathrm{p}K_a\) | 14.0 | 7.0 | 3.8 | 2.6 |
The hydrides of group 16 elements are weakly acidic. The acidic character of the hydrides increases, while the thermal stability decreases, from \(\mathrm{H_2O}\) to \(\mathrm{H_2Te}\). This is due to the decrease in the bond dissociation enthalpy of the H–E bond down the group (Table 7.5).
Both changes come from the same cause. A weaker H–E bond releases its hydrogen more readily as \(\mathrm{H^+}\), which means greater acidity; and a weaker bond also comes apart more easily on heating, which means lower thermal stability. Follow the \(\mathrm{p}K_a\) row of Table 7.5 and you can watch acidity rise as the number falls from 14.0 to 2.6.
All the hydrides except \(\mathrm{H_2O}\) possess reducing property, which increases in the order \(\mathrm{H_2S} < \mathrm{H_2Se} < \mathrm{H_2Te}\).
ii. Group 17 Elements
The elements of group 17 react with hydrogen to give hydrogen halides.
\[\mathrm{H_2} + \mathrm{X_2} \rightarrow 2\,\mathrm{HX} \qquad (\text{where } \mathrm{X} = \mathrm{F},\ \mathrm{Cl},\ \mathrm{Br},\ \mathrm{I})\]
Some of the properties of the hydrogen halides are given in Table 7.6.
Table 7.6 : Properties of hydrides of group 17 elements
| Property | HF | HCl | HBr | HI |
|---|---|---|---|---|
| m.p (K) | 190 | 159 | 185 | 222 |
| b.p (K) | 293 | 189 | 206 | 238 |
| Bond length (H–X) pm | 91.7 | 127.4 | 141.4 | 160.9 |
| \(\Delta_{\text{diss}}H^{\circ}\) kJ/mol | 574 | 432 | 363 | 295 |
| \(\mathrm{p}K_a\) | 3.2 | −7.0 | −9.5 | −10.0 |
The acidic strength of the halogen acids increases in the order
\[\mathrm{HF} < \mathrm{HCl} < \mathrm{HBr} < \mathrm{HI}\]
It is due to the decreasing bond dissociation enthalpy of the H–X bond in the order \(\mathrm{HF} > \mathrm{HCl} > \mathrm{HBr} > \mathrm{HI}\).
The thermal stability of the hydrogen halides decreases in the order \(\mathrm{HF} > \mathrm{HCl} > \mathrm{HBr} > \mathrm{HI}\). It is due to the decrease in bond dissociation enthalpy of the H–X bond down the group.
iii. Group 18 Elements (Noble Gases)
Noble gases are chemically inert towards hydrogen due to their stable electronic configuration.
7.6.3 Reactivity Towards Oxygen
i. Group 16 Elements
All the elements of group 16 form oxides of the type \(\mathrm{EO_2}\) and \(\mathrm{EO_3}\), where E = S, Se, Te, Po.
\(\mathrm{EO_2}\) type oxides: ozone (\(\mathrm{O_3}\)) and sulfur dioxide (\(\mathrm{SO_2}\)) are gases, while selenium dioxide (\(\mathrm{SeO_2}\)) is a solid. They are acidic in nature and react with water to form acids.
\[\mathrm{SO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_3}\ (\text{sulfurous acid})\]
\[\mathrm{SeO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2SeO_3}\ (\text{selenious acid})\]
The reducing property of the dioxides decreases from \(\mathrm{SO_2}\) to \(\mathrm{TeO_2}\). \(\mathrm{SO_2}\) is reducing while \(\mathrm{TeO_2}\) serves as an oxidising agent. This follows the stability pattern noted earlier: for sulfur the +6 state is easily reached, so \(\mathrm{SO_2}\) readily gives up electrons to move up to it, whereas for tellurium the +4 state is the comfortable one and \(\mathrm{TeO_2}\) would rather take electrons in.
\(\mathrm{EO_3}\) type oxides, \(\mathrm{SO_3}\), \(\mathrm{SeO_3}\) and \(\mathrm{TeO_3}\), are also acidic in nature. They dissolve in water to form acids.
\[\mathrm{SeO_3} + \mathrm{H_2O} \rightarrow \mathrm{H_2SeO_4}\ (\text{selenic acid})\]
\[\mathrm{TeO_3} + 3\,\mathrm{H_2O} \rightarrow \mathrm{H_6TeO_6}\ (\text{telluric acid})\]
ii. Group 17 Elements
Elements of group 17 (halogens) form many oxides with oxygen, but most of them are unstable.
Fluorine forms two oxides, \(\mathrm{OF_2}\) and \(\mathrm{O_2F_2}\). However, only \(\mathrm{OF_2}\) is thermally stable at 298 K. Both are strong fluorinating agents. \(\mathrm{O_2F_2}\) oxidises plutonium to \(\mathrm{PuF_6}\), and the reaction is used in removing plutonium as \(\mathrm{PuF_6}\) from spent nuclear fuel.
Chlorine oxides, \(\mathrm{Cl_2O}\), \(\mathrm{ClO_2}\), \(\mathrm{Cl_2O_6}\) and \(\mathrm{Cl_2O_7}\), are highly reactive oxidising agents and tend to explode.
\(\mathrm{ClO_2}\) is used as a bleaching agent for paper pulp and textiles, and in water treatment.
Bromine oxides, \(\mathrm{Br_2O}\), \(\mathrm{BrO_2}\) and \(\mathrm{BrO_3}\), are the least stable halogen oxides (middle row anomaly). They are very powerful oxidising agents.
Iodine oxides, \(\mathrm{I_2O_4}\), \(\mathrm{I_2O_5}\) and \(\mathrm{I_2O_7}\), are insoluble solids and decompose on heating. \(\mathrm{I_2O_5}\) is a very good oxidising agent and is used for the estimation of carbon monoxide.
The higher oxides of halogens are more stable than the lower ones.
iii. Group 18 Elements
Noble gas elements are chemically inert and do not directly react with oxygen.
7.6.4 Reactivity Towards Halogens
i. Group 16 Elements
Elements of group 16 react with halogens to give a large number of halides of the types \(\mathrm{EX_6}\), \(\mathrm{EX_4}\) and \(\mathrm{EX_2}\), where E = S, Se, Te.
Hexahalides: \(\mathrm{SF_6}\), \(\mathrm{SeF_6}\) and \(\mathrm{TeF_6}\) are formed by direct combination. They are colourless gases. They have \(\mathrm{sp^3d^2}\) hybridisation and possess octahedral structure. \(\mathrm{SF_6}\) is an exceptionally stable halide for steric reasons — the six fluorine atoms wrap the sulfur atom so completely that no reagent can reach it.
The stability of the halides decreases in the order
\[\text{fluorides} > \text{chlorides} > \text{bromides} > \text{iodides}\]
Tetrahalides: \(\mathrm{SF_4}\), \(\mathrm{SeF_4}\), \(\mathrm{TeF_4}\) and \(\mathrm{TeCl_4}\) have \(\mathrm{sp^3d}\) hybridisation and thus a trigonal bipyramidal geometry, with one equatorial position occupied by a lone pair.
Dihalides: \(\mathrm{SCl_2}\), \(\mathrm{SeCl_2}\) and \(\mathrm{TeCl_2}\) have \(\mathrm{sp^3}\) hybridisation, so their four electron pairs take up a tetrahedral arrangement in which two positions are occupied by lone pairs.
Diagram
Fig. 7.6 – Three structures in a row: SF₆ octahedral with six F atoms; SF₄ trigonal bipyramidal with a lone pair in an equatorial position, giving the see-saw shape; SCl₂ with a tetrahedral arrangement of four electron pairs, two of them lone pairs, giving a bent molecule. Label hybridisation under each: sp³d², sp³d, sp³.
Monohalides are dimeric in nature. For example, \(\mathrm{S_2F_2}\), \(\mathrm{S_2Cl_2}\), \(\mathrm{Se_2Cl_2}\) and \(\mathrm{SeBr_2}\). These dimeric halides undergo disproportionation.
\[2\,\mathrm{Se_2Cl_2} \rightarrow \mathrm{SeCl_4} + 3\,\mathrm{Se}\]
Disproportionation means that one and the same element is oxidised and reduced in the same reaction: selenium goes from +1 in \(\mathrm{Se_2Cl_2}\) up to +4 in \(\mathrm{SeCl_4}\) and down to 0 in free selenium.
ii. Group 17 Elements
Halogens (group 17 elements) combine amongst themselves to form a number of compounds known as interhalogen compounds.
These are of the following types: \(\mathrm{XX'}\), \(\mathrm{XX'_3}\), \(\mathrm{XX'_5}\) and \(\mathrm{XX'_7}\), where X is the halogen atom with the larger size and X’ is the halogen atom with the smaller size. More details of interhalogen compounds are included in section 7.12.
iii. Group 18 Elements
Group 18 elements (noble gases) are chemically inert. Krypton and xenon, however, react directly with fluorine to give their fluorides. For example,
\[\mathrm{Xe}\text{(g)} + \mathrm{F_2}\text{(g)} \xrightarrow{\ 673\ \text{K},\ 1\ \text{atm}\ } \mathrm{XeF_2}\text{(s)}\]
Xenon fluorides \(\mathrm{XeF_2}\), \(\mathrm{XeF_4}\) and \(\mathrm{XeF_6}\) are crystalline and colourless, and they sublime readily at 298 K. They are powerful fluorinating agents.
7.6.5 Reactivity Towards Metals
i. Group 16 Elements
Elements of group 16 react with metals to form the corresponding compounds — oxides, sulfides and selenides in which the group 16 element takes its −2 state.
\[4\,\mathrm{Al} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{Al_2O_3}\]
\[\mathrm{Cu} + \mathrm{S} \rightarrow \mathrm{CuS}\]
\[\mathrm{Mg} + \mathrm{Se} \rightarrow \mathrm{MgSe}\ (\text{magnesium selenide})\]
\[2\,\mathrm{Au} + 3\,\mathrm{Te} \rightarrow \mathrm{Au_2Te_3}\ (\text{gold telluride})\]
Gold telluride minerals are worth knowing about because prospectors have been fooled by them: the ore carries no golden glint at all, and gold has to be recovered from it by roasting off the tellurium.
ii. Group 17 Elements
Elements of group 17 (halogens) react with metals instantly to give metal halides.
\[2\,\mathrm{Na}\text{(s)} + \mathrm{Cl_2}\text{(g)} \rightarrow 2\,\mathrm{NaCl}\text{(s)}\]
\[\mathrm{Mg}\text{(s)} + \mathrm{Br_2}\text{(l)} \rightarrow \mathrm{MgBr_2}\text{(s)}\ (\text{magnesium bromide})\]
The ionic character of the halides decreases in the order \(\mathrm{MF} > \mathrm{MCl} > \mathrm{MBr} > \mathrm{MI}\), where M is a monovalent metal. Fluorine is the smallest and most electronegative halogen, so it takes the electron most completely and gives the most ionic bond; iodide is large and easily distorted, so its bond to the metal has more covalent character.
The metal halides having metals in their higher oxidation states are more covalent than the ones having metals in lower oxidation states. For example, \(\mathrm{SnCl_4}\), \(\mathrm{PbCl_4}\), \(\mathrm{SbCl_5}\) and \(\mathrm{UF_6}\) are more covalent than \(\mathrm{SnCl_2}\), \(\mathrm{PbCl_2}\), \(\mathrm{SbCl_3}\) and \(\mathrm{UF_4}\) respectively. A metal ion in a high oxidation state is small and highly charged, so it pulls the electron cloud of the halide ion towards itself and the bond becomes more covalent.
iii. Group 18 Elements
Noble gases do not directly react with metals.
7.7 Allotropy
Elements of the group 16 exhibit allotropy. The same element can be packaged in more than one way — as molecules of different size, or as rings stacked differently in a crystal — and each package has its own colour, melting point and density even though the atoms in it are identical.
It is worth separating this from a term that sounds similar. Polymorphism is the existence of a compound in more than one crystalline form, whereas allotropy refers to different forms of a single element. Calcium carbonate occurring as calcite and as aragonite is polymorphism; oxygen occurring as \(\mathrm{O_2}\) and as \(\mathrm{O_3}\) is allotropy.
Oxygen has two allotropes, \(\mathrm{O_2}\) and \(\mathrm{O_3}\) (ozone). Here the two forms differ in the number of atoms in the molecule itself, and their properties differ sharply, as section 7.9 will show.
Sulfur exists in a number of allotropic forms. Rhombic and monoclinic sulfur are the important allotropes of sulfur. Both are non metallic.
Selenium exists in two allotropic forms, red (non metallic) and grey (metallic).
Tellurium exists in two allotropic forms, (i) crystalline and (ii) amorphous.
Polonium reveals two allotropic forms, \(\alpha\) and \(\beta\), both metallic.
Read down that list and a pattern appears. In oxygen the allotropes are simply different molecules; by the time we reach tellurium and polonium, the difference is one of metallic character and crystal packing. This mirrors the shift from nonmetal to metal down the group.
Flowchart
FC 7.2 – Photocopying with a selenium drum: selenium-coated drum given a uniform positive charge (step 1) → drum exposed to the image; light discharges the lit areas, charge remains on the dark areas (step 2) → negatively charged toner particles attracted to the charged areas (step 3) → image transferred from drum to sheet of paper (step 4) → heating fixes the image; drum flooded with light and cleaned (step 5) → returns to step 1 for the next cycle.
7.7.1 Allotropes of Sulfur
Sulfur exhibits numerous allotropic forms. However, rhombic sulfur (\(\alpha\)-sulfur) and monoclinic sulfur (\(\beta\)-sulfur) are the most important allotropes of sulfur, and they are compared in Table 7.7.
Table 7.7 : Allotropes of sulfur
| Rhombic Sulfur (\(\alpha\)-Sulfur) | Monoclinic Sulfur (\(\beta\)-Sulfur) | |
|---|---|---|
| Color | Pale Yellow | Bright yellow solid |
| Shape | Orthorhombic crystals | Needle shaped monoclinic crystals |
| M. P. | 385.8 K | 393 K |
| Density | 2.069 g/cm³ | 1.989 g/cm³ |
| Solubility | Insoluble in water and soluble in \(\mathrm{CS_2}\) | Soluble in \(\mathrm{CS_2}\) |
| Stability | Stable below 369 K and transforms to \(\beta\)-sulphur above this temperature | Stable above 369 K and transforms into \(\alpha\)-sulphur below this temperature |
| Structure | \(\mathrm{S_8}\) molecules having puckered ring structure | \(\mathrm{S_8}\) molecules with puckered ring structure |
| Method of preparation | It is prepared by evaporation of roll sulphur in \(\mathrm{CS_2}\) | Rhombic sulphur melted in a dish and cooled till crust is formed. Two holes are made in the crust and the remaining liquid is poured out to give needle shaped crystals of \(\beta\)-sulphur |
The temperature 369 K is worth remembering, because it is the point at which the two forms change places: below it the rhombic form is the stable one, above it the monoclinic form takes over.
The Sulfur Rings
Several modifications of sulfur containing 6 to 20 sulfur atoms per ring have been synthesised. In the \(\mathrm{S_8}\) molecule the ring is puckered and has a crown shape, with an S–S bond length of 204 pm and a bond angle of about 107°, as shown in Fig. 7.7.
Diagram
Fig. 7.7 – Puckered crown-shaped S₈ ring drawn in perspective, with eight S atoms alternately above and below the mean plane; mark one S–S bond length as 204 pm and one S–S–S bond angle as 107°.
In cyclo-\(\mathrm{S_6}\) the ring adopts the chair form instead, with a bond length of 205.7 pm and a bond angle of 102.2°, drawn in Fig. 7.8.
Diagram
Fig. 7.8 – Chair form of the cyclo-S₆ ring shown in perspective like a cyclohexane chair; mark the S–S bond length 205.7 pm and the bond angle 102.2°.
At elevated temperature (about 1000 K), \(\mathrm{S_2}\) is the dominant species, which like \(\mathrm{O_2}\) is paramagnetic. The heat has broken the rings apart into small diatomic molecules, and those molecules turn out to have the same electronic peculiarity as dioxygen.
7.8 Oxoacids
7.8.1 Oxoacids of Sulfur
Sulfur forms a number of oxoacids. Some of them are unstable and cannot be isolated. They are known to exist in aqueous solutions or in the form of their salts.
Some important oxoacids of sulfur, together with their structures, are the following.
Sulfurous acid, \(\mathrm{H_2SO_3}\)
Sulfuric acid, \(\mathrm{H_2SO_4}\)
Di or pyrosulfuric acid, \(\mathrm{H_2S_2O_7}\)
Peroxy monosulfuric acid, \(\mathrm{H_2SO_5}\)
Peroxy disulfuric acid, \(\mathrm{H_2S_2O_8}\)
Thiosulfuric acid, \(\mathrm{H_2S_2O_3}\)
The structures of all six are collected in Fig. 7.9. Notice while reading them that sulfur sits at the centre of every one, holding some combination of –OH groups, doubly bonded oxygen atoms, and — in three cases — an extra link that gives the acid its name.
Diagram
Fig. 7.9 – Six structural formulae of the oxoacids of sulfur drawn side by side and labelled: (i) H₂SO₃, S with two OH, one S=O and a lone pair; (ii) H₂SO₄, S with two OH and two S=O; (iii) H₂S₂O₇, two SO₃H units joined through a bridging O; (iv) H₂SO₅, S with one OH, two S=O and one O–OH peroxo group; (v) H₂S₂O₈, two SO₃H units joined through an O–O peroxo link; (vi) H₂S₂O₃, like H₂SO₄ but with one doubly bonded O replaced by a terminal S.
7.8.2 Oxoacids of Halogens
Halogens form several oxoacids, which are set out in Table 7.8. Only four oxoacids have been isolated in pure form: hypofluorous acid (HOF), perchloric acid (\(\mathrm{HClO_4}\)), iodic acid (\(\mathrm{HIO_3}\)) and metaperiodic acid (\(\mathrm{H_2IO_6}\)). The others are stable only in aqueous solutions or in the form of their salts.
Read Table 7.8 across the rows and the pattern is clear at once: fluorine appears in one row only, while chlorine, bromine and iodine appear in several. This is the consequence of the point made in section 7.6 — fluorine has no d orbitals, cannot expand its octet, and cannot take a positive oxidation state beyond +1.
Table 7.8 : Oxoacids of halogens
| Oxidation state of X | Generic name | Oxoacids of fluorine | Oxoacids of chlorine | Oxoacids of bromine | Oxoacids of iodine |
|---|---|---|---|---|---|
| +1 | Hypohalous acid (HXO) | HOF | HOCl | HOBr | HOI |
| +3 | Halous acid (\(\mathrm{HXO_2}\)) | − | HOClO | − | − |
| +5 | Halic acid (\(\mathrm{HXO_3}\)) | − | \(\mathrm{HOClO_2}\) | \(\mathrm{HOBrO_2}\) | \(\mathrm{HOIO_2}\) |
| +7 | Perhalic acid (\(\mathrm{HXO_4}\)) | − | \(\mathrm{HOClO_3}\) | \(\mathrm{HOBrO_3}\) | \(\mathrm{HOIO_3}\) |
The acid strength of the halogen oxoacids increases with the increasing oxidation state of the halogen. For example, acid strength increases from \(\mathrm{HClO}\), a weak acid (\(K_a = 3.5 \times 10^{-8}\)), to \(\mathrm{HClO_4}\), a very strong acid (\(K_a \gg 1\)).
Structures of the Oxoacids of Chlorine
The four oxoacids of chlorine are built on the same H–O–Cl skeleton, with oxygen atoms added one at a time to the chlorine. Their structures are drawn in Fig. 7.10.
Hypochlorous acid, HOCl
Chlorous acid, HOClO or \(\mathrm{HClO_2}\)
Chloric acid, \(\mathrm{HClO_3}\)
Perchloric acid, \(\mathrm{HClO_4}\)
Diagram
Fig. 7.10 – Four structural formulae in a row: HOCl drawn as H–O–Cl bent at oxygen; HClO₂ as H–O–Cl with one Cl=O; HClO₃ as H–O–Cl with two Cl=O; HClO₄ as H–O–Cl with three Cl=O. Label the oxidation state of chlorine beneath each as +1, +3, +5 and +7.
7.9 Oxygen and Compounds of Oxygen
7.9.1 Dioxygen
a. Preparation
i. Laboratory methods
Oxygen is prepared in the laboratory by taking a compound that already contains oxygen and persuading it to let go.
By heating oxygen containing salts such as chlorates, nitrates and permanganates:
\[2\,\mathrm{KClO_3}\text{(s)} \xrightarrow{\ \text{Heat},\ \mathrm{MnO_2}\ } 2\,\mathrm{KCl}\text{(s)} + 3\,\mathrm{O_2}\text{(g)}\]
By thermal decomposition of oxides of metals:
\[2\,\mathrm{Ag_2O}\text{(s)} \xrightarrow{\ \Delta\ } 4\,\mathrm{Ag}\text{(s)} + \mathrm{O_2}\text{(g)}\]
\[2\,\mathrm{HgO}\text{(s)} \xrightarrow{\ \Delta\ } 2\,\mathrm{Hg}\text{(l)} + \mathrm{O_2}\text{(g)}\]
\[2\,\mathrm{PbO_2}\text{(s)} \xrightarrow{\ \Delta\ } 2\,\mathrm{PbO}\text{(s)} + \mathrm{O_2}\text{(g)}\]
By decomposition of hydrogen peroxide in presence of a catalyst such as finely divided metals and manganese dioxide:
\[2\,\mathrm{H_2O_2}\text{(aq)} \xrightarrow{\ \Delta,\ \mathrm{MnO_2}\ } 2\,\mathrm{H_2O}\text{(l)} + \mathrm{O_2}\text{(g)}\]
ii. Electrolysis
Dioxygen can be prepared on a large scale by the electrolysis of water, when hydrogen is liberated at the cathode and oxygen at the anode.
\[2\,\mathrm{H_2O} \xrightarrow{\ \text{Electrolysis}\ } 2\,\mathrm{H_2} + \mathrm{O_2}\]
iii. Industrial method
Dioxygen is obtained from air, by first removing carbon dioxide and water vapour. The remaining gases are liquified subsequently. This is followed by fractional distillation, which gives dinitrogen and dioxygen. The separation works because liquid nitrogen boils at a lower temperature than liquid oxygen, so nitrogen distils off first. The steps are summarised in FC 7.3.
Flowchart
FC 7.3 – Industrial manufacture of dioxygen: air → removal of CO₂ and water vapour → compression and cooling → liquefaction of the remaining gases → fractional distillation → dinitrogen distils off first (lower boiling point) → dioxygen left behind and collected.
b. Physical properties
Dioxygen is a colourless and odourless gas.
Dioxygen is sparingly soluble in water: 30.8 cm³ of \(\mathrm{O_2}\) dissolves in 1000 cm³ of water at 293 K. A small amount of dissolved dioxygen is sufficient to sustain marine and aquatic life.
It liquifies at 90 K and freezes at 55 K.
Oxygen has three stable isotopes, \(^{16}\mathrm{O}\), \(^{17}\mathrm{O}\) and \(^{18}\mathrm{O}\). All three have eight protons and differ only in the number of neutrons, so they are chemically alike.
Molecular oxygen, \(\mathrm{O_2}\), exhibits paramagnetism.
c. Chemical Properties
i. Reaction with metals
Dioxygen directly reacts with almost all metals except Au and Pt to form their oxides.
\[2\,\mathrm{Ca} + \mathrm{O_2} \rightarrow 2\,\mathrm{CaO}\]
\[4\,\mathrm{Al} + 3\,\mathrm{O_2} \rightarrow 2\,\mathrm{Al_2O_3}\]
Gold and platinum resist this attack, which is one reason they survive in nature as free metals and are used for jewellery and laboratory ware.
ii. Reaction with nonmetals
Dioxygen reacts with nonmetals (except noble gases) to form their oxides.
\[\mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2}\]
\[\mathrm{P_4} + 5\,\mathrm{O_2} \rightarrow \mathrm{P_4O_{10}}\]
iii. Reaction with some compounds
\[2\,\mathrm{ZnS} + 3\,\mathrm{O_2} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO} + 2\,\mathrm{SO_2}\]
\[\mathrm{CH_4} + 2\,\mathrm{O_2} \rightarrow \mathrm{CO_2} + 2\,\mathrm{H_2O}\]
\[2\,\mathrm{SO_2} + \mathrm{O_2} \xrightarrow{\ \mathrm{V_2O_5}\ } 2\,\mathrm{SO_3}\]
\[4\,\mathrm{HCl} + \mathrm{O_2} \xrightarrow{\ \mathrm{CuCl_2},\ \Delta\ } 2\,\mathrm{Cl_2} + 2\,\mathrm{H_2O}\]
The first of these is the roasting of a sulfide ore, and the last two will return later in this chapter as steps in the manufacture of sulfuric acid and of chlorine.
d. Uses
Dioxygen is important for respiration, to sustain animal and aquatic life.
It is used in the manufacture of steel.
It is used in the oxyacetylene flame for welding and cutting of metals.
Oxygen cylinders are widely used in hospitals, in high altitude flying and in mountaineering.
It is used in the combustion of fuels; for example, hydrazine in liquid oxygen provides tremendous thrust (energy) in rockets.
Diagram
Fig. 7.11 – Molecular orbital energy level diagram for O₂ with atomic orbitals of the two O atoms on either side; fill σ(2s)², σ(2s)², σ(2p_z)², π(2p_x)², π(2p_y)², and show single arrows of the same spin in π(2p_x) and π*(2p_y); label these two “unpaired electrons — paramagnetic”.
7.9.2 Simple Oxides
Oxides can be classified into (a) acidic oxides, (b) basic oxides, (c) amphoteric oxides and (d) neutral oxides. The classification is based on one simple test: what does the oxide do when it meets water, an acid or a base?
a. Acidic oxides
An oxide which dissolves in water to give an acid, or reacts with a base to give a salt, is called an acidic oxide. For example, \(\mathrm{SO_2}\), \(\mathrm{SO_3}\), \(\mathrm{CO_2}\), \(\mathrm{N_2O_5}\), \(\mathrm{Cl_2O_7}\) etc.
\[\mathrm{SO_2} + \mathrm{H_2O} \rightarrow \mathrm{H_2SO_3}\]
\[\mathrm{SO_3} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2SO_4} + \mathrm{H_2O}\]
Generally, oxides of nonmetals are acidic oxides.
b. Basic oxides
An oxide which dissolves in water to give a base, or reacts with an acid to give a salt, is called a basic oxide. For example, \(\mathrm{Na_2O}\), \(\mathrm{CaO}\), \(\mathrm{BaO}\) etc.
\[\mathrm{CaO} + \mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2}\]
\[\mathrm{BaO} + 2\,\mathrm{HCl} \rightarrow \mathrm{BaCl_2} + \mathrm{H_2O}\]
These are oxides of metals, in which the metal has given its electrons away fairly completely, so the oxide ion is free to accept a proton from water or acid.
c. Amphoteric oxides
The oxide which reacts with a base as well as with an acid to give a salt is called an amphoteric oxide. For example, \(\mathrm{Al_2O_3}\).
Acting as an acidic oxide, towards a base:
\[\mathrm{Al_2O_3}\text{(s)} + 6\,\mathrm{NaOH}\text{(aq)} + 3\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{Na_3[Al(OH)_6]}\text{(aq)}\]
Acting as a basic oxide, towards an acid:
\[\mathrm{Al_2O_3}\text{(s)} + 6\,\mathrm{HCl}\text{(aq)} + 9\,\mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{[Al(H_2O)_6]^{3+}}\text{(aq)} + 6\,\mathrm{Cl^{-}}\text{(aq)}\]
d. Neutral oxides
The oxides which are neither acidic nor basic are called neutral oxides. For example, \(\mathrm{CO}\), \(\mathrm{NO}\), \(\mathrm{N_2O}\) etc.
7.9.3 Ozone
Ozone (\(\mathrm{O_3}\)) is an allotrope of oxygen. Oxygen in the upper atmosphere absorbs energy in the form of ultra-violet light and changes to atomic oxygen, which combines with molecular oxygen to form \(\mathrm{O_3}\).
\[\mathrm{O_2} \xrightarrow{\ \text{U.V. light}\ } \mathrm{O} + \mathrm{O}\]
\[\mathrm{O_2} + \mathrm{O} \rightarrow \mathrm{O_3}\]
The layer of ozone protects the earth’s surface from harmful ultraviolet (U.V.) radiations. Hence it is called the ‘ozone umbrella’.
a. Preparation of Ozone
Ozone is prepared in the laboratory by passing a silent electric discharge through pure and dry oxygen in an apparatus called an ozoniser, drawn in Fig. 7.12. As the conversion of oxygen to ozone is only 10 %, the product is known as ozonised oxygen. It is an endothermic process.
\[3\,\mathrm{O_2}\text{(g)} \rightleftharpoons 2\,\mathrm{O_3} \qquad \Delta H = +142\ \text{kJ mol}^{-1}\]
The discharge is called silent because it passes as a diffuse glow between the glass walls rather than as a spark; a spark would supply so much heat that the ozone formed would break down again.
Diagram
Fig. 7.12 – Ozoniser: two concentric glass tubes with tin foil coatings connected to a high voltage induction coil, dry oxygen entering at one end and ozonised oxygen leaving at the other; label “silent electric discharge”.
b. Physical properties of ozone
Pure ozone is a pale-blue gas, a dark blue liquid and a violet-black solid.
Ozone has a characteristic smell. When inhaled in concentrations above 100 ppm, it causes nausea and headache.
It is diamagnetic in nature — unlike dioxygen, all the electrons in the ozone molecule are paired.
c. Chemical Properties
i. Oxidising property
Ozone is a powerful oxidising agent, as it easily decomposes to liberate nascent oxygen.
\[\mathrm{O_3} \rightarrow \mathrm{O_2} + \mathrm{O}\]
Nascent oxygen means a single, freshly formed oxygen atom, which is far more reactive than the paired-up \(\mathrm{O_2}\) molecule; this atom is what does the oxidising.
Ozone oxidises lead sulfide to lead sulfate, and iodide ions to iodine.
\[\mathrm{PbS}\text{(s)} + 4\,\mathrm{O_3}\text{(g)} \rightarrow \mathrm{PbSO_4}\text{(s)} + 4\,\mathrm{O_2}\text{(g)}\]
\[2\,\mathrm{KI}\text{(aq)} + \mathrm{H_2O}\text{(l)} + \mathrm{O_3}\text{(g)} \rightarrow 2\,\mathrm{KOH}\text{(aq)} + \mathrm{I_2}\text{(g)} + \mathrm{O_2}\text{(g)}\]
Ozone oxidises nitrogen oxide and gives nitrogen dioxide.
\[\mathrm{NO}\text{(g)} + \mathrm{O_3}\text{(g)} \rightarrow \mathrm{NO_2}\text{(g)} + \mathrm{O_2}\text{(g)}\]
Hence the nitrogen oxide emitted from the exhaust systems of supersonic jet aeroplanes can bring forth depletion of the ozone layer in the upper atmosphere.
ii. Bleaching property
Ozone acts as a good bleaching agent due to its oxidising nature.
\[\mathrm{O_3} \rightarrow \mathrm{O} + \mathrm{O_2}\]
\[\text{Coloured matter} + \mathrm{O} \rightarrow \text{colourless matter}\]
Ozone bleaches in the absence of moisture, so it is also known as dry bleach. This is a useful contrast to keep in mind, because chlorine — as section 7.11 will show — bleaches only when moisture is present.
iii. Reducing property
Ozone reduces peroxides to oxides.
\[\mathrm{H_2O_2} + \mathrm{O_3} \rightarrow \mathrm{H_2O} + 2\,\mathrm{O_2}\]
\[\mathrm{BaO_2} + \mathrm{O_3} \rightarrow \mathrm{BaO} + 2\,\mathrm{O_2}\]
iv. Ozone depletion
The thinning of the ozone layer in the upper atmosphere is called ozone depletion.
The ozone (\(\mathrm{O_3}\)) layer in the upper atmosphere absorbs harmful UV radiations from the sun, thus protecting people on the earth.
Depletion of the ozone layer in the upper atmosphere is caused by nitrogen oxide released from the exhaust systems of cars or supersonic jet aeroplanes.
\[\mathrm{NO}\text{(g)} + \mathrm{O_3}\text{(g)} \rightarrow \mathrm{NO_2}\text{(g)} + \mathrm{O_2}\text{(g)}\]
Depletion (thinning) of the ozone layer can also be caused by chlorofluoro carbons (freons), used in aerosols and refrigerators, and their subsequent escape into the atmosphere.
The depletion of the ozone layer has been most pronounced in the polar regions, especially over Antarctica.
Ozone depletion is a major environmental problem because it increases the amount of ultraviolet (UV) radiation that reaches the earth’s surface, thus causing an increase in the rate of skin cancer, eye cataracts, and genetic as well as immune system damage among people. The cycle of formation and destruction is set out in FC 7.4.
Flowchart
FC 7.4 – Ozone in the stratosphere: O₂ absorbs UV → splits into O + O → O + O₂ → O₃ (ozone umbrella formed); destruction branch: O₃ + NO from jet and car exhausts → NO₂ + O₂, and O₃ attacked by chlorine atoms released from CFCs → ozone layer thins → more UV reaches the earth’s surface → skin cancer, cataracts, genetic and immune damage.
d. Structure of Ozone
Ozone is an angular molecule. The two O—O bond lengths in the ozone molecule are identical, 128 pm, and the O—O—O bond angle is about 117°. It is a resonance hybrid of two canonical forms, drawn in Fig. 7.13.
The word identical is the key to the structure. If ozone had one single bond and one double bond, the two lengths would differ; because they are equal, neither canonical form is the real molecule. The true molecule is the resonance hybrid, in which the extra pair of electrons is spread evenly over both bonds.
Diagram
Fig. 7.13 – Two canonical forms of ozone drawn side by side with a double-headed resonance arrow between them, each showing one O=O and one O–O with lone pairs and formal charges; to the right, the resonance hybrid with both bond lengths marked 128 pm and the bond angle 117°.
e. Uses of ozone
Ozone is used for air purification at crowded places like cinema halls, tunnels and railways.
It is used in sterilizing drinking water by oxidising all germs and bacteria.
It is used for bleaching ivory, oils, starch, wax and delicate fabrics such as silk.
It is used in the manufacture of synthetic camphor, potassium permanganate, etc.
7.10 Compounds of Sulfur
7.10.1 Sulfur Dioxide
a. Preparation
i. From sulfur
Sulfur dioxide gas can be prepared by burning sulfur in air.
\[\mathrm{S}\text{(s)} + \mathrm{O_2}\text{(g)} \rightarrow \mathrm{SO_2}\text{(g)}\]
ii. From sulfite
In the laboratory, sulfur dioxide is prepared by treating an aqueous solution of sodium sulfite with dilute sulfuric acid.
\[\mathrm{Na_2SO_3} + \mathrm{H_2SO_4}\text{(aq)} \rightarrow \mathrm{Na_2SO_4} + \mathrm{H_2O}\text{(l)} + \mathrm{SO_2}\text{(g)}\]
This is the usual bench method because the gas comes off steadily and needs no heating — the stronger, less volatile acid simply displaces the weaker, volatile one from its salt.
iii. From sulfides (industrial method)
Sulfur dioxide can be prepared by roasting zinc sulfide and iron pyrites.
\[2\,\mathrm{ZnS}\text{(s)} + 3\,\mathrm{O_2}\text{(g)} \xrightarrow{\ \Delta\ } 2\,\mathrm{ZnO}\text{(s)} + 2\,\mathrm{SO_2}\text{(g)}\]
\[4\,\mathrm{FeS_2}\text{(s)} + 11\,\mathrm{O_2}\text{(g)} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe_2O_3}\text{(s)} + 8\,\mathrm{SO_2}\text{(g)}\]
Roasting means heating a sulfide ore strongly in a plentiful supply of air. Industry gets its sulfur dioxide almost free this way, as a by-product of extracting zinc and iron.
b. Physical properties of \(\mathrm{SO_2}\)
Sulfur dioxide is a colourless gas with a pungent smell.
It is poisonous in nature.
\(\mathrm{SO_2}\) is highly soluble in water, and its solution in water is called sulfurous acid.
It liquifies at room temperature under a pressure of 2 atm and boils at 263 K.
c. Chemical Properties
i. Reaction with \(\mathrm{Cl_2}\)
Sulfur dioxide reacts with chlorine in the presence of charcoal (catalyst) to form sulfuryl chloride.
\[\mathrm{SO_2}\text{(g)} + \mathrm{Cl_2}\text{(g)} \xrightarrow{\ \text{charcoal}\ } \mathrm{SO_2Cl_2}\text{(l)}\]
ii. Reaction with \(\mathrm{O_2}\)
Sulfur dioxide is oxidised by dioxygen in the presence of vanadium(V) oxide to sulfur trioxide.
\[2\,\mathrm{SO_2}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ \mathrm{V_2O_5}\ } 2\,\mathrm{SO_3}\text{(g)}\]
iii. Reaction with NaOH
Sulfur dioxide readily reacts with sodium hydroxide solution to form sodium sulfite.
\[2\,\mathrm{NaOH} + \mathrm{SO_2} \rightarrow \mathrm{Na_2SO_3} + \mathrm{H_2O}\]
iv. Reaction with \(\mathrm{Na_2SO_3}\)
Sulfur dioxide reacts with sodium sulfite solution to form sodium hydrogen sulfite.
\[\mathrm{Na_2SO_3} + \mathrm{H_2O}\text{(l)} + \mathrm{SO_2} \rightarrow 2\,\mathrm{NaHSO_3}\]
These last two reactions together confirm that \(\mathrm{SO_2}\) is an acidic oxide: it neutralises a base, and then reacts further with its own salt.
v. Reducing property
Sulfur dioxide acts as a reducing agent in the presence of moisture. In each of the reactions below, sulfur moves from the +4 state to the +6 state, giving up two electrons to whatever it reduces.
Moist sulfur dioxide reduces ferric salts into ferrous salts.
\[2\,\mathrm{Fe^{3+}} + \mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{Fe^{2+}} + \mathrm{SO_4^{2-}} + 4\,\mathrm{H^{+}}\]
Moist sulfur dioxide decolourises acidified potassium permanganate(VII) solution.
\[2\,\mathrm{KMnO_4} + 5\,\mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow \mathrm{K_2SO_4} + 2\,\mathrm{MnSO_4} + 2\,\mathrm{H_2SO_4}\]
The loss of the purple colour is a standard laboratory test for a reducing gas, since the deep purple \(\mathrm{MnO_4^{-}}\) ion is reduced to the almost colourless \(\mathrm{Mn^{2+}}\) ion.
Moist sulfur dioxide reduces halogens to halogen acids.
\[\mathrm{I_2} + \mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow \mathrm{H_2SO_4} + 2\,\mathrm{HI}\]
d. Structure of \(\mathrm{SO_2}\)
Sulfur dioxide is angular, with an O—S—O bond angle of 119.5°. The S—O double bond arises from \(\mathrm{d\pi\!-\!p\pi}\) bonding, in which a filled p orbital of oxygen overlaps sideways with an empty d orbital of sulfur — an option open to sulfur precisely because its valence shell has d orbitals, and closed to oxygen. It is a resonance hybrid of two canonical forms, shown in Fig. 7.14.
Diagram
Fig. 7.14 – Two canonical forms of SO₂ side by side with a double-headed resonance arrow between them; each drawn bent at sulfur with one S=O and one S–O, lone pairs marked on S and on both O atoms; mark the O–S–O bond angle 119.5°.
e. Uses
Sulfur dioxide is used in refining of petroleum and sugar, in bleaching wool and silk, as an anti-chlor and disinfectant, and as a preservative.
It is used in the manufacture of \(\mathrm{H_2SO_4}\) and \(\mathrm{NaHSO_3}\).
Liquid \(\mathrm{SO_2}\) is used as a solvent to dissolve a number of organic and inorganic chemicals.
7.10.2 Sulfuric Acid, \(\mathrm{H_2SO_4}\)
a. Preparation
Sulfuric acid is manufactured by the Contact process, which involves the following three steps. The plant is drawn in Fig. 7.15, and the chemical sequence is summarised in FC 7.5.
Diagram
Fig. 7.15 – Flow diagram for the manufacture of sulfuric acid: sulfur burner with air inlet → dust precipitator → washing and cooling tower with water spray → drying tower with conc. H₂SO₄ spray → arsenic purifier containing gelatinous hydrated ferric oxide → preheater → catalytic converter packed with V₂O₅ → absorption tower with conc. H₂SO₄ spray → oleum outlet; mark waste water and waste acid outlets.
Flowchart
FC 7.5 – Contact process: S or FeS₂ roasted in air → SO₂ → purification of the gas (dust, moisture and arsenic removed, since arsenic poisons the catalyst) → catalytic oxidation over V₂O₅ at 720 K and 2 bar → SO₃ → absorbed in concentrated H₂SO₄ → oleum (H₂S₂O₇) → diluted with water → H₂SO₄ of the desired concentration, 96–98 % pure.
i. Formation of sulfur dioxide
Sulfur or sulfide ore (iron pyrites) on burning or roasting in air produces sulfur dioxide.
\[\mathrm{S}\text{(s)} + \mathrm{O_2}\text{(g)} \xrightarrow{\ \Delta\ } \mathrm{SO_2}\text{(g)}\]
\[4\,\mathrm{FeS_2}\text{(s)} + 11\,\mathrm{O_2}\text{(g)} \xrightarrow{\ \Delta\ } 2\,\mathrm{Fe_2O_3}\text{(s)} + 8\,\mathrm{SO_2}\text{(g)}\]
ii. Catalytic oxidation to sulfur trioxide
Sulfur dioxide is oxidised catalytically with oxygen to sulfur trioxide, in the presence of \(\mathrm{V_2O_5}\) catalyst.
\[2\,\mathrm{SO_2}\text{(g)} + \mathrm{O_2} \xrightarrow{\ \mathrm{V_2O_5}\ } 2\,\mathrm{SO_3}\text{(g)}\]
The reaction is exothermic and reversible, and the forward reaction leads to a decrease in volume. Therefore low temperature (720 K) and high pressure (2 bar) are favourable conditions for the maximum yield of \(\mathrm{SO_3}\).
Count the molecules on each side and the reasoning becomes clear: three molecules of gas become two, so squeezing the mixture pushes the equilibrium forward. Since heat is given out, keeping the temperature down also favours the product. In practice the temperature cannot be lowered too far, because the reaction would then become impractically slow — which is why a catalyst is used to obtain a workable rate at a moderate temperature.
iii. Absorption to oleum, and dilution
Sulfur trioxide gas (from the catalytic converter) is absorbed in concentrated \(\mathrm{H_2SO_4}\) to produce oleum. Dilution of oleum with water gives sulfuric acid of the desired concentration.
\[\mathrm{SO_3}\text{(g)} + \mathrm{H_2SO_4} \rightarrow \mathrm{H_2S_2O_7}\ (\text{oleum})\]
\[\mathrm{H_2S_2O_7} + \mathrm{H_2O} \rightarrow 2\,\mathrm{H_2SO_4}\]
The sulfuric acid obtained by the contact process is 96–98 % pure.
b. Physical properties of \(\mathrm{H_2SO_4}\)
Sulfuric acid is a colourless, dense, oily liquid.
It has a density (specific gravity) of 1.84 g/cm³ at 298 K.
It freezes at 283 K and boils at 611 K.
It is highly corrosive and produces severe burns on the skin.
c. Chemical Properties
i. Acidic property
Sulfuric acid ionises in aqueous solution in two steps.
\[\mathrm{H_2SO_4}\text{(aq)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{H_3O^{+}}\text{(aq)} + \mathrm{HSO_4^{-}}\text{(aq)} \qquad K_a > 10\]
\[\mathrm{HSO_4^{-}}\text{(aq)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{H_3O^{+}}\text{(aq)} + \mathrm{SO_4^{2-}}\text{(aq)} \qquad K_a = 1.2 \times 10^{-2}\]
The greater value of \(K_a\) (\(K_a > 10\)) means that \(\mathrm{H_2SO_4}\) is largely dissociated into \(\mathrm{H^{+}}\) and \(\mathrm{HSO_4^{-}}\) ions. Thus \(\mathrm{H_2SO_4}\) is a strong acid.
Compare the two constants and you can see that the second proton is released far less readily than the first. Pulling a positive proton away from an already negative \(\mathrm{HSO_4^{-}}\) ion is much harder than pulling it from the neutral molecule.
ii. Reaction with metals and nonmetals (oxidising property)
Metals and nonmetals are both oxidised by hot, concentrated sulfuric acid, which itself gets reduced to \(\mathrm{SO_2}\).
\[\mathrm{Cu} + 2\,\mathrm{H_2SO_4}\ (\text{Conc.}) \rightarrow \mathrm{CuSO_4} + \mathrm{SO_2} + 2\,\mathrm{H_2O}\]
\[\mathrm{S} + 2\,\mathrm{H_2SO_4}\ (\text{Conc.}) \rightarrow 3\,\mathrm{SO_2} + 2\,\mathrm{H_2O}\]
\[\mathrm{C} + 2\,\mathrm{H_2SO_4}\ (\text{Conc.}) \rightarrow \mathrm{CO_2} + 2\,\mathrm{SO_2} + 2\,\mathrm{H_2O}\]
The oxidising properties of sulfuric acid depend on its concentration and temperature. In dilute solutions at room temperature, \(\mathrm{H_2SO_4}\) behaves like HCl, oxidising only those metals that stand above hydrogen in the e.m.f. series.
\[\mathrm{Fe}\text{(s)} + 2\,\mathrm{H^{+}}\text{(aq)} \rightarrow \mathrm{Fe^{2+}}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
Hot, concentrated \(\mathrm{H_2SO_4}\) is a better oxidising agent than the dilute, cold acid. It oxidises metals like copper, which dilute acid leaves untouched.
iii. Dehydrating property
Concentrated sulfuric acid is a strong dehydrating agent.
Sulfuric acid removes water from sugar and carbohydrates. The carbon left behind is called sugar charcoal, and the process is called charring.
\[\mathrm{C_{12}H_{22}O_{11}} \xrightarrow{\ \text{conc. } \mathrm{H_2SO_4}\ } 12\,\mathrm{C} + 11\,\mathrm{H_2O}\]
The acid does not simply drive off water that was already present — it strips hydrogen and oxygen atoms out of the sugar molecule itself in the ratio in which they occur in water, which is why a white solid turns into a rising black column of carbon.
iv. Reaction with salts
Concentrated sulfuric acid decomposes the salts of more volatile acids to the corresponding acid.
\[\mathrm{NaCl} + \mathrm{H_2SO_4} \rightarrow \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[\mathrm{KNO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{KHSO_4} + \mathrm{HNO_3}\]
\[\mathrm{CaF_2} + \mathrm{H_2SO_4} \rightarrow \mathrm{CaSO_4} + 2\,\mathrm{HF}\]
Sulfuric acid succeeds here not because it is the stronger acid in every sense, but because it has a very high boiling point (611 K). On heating, the volatile acid escapes as a gas and cannot come back, so the reaction is driven to completion.
d. Uses
Sulfuric acid is a very important industrial chemical. It is used in the manufacture of fertilizers — for example, ammonium sulfate and superphosphate.
It is used in the manufacture of pigments, paints and dyestuff intermediates.
It is used in petroleum refining and in the detergent industry.
It is used in metallurgy, for cleaning of metals, electroplating and galvanising.
It is used in storage batteries, as a laboratory reagent, and in the manufacture of nitrocellulose products.
7.11 Chlorine and Compounds of Chlorine
7.11.1 Chlorine
Chlorine was discovered by Scheele, a German-Swedish chemist, in 1774, by the action of HCl on \(\mathrm{MnO_2}\). In 1810, Davy established its elementary nature and suggested the name chlorine on account of its colour, from the Greek chloros, meaning greenish yellow.
a. Preparation
1. Chlorine can be prepared by the oxidation of hydrochloric acid with any of the following oxidising agents.
i. Manganese dioxide
\[\mathrm{MnO_2} + 4\,\mathrm{HCl} \rightarrow \mathrm{MnCl_2} + \mathrm{Cl_2} + 2\,\mathrm{H_2O}\]
ii. Potassium permanganate
\[2\,\mathrm{KMnO_4} + 16\,\mathrm{HCl} \rightarrow 2\,\mathrm{KCl} + 2\,\mathrm{MnCl_2} + 8\,\mathrm{H_2O} + 5\,\mathrm{Cl_2}\]
In both reactions the chloride ion loses an electron and is oxidised to free chlorine, while manganese is reduced — from +4 to +2 in the first case and from +7 to +2 in the second. Potassium permanganate is the more powerful oxidant, which is why it needs no heating.
2. Chlorine can also be prepared by the action of concentrated sulfuric acid on a mixture of sodium chloride (common salt) and manganese dioxide. The reaction takes place in two steps.
\[4\,\mathrm{NaCl} + 4\,\mathrm{H_2SO_4} \rightarrow 4\,\mathrm{NaHSO_4} + 4\,\mathrm{HCl}\]
\[\mathrm{MnO_2} + 4\,\mathrm{HCl} \rightarrow \mathrm{MnCl_2} + 2\,\mathrm{H_2O} + \mathrm{Cl_2}\]
Adding the two steps gives the overall change:
\[4\,\mathrm{NaCl} + \mathrm{MnO_2} + 4\,\mathrm{H_2SO_4} \rightarrow 4\,\mathrm{NaHSO_4} + \mathrm{MnCl_2} + 2\,\mathrm{H_2O} + \mathrm{Cl_2}\]
The sulfuric acid first sets HCl free from common salt, and the manganese dioxide then oxidises that HCl — so the second method is really the first method with the hydrochloric acid generated in the flask.
b. Manufacture of Chlorine
i. Deacon’s process
Chlorine is manufactured by the oxidation of hydrogen chloride gas by atmospheric oxygen in the presence of \(\mathrm{CuCl_2}\) as catalyst at 723 K.
\[4\,\mathrm{HCl} + \mathrm{O_2} \xrightarrow{\ \mathrm{CuCl_2}\ } 2\,\mathrm{Cl_2} + 2\,\mathrm{H_2O}\]
ii. Electrolytic process
By the electrolysis of brine (concentrated NaCl solution), chlorine is liberated at the anode.
\[\mathrm{NaCl} \rightleftharpoons \mathrm{Na^{+}} + \mathrm{Cl^{-}}\]
At the cathode:
\[2\,\mathrm{H_2O} + 2\,\mathrm{e^{-}} \rightarrow \mathrm{H_2} + 2\,\mathrm{OH^{-}}\]
\[\mathrm{Na^{+}} + \mathrm{OH^{-}} \rightarrow \mathrm{NaOH}\]
At the anode:
\[\mathrm{Cl^{-}} \rightarrow \mathrm{Cl} + \mathrm{e^{-}}\]
\[\mathrm{Cl} + \mathrm{Cl} \rightarrow \mathrm{Cl_2}\]
Three saleable products come out of this one cell — chlorine at the anode, hydrogen at the cathode and sodium hydroxide left in the solution — which is why the process is the backbone of the chemical industry. The arrangement is shown in Fig. 7.16 and the products are traced in FC 7.6.
Diagram
Fig. 7.16 – Electrolytic cell for brine: tank of concentrated NaCl solution with an anode and a cathode, chlorine gas bubbling off at the anode, hydrogen gas at the cathode, and NaOH accumulating in the solution; label the ion movements Cl⁻ → anode and Na⁺ → cathode.
Flowchart
FC 7.6 – Electrolysis of brine: concentrated NaCl solution → dissociates into Na⁺ and Cl⁻ → branch 1: Cl⁻ migrates to anode → loses an electron → Cl atoms pair up → Cl₂ gas collected; branch 2: at the cathode water is reduced → H₂ gas + OH⁻ → OH⁻ combines with Na⁺ → NaOH remains in solution.
c. Physical Properties of Chlorine
Chlorine is a greenish-yellow gas having a pungent and suffocating odour.
It is poisonous in nature.
It can be easily liquified into a greenish yellow liquid, which boils at 239 K.
It dissolves in water to give chlorine water.
It is about 2.5 times heavier than air, which is why the gas collects near the floor and is always handled in a fume cupboard.
d. Chemical Properties of Chlorine
i. Reaction with metals
Chlorine reacts with metals to form chlorides.
\[2\,\mathrm{Al} + 3\,\mathrm{Cl_2} \rightarrow 2\,\mathrm{AlCl_3}\]
\[2\,\mathrm{Na} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{NaCl}\]
\[2\,\mathrm{Fe} + 3\,\mathrm{Cl_2} \rightarrow 2\,\mathrm{FeCl_3}\]
Notice that iron is taken to its higher (+3) state. Chlorine is a strong enough oxidant to push a metal to its upper oxidation state, which is a useful point of contrast with dilute acids, where iron stops at +2.
ii. Reaction with nonmetals
Chlorine reacts with nonmetals to form their chlorides.
\[\mathrm{P_4} + 6\,\mathrm{Cl_2} \rightarrow 4\,\mathrm{PCl_3}\]
\[\mathrm{S_8} + 4\,\mathrm{Cl_2} \rightarrow 4\,\mathrm{S_2Cl_2}\]
iii. Affinity for hydrogen
Chlorine has a great affinity for hydrogen. It reacts with hydrogen, and with compounds containing hydrogen, to form HCl.
\[\mathrm{H_2} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{HCl}\]
\[\mathrm{H_2S} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{HCl} + \mathrm{S}\]
In the second reaction chlorine strips the hydrogen out of hydrogen sulfide and leaves sulfur behind — a direct demonstration of that affinity.
iv. Reaction with \(\mathrm{NH_3}\)
Chlorine, when reacted with an excess of ammonia, gives ammonium chloride and nitrogen.
\[8\,\mathrm{NH_3}\ (\text{Excess}) + 3\,\mathrm{Cl_2} \rightarrow 6\,\mathrm{NH_4Cl} + \mathrm{N_2}\]
An excess of chlorine reacts with ammonia to give nitrogen trichloride, which is explosive.
\[\mathrm{NH_3} + 3\,\mathrm{Cl_2}\ (\text{Excess}) \rightarrow \mathrm{NCl_3} + 3\,\mathrm{HCl}\]
v. Reaction with alkali
Chlorine reacts with cold and dilute alkali to produce a mixture of chloride and hypochlorite. When reacted with hot concentrated alkali, chloride and chlorate are produced.
\[2\,\mathrm{NaOH}\ (\text{Cold and dilute}) + \mathrm{Cl_2} \rightarrow \mathrm{NaCl} + \mathrm{NaOCl} + \mathrm{H_2O}\]
\[6\,\mathrm{NaOH}\ (\text{Hot and conc.}) + 3\,\mathrm{Cl_2} \rightarrow 5\,\mathrm{NaCl} + \mathrm{NaClO_3} + 3\,\mathrm{H_2O}\]
Both are disproportionation reactions: the chlorine that entered as \(\mathrm{Cl_2}\) (oxidation state 0) leaves partly as chloride (−1) and partly as hypochlorite (+1) or chlorate (+5).
Chlorine, when reacted with dry slaked lime, gives bleaching powder.
\[2\,\mathrm{Ca(OH)_2} + 2\,\mathrm{Cl_2} \rightarrow \mathrm{Ca(OCl)_2} + \mathrm{CaCl_2} + 2\,\mathrm{H_2O}\]
vi. Reaction with hydrocarbons
Chlorine reacts with saturated hydrocarbons to give substitution products, and with unsaturated hydrocarbons gives addition products.
\[\mathrm{CH_4} + \mathrm{Cl_2} \xrightarrow{\ \text{U.V.}\ } \mathrm{CH_3Cl} + \mathrm{HCl}\]
Methane gives methyl chloride, and a hydrogen atom has been replaced — that is substitution.
\[\mathrm{H_2C}\!=\!\mathrm{CH_2} + \mathrm{Cl_2} \rightarrow \mathrm{CH_2Cl\!-\!CH_2Cl}\]
Ethene gives 1,2-dichloroethane, and nothing has been lost — the double bond has simply opened to accept both chlorine atoms. That is addition.
vii. Oxidising property
Chlorine oxidises ferrous salts to ferric salts and sulfites to sulfates.
\[2\,\mathrm{FeSO_4} + \mathrm{H_2SO_4} + \mathrm{Cl_2} \rightarrow \mathrm{Fe_2(SO_4)_3} + 2\,\mathrm{HCl}\]
\[\mathrm{Na_2SO_3} + \mathrm{Cl_2} + \mathrm{H_2O} \rightarrow \mathrm{Na_2SO_4} + 2\,\mathrm{HCl}\]
It oxidises sulfur dioxide to sulfur trioxide, and iodine to iodate. In the presence of water they form sulfuric acid and iodic acid respectively.
\[\mathrm{SO_2} + 2\,\mathrm{H_2O} + \mathrm{Cl_2} \rightarrow \mathrm{H_2SO_4} + 2\,\mathrm{HCl}\]
\[\mathrm{I_2} + 6\,\mathrm{H_2O} + 5\,\mathrm{Cl_2} \rightarrow 2\,\mathrm{HIO_3} + 10\,\mathrm{HCl}\]
The last reaction is worth pausing over: chlorine, being the more electronegative halogen, forces iodine all the way up to the +5 state. A halogen higher in the group can always oxidise one below it.
viii. Bleaching property
Chlorine requires the presence of moisture (water) for bleaching. It liberates nascent oxygen from water, which is responsible for its oxidising and bleaching property.
\[\mathrm{Cl_2} + \mathrm{H_2O} \rightarrow \mathrm{HCl} + \mathrm{HOCl}\]
\[\mathrm{HOCl} \rightarrow \mathrm{HCl} + [\mathrm{O}]\]
Chlorine bleaches vegetable matter or coloured organic matter in the presence of moisture to colourless matter.
\[\text{Coloured organic matter} + [\mathrm{O}] \rightarrow \text{Colourless organic matter}\]
Compare this with ozone, which bleaches without moisture and is therefore called a dry bleach. Chlorine has to make its oxidising agent out of water first.
e. Uses
Chlorine is used for the purification (sterilizing) of drinking water.
It is used for bleaching wood pulp required for the manufacture of paper and rayon, and for bleaching cotton and textiles.
It is used for the extraction of metals like gold and platinum.
It is used in the manufacture of dyes, drugs and organic compounds such as \(\mathrm{CCl_4}\), \(\mathrm{CHCl_3}\), DDT, refrigerants, etc.
It is used in the preparation of poisonous gases such as phosgene (\(\mathrm{COCl_2}\)), tear gas (\(\mathrm{CCl_3NO_2}\)) and mustard gas (\(\mathrm{ClCH_2CH_2SCH_2CH_2Cl}\)).
7.11.2 Hydrogen Chloride
Hydrogen chloride was prepared by Glauber in 1648 by heating common salt with concentrated sulfuric acid. Davy, in 1810, showed that it is a compound of hydrogen and chlorine.
a. Preparation
In the laboratory, hydrogen chloride is prepared by heating sodium chloride (common salt) with concentrated sulfuric acid.
\[\mathrm{NaCl} + \mathrm{H_2SO_4} \xrightarrow{\ 420\ \text{K}\ } \mathrm{NaHSO_4} + \mathrm{HCl}\]
\[\mathrm{NaHSO_4} + \mathrm{NaCl} \xrightarrow{\ 420\ \text{K}\ } \mathrm{Na_2SO_4} + \mathrm{HCl}\]
HCl gas can be dried by passing it through concentrated sulfuric acid. This works because sulfuric acid absorbs water strongly but does not react with hydrogen chloride — recall from section 7.10 that concentrated \(\mathrm{H_2SO_4}\) oxidises HBr and HI, but not HCl, which is why the same drying trick cannot be used for those two gases.
b. Physical properties of HCl
Hydrogen chloride is a colourless and pungent smelling gas.
It can be easily liquified to a colourless liquid (b.p. 189 K), which freezes to a white crystalline solid (m.p. 159 K).
It is highly soluble in water.
c. Chemical properties
i. Acidic property
Hydrogen chloride is highly soluble in water and ionises as follows:
\[\mathrm{HCl}\text{(g)} + \mathrm{H_2O}\text{(l)} \rightarrow \mathrm{H_3O^{+}}\text{(aq)} + \mathrm{Cl^{-}}\text{(aq)} \qquad K_a = 10^{7}\]
The aqueous solution of HCl gas is called hydrochloric acid. The high value of the dissociation constant (\(K_a\)) indicates that it is a strong acid in water — the equilibrium lies so far to the right that practically every molecule has given up its proton.
ii. Reaction with \(\mathrm{NH_3}\)
Hydrochloric acid reacts with ammonia and gives white fumes of ammonium chloride.
\[\mathrm{NH_3} + \mathrm{HCl} \rightarrow \mathrm{NH_4Cl}\]
The dense white cloud formed when the two gases meet is the standard test for either of them.
iii. Reaction with noble metals
When three parts of concentrated HCl and one part of concentrated \(\mathrm{HNO_3}\) are mixed, aqua regia is formed.
Noble metals like gold and platinum get dissolved in aqua regia.
\[\mathrm{Au} + 4\,\mathrm{H^{+}} + \mathrm{NO_3^{-}} + 4\,\mathrm{Cl^{-}} \rightarrow \mathrm{AuCl_4^{-}} + \mathrm{NO} + 2\,\mathrm{H_2O}\]
\[3\,\mathrm{Pt} + 16\,\mathrm{H^{+}} + 4\,\mathrm{NO_3^{-}} + 18\,\mathrm{Cl^{-}} \rightarrow 3\,\mathrm{PtCl_6^{2-}} + 4\,\mathrm{NO} + 8\,\mathrm{H_2O}\]
d. Uses
Hydrochloric acid is used as an important laboratory reagent, in cleaning and pickling metal surfaces, and in the preparation of chlorides and of aqua regia as described above.
7.12 Interhalogen Compounds
We have seen that all halogen molecules are diatomic. They form binary compounds with hydrogen; with oxygen they form oxyacids; and with halogens they form interhalogen compounds.
Although all halogens belong to the same group, they have different electronegativities. Due to this difference in electronegativity, two or more halogen atoms combine to form species which may be ionic or neutral. The neutral molecules are called interhalogen compounds — for example, ClF and \(\mathrm{BrF_3}\).
The ionic species are worth noting alongside. Ions formed by the combination of different halogens are called polyhalide ions, or interhalogen ions — for example \(\mathrm{K^{+}[Cl_3^{-}]}\) and \(\mathrm{[NH_4]^{+}[I_5]^{-}}\), which contain the \(\mathrm{Cl_3^{-}}\) and \(\mathrm{I_5^{-}}\) ions.
Table 7.9 : Various types of interhalogen compounds
| Element | Fluoride | Chloride | Bromide | Iodide |
|---|---|---|---|---|
| Chlorine | ClF, \(\mathrm{ClF_3}\), \(\mathrm{ClF_5}\) | − | − | − |
| Bromine | BrF, \(\mathrm{BrF_3}\), \(\mathrm{BrF_5}\) | BrCl | − | − |
| Iodine | IF, \(\mathrm{IF_3}\), \(\mathrm{IF_5}\), \(\mathrm{IF_7}\) | ICl, \(\mathrm{ICl_3}\) | IBr | − |
Read Table 7.9 down the left column and across the top and the rule stated in the definition shows itself. Every entry pairs a larger halogen with a smaller one; there is no fluorine row and no iodide column, because fluorine is never the central atom and iodine is never the attached one.
7.12.1 Classification
Depending on their composition, interhalogen compounds are classified into four types.
| Type | Example |
|---|---|
| \(\mathrm{XX'}\) | ClF, BrF, BrCl, ICl, IBr |
| \(\mathrm{XX'_3}\) | \(\mathrm{ClF_3}\), \(\mathrm{BrF_3}\), \(\mathrm{IF_3}\) |
| \(\mathrm{XX'_5}\) | \(\mathrm{ClF_5}\), \(\mathrm{BrF_5}\), \(\mathrm{IF_5}\) |
| \(\mathrm{XX'_7}\) | \(\mathrm{IF_7}\) |
In the general formula \(\mathrm{XX'_n}\), X is the halogen having the larger size and is more electropositive. X’ is the halogen having the smaller size and more electronegativity.
7.12.2 General Characteristics of Interhalogen Compounds
1. The compound is considered as the halide of X. For example, in ClF the halogen having the larger size is chlorine, which is more electropositive than F, and hence the interhalogen compound is named as chlorine monofluoride. Here \(n\) is the number of atoms of X’ attached to X.
As the ratio [radius of X : radius of X’] increases, the value of \(n\) also increases. This is simply a matter of room — a large central atom with small partners can fit more of them around itself, which is why \(\mathrm{IF_7}\) exists while \(\mathrm{ClF_7}\) does not.
2. Interhalogen compounds have an even number of atoms — 2, 4, 6 or 8. For example, \(\mathrm{ClF_3}\) has 4 atoms.
3. The properties of interhalogen compounds are generally intermediate between those of the halogens from which they are made.
4. The central halogen exhibits different oxidation states in different interhalogen compounds.
5. The number of X’ atoms in the compounds is always odd.
6. They are all diamagnetic — every electron in the molecule is paired.
| Formula | Name |
|---|---|
| ClF | Chlorine monofluoride |
| \(\mathrm{ClF_3}\) | …… |
| …… | Chlorine pentafluoride |
| BrF | …… |
| …… | Bromine pentafluoride |
| ICl | …… |
| \(\mathrm{ICl_3}\) | …… |
Table 7.10 : States of interhalogen compounds at 25 °C
| Type | Compound | State and appearance |
|---|---|---|
| \(\mathrm{XX'}\) | ClF | Colorless gas |
| BrF | Pale brown gas | |
| BrCl | Gas | |
| ICl | Ruby red solid (\(\alpha\)-form); brown red solid (\(\beta\)-form) | |
| IBr | Black solid | |
| \(\mathrm{XX'_3}\) | \(\mathrm{ClF_3}\) | Colorless gas |
| \(\mathrm{BrF_3}\) | Yellow green liquid | |
| \(\mathrm{IF_3}\) | Yellow powder | |
| \(\mathrm{ICl_3}\) | Orange solid; dimerises to form \(\mathrm{I_2Cl_6}\) having Cl-bridges | |
| \(\mathrm{XX'_5}\) | \(\mathrm{IF_5}\) | Colorless gas at R.T. but solid below 77 K |
| \(\mathrm{BrF_5}\) | Colorless liquid | |
| \(\mathrm{ClF_5}\) | Colorless liquid | |
| \(\mathrm{XX'_7}\) | \(\mathrm{IF_7}\) | Colorless gas |
Methods of Preparation
Table : Methods of preparation of interhalogen compounds
| Method | \(\mathrm{XX'}\) | \(\mathrm{XX'_3}\) |
|---|---|---|
| 1. Direct combination | Both in equal volumes: \(\mathrm{Cl_2} + \mathrm{F_2} \xrightarrow{\ 523\ \text{K},\ \text{Cu-vessel}\ } 2\,\mathrm{ClF}\); \(\mathrm{I_2} + \mathrm{Cl_2} \xrightarrow{\ \mathrm{HCl} + \mathrm{HNO_3},\ \text{boil}\ } 2\,\mathrm{ICl}\); \(\mathrm{Br_2} + \mathrm{Cl_2} \rightarrow 2\,\mathrm{BrCl}\) | \(\mathrm{Cl_2} + 3\,\mathrm{F_2}\ (\text{excess}) \rightarrow 2\,\mathrm{ClF_3}\); \(\mathrm{Br_2} + 3\,\mathrm{F_2} \rightarrow 2\,\mathrm{BrF_3}\); \(\mathrm{I_2} + 3\,\mathrm{Cl_2}\ (\text{excess}) \rightarrow 2\,\mathrm{ICl_3}\) |
| 2. Reaction of a halogen with an interhalogen compound | \(\mathrm{Br_2} + \mathrm{BrF_3} \rightarrow 3\,\mathrm{BrF}\) | \(\mathrm{Br_2} + \mathrm{ClF_3} \rightarrow 2\,\mathrm{BrF_3} + \mathrm{BrCl}\) |
| Special reaction for ICl | \(\mathrm{I_2} + \mathrm{KClO_3} \xrightarrow{\ \Delta\ } \mathrm{ICl} + \mathrm{KIO_3}\) | ——- |
Look closely at the first row and you will see how the product is controlled: equal volumes of the two halogens give the 1 : 1 compound, while an excess of the smaller halogen drives the reaction on to \(\mathrm{XX'_3}\).
Some Properties
Table : Some properties of interhalogen compounds
| Reaction/Property | \(\mathrm{XX'}\) | \(\mathrm{XX'_3}\) | |
|---|---|---|---|
| 1 | Thermal stability | \(\mathrm{ClF} > \mathrm{ICl} > \mathrm{IBr} > \mathrm{BrCl} > \mathrm{BrF}\) | |
| 2 | Hydrolysis gives oxoacids | \(\mathrm{BrCl} + \mathrm{H_2O} \rightarrow \mathrm{HOBr} + \mathrm{HCl}\); \(5\,\mathrm{ICl} + 3\,\mathrm{H_2O} \rightarrow \mathrm{HIO_3} + 5\,\mathrm{HCl} + 2\,\mathrm{I_2}\) | \(2\,\mathrm{ICl_3} + 3\,\mathrm{H_2O} \rightarrow \mathrm{HIO_3} + 5\,\mathrm{HCl} + \mathrm{ICl}\) |
| 3 | Disproportionation / autoionisation | \(5\,\mathrm{BrF} \rightarrow 2\,\mathrm{Br_2} + \mathrm{BrF_5}\) | \(2\,\mathrm{ClF_3} \rightarrow \mathrm{ClF_2^{+}} + \mathrm{ClF_4^{-}}\); \(\mathrm{ICl_3} \xrightarrow{\ 341\ \text{K}\ } \mathrm{ICl} + \mathrm{Cl_2}\) |
| 4 | Fluorination | ——— | \(\mathrm{U} + \mathrm{ClF_3} \rightarrow \mathrm{UF_6}\text{(l)} + \mathrm{ClF}\text{(g)}\) |
| 5 | Addition across olefins | \(\mathrm{H_2C}\!=\!\mathrm{CH_2} + \mathrm{ICl} \rightarrow \mathrm{ICH_2\!-\!CH_2Cl}\) |
In hydrolysis the more electropositive halogen ends up in the oxoacid and the more electronegative one leaves as the halide acid — which is exactly what the definition of an interhalogen compound predicts.
Liquid ICl also undergoes autoionization, like water, to form a cation and an anion:
\[2\,\mathrm{ICl} \rightleftharpoons \mathrm{I^{+}}\ (\text{solvent cation}) + \mathrm{ICl_2^{-}}\ (\text{solvent anion})\]
This is what makes some interhalogen compounds useful as nonaqueous solvents: like water, they can supply their own ions and dissolve salts.
Uses
Table : Uses of interhalogen compounds
| \(\mathrm{XX'}\) | \(\mathrm{XX'_3}\) |
|---|---|
| ICl is used to determine the iodine value of oils | For preparation of polyhalides |
| As catalyst for oxidation of As(III) | As fluorinating agent |
| For preparation of polyhalides | As nonaqueous solvent |
Structures
The shapes of these molecules follow straight from VSEPR theory, once the lone pairs on the central atom are counted. Table 7.11 lists the structures, which are drawn in Fig. 7.17.
Table 7.11 : Structures of some interhalogen compounds
| Formula | Name | Shape |
|---|---|---|
| ICl | Iodine monochloride | Linear |
| \(\mathrm{ClF_3}\) | Chlorine trifluoride | Bent T-shaped (bond angle 89°) |
| \(\mathrm{BrF_3}\) | Bromine trifluoride | Bent T-shaped (bond angle 86°) |
| \(\mathrm{BrF_5}\) | Bromine pentafluoride | Square pyramidal |
| \(\mathrm{ClF_5}\) | Chlorine pentafluoride | Square pyramidal |
Diagram
Fig. 7.17 – Five structures in a row with lone pairs shown on the central atom: ICl linear; ClF₃ bent T-shape with the 89° angle marked; BrF₃ bent T-shape with the 86° angle marked; BrF₅ square pyramidal with one lone pair above the Br; ClF₅ square pyramidal with one lone pair above the Cl.
Table 7.12 : Oxidation states of the central halogen atom in interhalogen compounds
| O.S. of central halogen | No. of lone pairs of electrons | Examples |
|---|---|---|
| +7 | 0 | \(\mathrm{IF_7}\) |
| +5 | 1 | \(\mathrm{ClF_5}\), \(\mathrm{BrF_5}\), \(\mathrm{IF_5}\) |
| +3 | 2 | \(\mathrm{ClF_3}\), \(\mathrm{BrF_3}\), \(\mathrm{IF_3}\), \(\mathrm{I_2Cl_6}\) |
| +1 | 3 | ClF, BrF, IF, BrCl, ICl, IBr |
Table 7.12 is more useful than it looks. Read it as a rule: each step down in oxidation state adds one lone pair to the central atom, because the two electrons that are no longer used in bonding stay behind as a pair. Once you know the oxidation state, you know the number of lone pairs, and once you know that, VSEPR gives you the shape.
7.13 Compounds of Xenon
We have studied earlier in this chapter that group 18 elements have very high ionisation energies, and due to this property they are unreactive. Each noble gas atom has a completely filled valence electron shell, which makes it inert.
The first ionization potential decreases down the group. Hence the heavier noble gases Kr, Xe and Rn can form compounds, because of their low ionization energy.
Only xenon reacts directly with fluorine to form xenon fluorides, and the stable fluorides are \(\mathrm{XeF_2}\), \(\mathrm{XeF_4}\) and \(\mathrm{XeF_6}\). Xenon also forms compounds with oxygen, such as \(\mathrm{XeO_3}\), \(\mathrm{XeOF_2}\), \(\mathrm{XeOF_4}\) and \(\mathrm{XeO_2F_2}\).
\[\mathrm{Xe} + 2\,\mathrm{[PtF_6]} \rightleftharpoons \mathrm{[XeF]^{+}[Pt_2F_{11}]^{-}}\]
Sixty years of textbooks calling this group the “inert gases” ended with that experiment, and the family was renamed the noble gases instead.
7.13.1 Xenon Fluorides
a. Preparation of Xenon Fluorides
Xenon fluorides are generally prepared by the direct reaction of xenon and fluorine in different ratios and conditions, such as temperature, electric discharge and photochemical reaction.
\[\mathrm{Xe} + \mathrm{F_2} \xrightarrow{\ \text{sealed Ni tube},\ 400\ ^\circ\text{C}\ } \mathrm{XeF_2}\]
\[\mathrm{Xe} + \mathrm{F_2}\ (1:5) \xrightarrow{\ \text{Ni tube},\ 400\ ^\circ\text{C},\ 5\text{–}6\ \text{atm}\ } \mathrm{XeF_4}\]
\[\mathrm{Xe} + 2\,\mathrm{F_2} \xrightarrow{\ \text{electric discharge},\ -80\ ^\circ\text{C}\ } \mathrm{XeF_4}\]
\[\mathrm{Xe} + 3\,\mathrm{F_2} \xrightarrow{\ \text{electric discharge},\ \text{low temp.}\ } \mathrm{XeF_6}\]
The pattern is worth reading as a whole: the more fluorine that is supplied, and the harsher the conditions, the further the reaction is pushed — \(\mathrm{XeF_2}\), then \(\mathrm{XeF_4}\), then \(\mathrm{XeF_6}\). Controlling the ratio of the two gases is what controls the product.
b. Important chemical reactions of \(\mathrm{XeF_2}\)
i. Hydrolysis: \(\mathrm{XeF_2}\) undergoes hydrolysis to form HF.
\[2\,\mathrm{XeF_2} + 2\,\mathrm{H_2O} \rightarrow 4\,\mathrm{HF} + 2\,\mathrm{Xe} + \mathrm{O_2}\]
Water pulls the compound apart completely here, releasing the xenon as free gas. This is why xenon fluorides must be prepared and stored under strictly dry conditions.
ii. Reaction with \(\mathrm{PF_5}\): \(\mathrm{XeF_2}\) forms adducts on reaction with \(\mathrm{PF_5}\).
\[\mathrm{XeF_2} + \mathrm{PF_5} \rightarrow \mathrm{XeF_2{\cdot}PF_5}\]
7.13.2 Xenon Trioxide
a. Preparation of xenon trioxide (\(\mathrm{XeO_3}\))
Fluorides of xenon react with water to form \(\mathrm{XeO_3}\).
\[3\,\mathrm{XeF_4} + 6\,\mathrm{H_2O} \rightarrow 2\,\mathrm{Xe} + \mathrm{XeO_3} + 12\,\mathrm{HF} + 1\tfrac{1}{2}\,\mathrm{O_2}\]
\[\mathrm{XeF_6} + 3\,\mathrm{H_2O} \rightarrow \mathrm{XeO_3} + 6\,\mathrm{HF}\]
Notice the difference between the two. Complete hydrolysis of \(\mathrm{XeF_6}\) replaces every fluorine by oxygen in a single clean step, while \(\mathrm{XeF_4}\) partly falls apart to free xenon at the same time.
7.13.3 Oxyfluorides of Xenon
Xenon forms the following oxyfluorides: \(\mathrm{XeOF_2}\), \(\mathrm{XeO_2F_2}\), \(\mathrm{XeO_3F_2}\), \(\mathrm{XeO_2F_4}\) and \(\mathrm{XeOF_4}\).
a. Preparation
i. Partial hydrolysis of xenon fluorides yields different oxyfluorides.
\[\mathrm{XeF_4} + \mathrm{H_2O} \xrightarrow{\ 80\ ^\circ\text{C}\ } \mathrm{XeOF_2} + 2\,\mathrm{HF}\]
\[\mathrm{XeF_6} + \mathrm{H_2O} \rightarrow \mathrm{XeOF_4} + 2\,\mathrm{HF}\]
The word partial is doing the work here. One molecule of water replaces just two fluorine atoms by one oxygen atom; with three molecules of water the replacement goes all the way and \(\mathrm{XeO_3}\) is obtained instead, as in section 7.13.2.
ii. Reaction of a xenon oxyfluoride with \(\mathrm{SiO_2}\), or its hydrolysis, yields xenon dioxydifluoride.
\[2\,\mathrm{XeOF_4} + \mathrm{SiO_2} \rightarrow 2\,\mathrm{XeO_2F_2} + \mathrm{SiF_4}\]
\[\mathrm{XeOF_4} + \mathrm{H_2O} \rightarrow \mathrm{XeO_2F_2} + 2\,\mathrm{HF}\]
The reaction with silica is a warning as much as a preparation: it means xenon oxyfluorides attack ordinary glassware, and so must be handled in nickel or fluoropolymer vessels. The whole hydrolysis family is collected in FC 7.7.
Flowchart
FC 7.7 – Hydrolysis routes of xenon fluorides: XeF₆ → with 1 H₂O (partial) → XeOF₄ → with more H₂O or with SiO₂ → XeO₂F₂ → further hydrolysis → XeO₃; parallel branch: XeF₆ + 3H₂O (complete) → XeO₃ + 6HF; second branch: XeF₄ + H₂O at 80 °C → XeOF₂; third branch: XeF₄ + excess H₂O → Xe + XeO₃ + HF + O₂; fourth branch: XeF₂ + H₂O → Xe + HF + O₂ (complete breakdown, no oxide formed).
| Formula | Name |
|---|---|
| \(\mathrm{XeOF_2}\) | Xenon mono oxyfluoride |
| …… | Xenon dioxydifluoride |
| \(\mathrm{XeO_3F_2}\) | …… |
| \(\mathrm{XeO_2F_4}\) | …… |
Structures of the Xenon Compounds
The shapes follow from VSEPR theory once the lone pairs left on xenon are counted, exactly as they did for the interhalogen compounds. The structures are drawn in Fig. 7.18.
Table 7.13 : Structure of xenon compounds
| Sr. No. | Formula | Oxidation state of Xe | Structure |
|---|---|---|---|
| 1. | \(\mathrm{XeF_2}\) | +2 | Linear |
| 2. | \(\mathrm{XeF_4}\) | +4 | Square planar |
| 3. | \(\mathrm{XeF_6}\) | +6 | Distorted octahedral |
| 4. | \(\mathrm{XeO_3}\) | +6 | Trigonal pyramidal |
| 5. | \(\mathrm{XeOF_4}\) | +6 | Square pyramidal |
Diagram
Fig. 7.18 – Five structures in a row with lone pairs on xenon shown: XeF₂ linear with three lone pairs in the equatorial plane; XeF₄ square planar with one lone pair above and one below; XeF₆ distorted octahedral with a lone pair pushing one face out of shape; XeO₃ trigonal pyramidal with one lone pair on top; XeOF₄ square pyramidal with the oxygen at the apex and a lone pair below.
7.14 Summary
In the p-block elements the differentiating electron enters a p-orbital of the outermost shell. A p-subshell holds a maximum of six electrons, which is why the p-block is six columns wide, giving groups 13 to 18.
Group 16 (O, S, Se, Te, Po) is the oxygen family, also called the chalcogens or ore-forming elements. Group 17 (F, Cl, Br, I, At) are the halogens, or salt producers. Group 18 (He, Ne, Ar, Kr, Xe, Rn) are the noble gases.
Oxygen is the most abundant element on earth — 20.95 % by volume of air and about 46.6 % by mass of the crust. Sulfur (0.034 % by mass of the crust) occurs as sulfates such as gypsum, epsom salt and baryte, and as sulfides such as galena, zinc blende and copper pyrites.
Halogens are too reactive to occur free. Fluorine occurs as fluorspar, cryolite and fluorapatite; sea water supplies chloride, bromide and iodide, and contains 2.5 % NaCl by mass. Astatine is radioactive with a half life of 8.1 hours.
All noble gases except radon occur in the atmosphere, about 1 % by volume of dry air, with argon as the major constituent. Helium comes commercially from natural gas.
General valence configurations: group 16 is \(\mathrm{ns^2np^4}\), group 17 is \(\mathrm{ns^2np^5}\), group 18 is \(\mathrm{ns^2np^6}\). Groups 16 and 17 are two and one electrons short of the nearest noble gas respectively.
Atomic and ionic radii increase down each group as quantum shells are added, and decrease across a period as effective nuclear charge rises. Halogens are the smallest atoms in their periods.
Ionisation enthalpy is high for all three groups; it decreases down a group with increasing size and increases across a period. Group 16 values fall below those of group 15 in the same period because the half filled \(\mathrm{ns^2np^3}\) configuration of group 15 is extra stable.
Electronegativity decreases down a group. Fluorine is the most electronegative element, and oxygen is next after it.
Electron gain enthalpy becomes less negative down groups 16 and 17. Oxygen is less negative than sulfur, and fluorine less negative than chlorine, because the first member is so small that the incoming electron meets strong repulsion. Noble gases have large positive values.
Physically: oxygen is a gas and the rest of group 16 are solids; O and S are nonmetals, Se and Te metalloids, Po a metal. In group 17, \(\mathrm{F_2}\) and \(\mathrm{Cl_2}\) are gases, \(\mathrm{Br_2}\) a liquid and \(\mathrm{I_2}\) a solid. Noble gases are monoatomic with very low melting and boiling points; helium has the lowest boiling point known, 4.2 K.
Bond dissociation enthalpy of the halogens runs \(\mathrm{Cl\!-\!Cl} > \mathrm{Br\!-\!Br} > \mathrm{F\!-\!F} > \mathrm{I\!-\!I}\). \(\mathrm{F_2}\) is out of place because lone pair–lone pair repulsion in the tiny molecule weakens the bond.
Oxygen behaves anomalously: it is diatomic and paramagnetic, shows only −2, −1 and +2 oxidation states, its hydride is a liquid, and its covalency is limited to 2. The causes are its small size, high electronegativity and absence of d-orbitals.
Fluorine behaves anomalously for the same three reasons together with its low F–F bond dissociation enthalpy. It shows only the −1 state, forms just one oxoacid (HOF), reacts exothermically, and HF is a liquid because of hydrogen bonding.
Group 16 shows −2 and +2 states, and also +4 and +6 for the heavier members; the +6 state becomes less stable and +4 more stable down the group because of the inert pair effect. Halogens all show −1, while Cl, Br and I also show +1, +3, +5 and +7 by expanding the octet into empty d orbitals.
Xenon is the noble gas that reacts, because it is large and its ionisation enthalpy is low enough for paired electrons to be promoted into empty d orbitals.
Hydrides \(\mathrm{H_2E}\) are angular and \(\mathrm{sp^3}\) hybridised. Down the group the H–E bond weakens, so acidic character and reducing power increase while thermal stability decreases, from \(\mathrm{H_2O}\) to \(\mathrm{H_2Te}\).
Hydrogen halides follow the same logic: acid strength increases \(\mathrm{HF} < \mathrm{HCl} < \mathrm{HBr} < \mathrm{HI}\) while thermal stability decreases in the reverse order, both because the H–X bond dissociation enthalpy falls down the group.
Group 16 forms acidic oxides of types \(\mathrm{EO_2}\) and \(\mathrm{EO_3}\) that give acids with water. \(\mathrm{SO_2}\) is reducing while \(\mathrm{TeO_2}\) is oxidising. Halogen oxides are mostly unstable, and the higher oxides are more stable than the lower ones.
Group 16 forms hexahalides (\(\mathrm{sp^3d^2}\), octahedral), tetrahalides (\(\mathrm{sp^3d}\), trigonal bipyramidal with one lone pair) and dihalides (\(\mathrm{sp^3}\), two lone pairs). Stability falls in the order fluorides > chlorides > bromides > iodides.
Group 16 elements combine with metals to give oxides, sulfides and selenides; halogens give metal halides whose ionic character falls in the order \(\mathrm{MF} > \mathrm{MCl} > \mathrm{MBr} > \mathrm{MI}\), and which become more covalent as the oxidation state of the metal rises.
All group 16 elements show allotropy. Oxygen has \(\mathrm{O_2}\) and \(\mathrm{O_3}\); sulfur has rhombic (\(\alpha\)) and monoclinic (\(\beta\)) forms, both built from puckered \(\mathrm{S_8}\) rings and interconverting at 369 K. Allotropy applies to elements, polymorphism to compounds.
Sulfur forms the oxoacids \(\mathrm{H_2SO_3}\), \(\mathrm{H_2SO_4}\), \(\mathrm{H_2S_2O_7}\), \(\mathrm{H_2SO_5}\), \(\mathrm{H_2S_2O_8}\) and \(\mathrm{H_2S_2O_3}\); several exist only in solution or as salts.
Halogen oxoacid strength increases with the oxidation state of the halogen, from the weak \(\mathrm{HClO}\) (\(K_a = 3.5 \times 10^{-8}\)) to the very strong \(\mathrm{HClO_4}\).
Dioxygen is prepared in the laboratory from chlorates, from metal oxides and from hydrogen peroxide, and industrially by the fractional distillation of liquefied air. It is colourless, sparingly soluble in water, and paramagnetic — a fact only molecular orbital theory explains, through two unpaired electrons in \(\pi^*\) orbitals.
Oxides are classified as acidic (mostly of nonmetals), basic (of metals), amphoteric (such as \(\mathrm{Al_2O_3}\)) and neutral (CO, NO, \(\mathrm{N_2O}\)).
Ozone is prepared by silent electric discharge through dry oxygen in an ozoniser, only 10 % conversion being achieved: \(3\,\mathrm{O_2} \rightleftharpoons 2\,\mathrm{O_3}\), \(\Delta H = +142\ \text{kJ mol}^{-1}\). It is angular, a resonance hybrid with equal O–O bond lengths of 128 pm and a bond angle of about 117°.
Ozone is a powerful oxidising agent because it releases nascent oxygen; it bleaches without moisture (dry bleach) and reduces peroxides to oxides. Its layer shields the earth from UV radiation, and it is depleted by nitrogen oxides and by chlorofluorocarbons.
Sulfur dioxide is prepared by burning sulfur, by acidifying a sulfite, and industrially by roasting sulfide ores. It is angular (119.5°), a resonance hybrid with \(\mathrm{d\pi\!-\!p\pi}\) bonding, and a reducing agent in moist conditions.
Sulfuric acid is manufactured by the contact process: \(\mathrm{SO_2}\) is oxidised over \(\mathrm{V_2O_5}\) at 720 K and 2 bar, the \(\mathrm{SO_3}\) is absorbed in concentrated acid to give oleum, and the oleum is diluted. The product is 96–98 % pure.
\[2\,\mathrm{SO_2}\text{(g)} + \mathrm{O_2} \xrightarrow{\ \mathrm{V_2O_5}\ } 2\,\mathrm{SO_3}\text{(g)}\]
Sulfuric acid is a strong dibasic acid, an oxidising agent when hot and concentrated, a powerful dehydrating agent that chars sugar, and it displaces more volatile acids from their salts.
Chlorine, discovered by Scheele in 1774 and named by Davy in 1810, is prepared by oxidising HCl with \(\mathrm{MnO_2}\) or \(\mathrm{KMnO_4}\), and manufactured by Deacon’s process and by the electrolysis of brine.
Chlorine reacts with metals, nonmetals, hydrogen, ammonia, alkalis and hydrocarbons. With cold dilute alkali it gives chloride and hypochlorite; with hot concentrated alkali, chloride and chlorate; with dry slaked lime, bleaching powder. It bleaches only in the presence of moisture, through nascent oxygen.
Hydrogen chloride, first made by Glauber in 1648, is prepared by heating NaCl with concentrated \(\mathrm{H_2SO_4}\) and dried over the same acid. Its aqueous solution is a strong acid, and mixed 3 : 1 with concentrated \(\mathrm{HNO_3}\) it forms aqua regia, which dissolves gold and platinum.
Interhalogen compounds \(\mathrm{XX'_n}\) form because halogens differ in electronegativity; X is the larger, more electropositive atom and \(n\) is 1, 3, 5 or 7. They are diamagnetic, more reactive than the parent halogens, and hydrolyse to oxoacids.
Their shapes follow VSEPR: \(\mathrm{XX'}\) linear, \(\mathrm{XX'_3}\) bent T-shaped, \(\mathrm{XX'_5}\) square pyramidal, \(\mathrm{XX'_7}\) pentagonal bipyramidal, with the number of lone pairs rising as the oxidation state of the central halogen falls.
The first noble gas compound was made by Neil Bartlett in 1962. Xenon forms \(\mathrm{XeF_2}\), \(\mathrm{XeF_4}\) and \(\mathrm{XeF_6}\) by direct combination with fluorine under different ratios and conditions, and these hydrolyse to \(\mathrm{XeO_3}\) and, partially, to oxyfluorides such as \(\mathrm{XeOF_4}\) and \(\mathrm{XeO_2F_2}\).
Xenon compound shapes: \(\mathrm{XeF_2}\) linear, \(\mathrm{XeF_4}\) square planar, \(\mathrm{XeF_6}\) distorted octahedral, \(\mathrm{XeO_3}\) trigonal pyramidal, \(\mathrm{XeOF_4}\) square pyramidal.
Helium, neon and argon are used wherever chemical inertness or extreme cold is needed — breathing mixtures, welding atmospheres, discharge lamps, filament bulbs, lasers, cryogenic research and magnetic resonance imaging.
Flowchart
FC 7.8 – Chapter overview: valence configuration (ns²np⁴ / ns²np⁵ / ns²np⁶) → decides electrons short of the octet → decides atomic properties (radius, ionisation enthalpy, electronegativity, electron gain enthalpy) → decides oxidation states, with the branch “first member: no d orbitals → anomalous behaviour of O and F” → leads to compound classes: hydrides H₂E and HX → oxides and oxoacids → halides and interhalogens → xenon fluorides, oxides and oxyfluorides → industrial routes: contact process for H₂SO₄, Deacon’s process and brine electrolysis for Cl₂.
7.15 Objective Questions
- Which of the following has highest electron gain enthalpy?
- Fluorine
- Chlorine
- Bromine
- Iodine
- Hydrides of group 16 are weakly acidic. The correct order of acidity is
- \(\mathrm{H_2O} > \mathrm{H_2S} > \mathrm{H_2Se} > \mathrm{H_2Te}\)
- \(\mathrm{H_2Te} > \mathrm{H_2O} > \mathrm{H_2S} > \mathrm{H_2Se}\)
- \(\mathrm{H_2Te} > \mathrm{H_2Se} > \mathrm{H_2S} > \mathrm{H_2O}\)
- \(\mathrm{H_2Te} > \mathrm{H_2Se} > \mathrm{H_2O} > \mathrm{H_2S}\)
- Which of the following elements does not show oxidation state of +4?
- O
- S
- Se
- Te
- HI acid when heated with conc. \(\mathrm{H_2SO_4}\) forms
- \(\mathrm{HIO_3}\)
- \(\mathrm{KIO_3}\)
- \(\mathrm{I_2}\)
- KI
- Ozone layer is depleted by
- NO
- \(\mathrm{NO_2}\)
- \(\mathrm{NO_3}\)
- \(\mathrm{N_2O_5}\)
- Which of the following occurs in liquid state at room temperature?
- \(\mathrm{HIO_3}\)
- HBr
- HCl
- HF
- In pyrosulfurous acid the oxidation state of sulfur is
- Only +2
- Only +4
- +2 and +6
- Only +6
- Stability of interhalogen compounds follows the order
- \(\mathrm{BrF} > \mathrm{IBr} > \mathrm{ICl} > \mathrm{ClF} > \mathrm{BrCl}\)
- \(\mathrm{IBr} > \mathrm{BeF} > \mathrm{ICl} > \mathrm{ClF} > \mathrm{BrCl}\)
- \(\mathrm{ClF} > \mathrm{ICl} > \mathrm{IBr} > \mathrm{BrCl} > \mathrm{BrF}\)
- \(\mathrm{ICl} > \mathrm{ClF} > \mathrm{BrCl} > \mathrm{IBr} > \mathrm{BrF}\)
- BrCl reacts with water to form
- HBr
- \(\mathrm{Br_2} + \mathrm{Cl_2}\)
- HOBr
- HOBr + HCl
- Chlorine reacts with excess of fluorine to form
- ClF
- \(\mathrm{ClF_3}\)
- \(\mathrm{ClF_2}\)
- \(\mathrm{Cl_2F_3}\)
- In interhalogen compounds, which of the following halogens is never the central atom?
- I
- Cl
- Br
- F
- Which of the following has one lone pair of electrons?
- \(\mathrm{IF_3}\)
- ICl
- \(\mathrm{IF_5}\)
- \(\mathrm{ClF_3}\)
- In which of the following pairs are molecules paired with their correct shapes?
- \(\mathrm{[I_3]}\) : bent
- \(\mathrm{BrF_5}\) : trigonal bipyramid
- \(\mathrm{ClF_3}\) : trigonal planar
- \(\mathrm{[BrF_4]}\) : square planar
- Among the known interhalogen compounds, the maximum number of atoms is
- 3
- 6
- 7
- 8
7.16 Check Your Concepts
Assertion: The first ionisation enthalpy of sulfur is lower than that of phosphorus. Reason: Sulfur has a larger atomic radius than phosphorus. State whether the assertion and the reason are true, and whether the reason correctly explains the assertion.
Why does the electron gain enthalpy of fluorine turn out to be less negative than that of chlorine, even though fluorine is the more electronegative of the two?
Arrange \(\mathrm{F_2}\), \(\mathrm{Cl_2}\), \(\mathrm{Br_2}\) and \(\mathrm{I_2}\) in decreasing order of bond dissociation enthalpy. Justify the position of \(\mathrm{F_2}\), and explain why \(\mathrm{I_2}\) occupies its position for an entirely different reason.
Oxygen and sulfur belong to the same group, yet \(\mathrm{H_2O}\) is a liquid and \(\mathrm{H_2S}\) is a gas at room temperature. Account for this difference, and predict where \(\mathrm{H_2Se}\) would fit.
Explain why sulfur forms \(\mathrm{SF_6}\) while oxygen forms only \(\mathrm{OF_2}\) with fluorine. Which species out of \(\mathrm{OF_6}\), \(\mathrm{SeF_6}\) and \(\mathrm{OCl_4}\) cannot exist, and why?
What happens if chlorine is passed into (a) cold dilute NaOH and (b) hot concentrated NaOH? Write the balanced equation in each case and state why both are called disproportionation reactions.
Complete the following reactions and name the main product in each case:
- \(\mathrm{SeO_3} + \mathrm{H_2O} \rightarrow\)
- \(2\,\mathrm{Ca(OH)_2} + 2\,\mathrm{Cl_2} \rightarrow\)
- \(\mathrm{Na_2SO_3} + \mathrm{H_2O} + \mathrm{SO_2} \rightarrow\)
- \(\mathrm{XeF_2} + \mathrm{PF_5} \rightarrow\)
Assertion: Ozone acts as a powerful oxidising agent. Reason: Ozone is thermodynamically more stable than dioxygen. Comment on the truth of each statement and on whether the reason explains the assertion.
Ozone is called a dry bleach while chlorine bleaches only in the presence of moisture. Explain the difference using the chemistry of each bleaching action.
Why is concentrated sulfuric acid always added slowly to water, and never water to the acid? Explain what would physically happen if the order were reversed.
Concentrated \(\mathrm{H_2SO_4}\) can be used to prepare HCl from NaCl but not HBr from NaBr. Explain, writing the equation that shows what goes wrong in the bromide case.
Arrange HF, HCl, HBr and HI in increasing order of (a) acid strength and (b) thermal stability. Explain why the two orders are the reverse of each other and why one single property accounts for both.
State what you would observe, and write the equation, when moist \(\mathrm{SO_2}\) is passed through acidified potassium permanganate solution. Which property of \(\mathrm{SO_2}\) does this demonstrate?
In the contact process, the oxidation of \(\mathrm{SO_2}\) to \(\mathrm{SO_3}\) is favoured by low temperature and high pressure. Explain both conditions from the equation, and state why in practice the temperature cannot be lowered indefinitely.
Why is sulfur trioxide absorbed in concentrated sulfuric acid rather than in water during the manufacture of sulfuric acid?
Predict the shape and hybridisation of each of the following, showing the count of bond pairs and lone pairs: \(\mathrm{ClF_3}\), \(\mathrm{BrF_5}\), \(\mathrm{XeF_4}\) and \(\mathrm{XeO_3}\).
Why is \(\mathrm{XeF_6}\) described as distorted octahedral while \(\mathrm{SF_6}\) is a regular octahedron?
Explain why interhalogen compounds are more reactive than the halogens from which they are made, and why the number of atoms in an interhalogen molecule is always even.
Why does iodine form \(\mathrm{IF_7}\) while chlorine forms no compound beyond \(\mathrm{ClF_5}\)? Refer to the ratio of the radii of the two halogens in your answer.
Match the following.
| A | B |
|---|---|
| a. \(\mathrm{I_2O_5}\) | i. Bleaching agent for paper pulp |
| b. \(\mathrm{ClO_2}\) | ii. Estimation of carbon monoxide |
| c. ICl in glacial acetic acid | iii. Removal of plutonium from spent nuclear fuel |
| d. \(\mathrm{O_2F_2}\) | iv. Determination of iodine value of an oil |
Helium and xenon are both noble gases, yet only xenon forms compounds. Give two reasons, and explain why fluorine and oxygen are the only partners xenon accepts.
Both rhombic and monoclinic sulfur are built from \(\mathrm{S_8}\) rings. Then why do they differ in colour, density and melting point? What happens to a sample of rhombic sulfur kept at 380 K?
Explain why (a) polonium has no ore of its own, (b) astatine cannot be collected in weighable amounts, and (c) helium is obtained from natural gas rather than from air.
The stability of the +6 oxidation state falls down group 16 while that of the +4 state rises. Name the effect responsible and predict, with reasoning, which of \(\mathrm{SO_2}\) and \(\mathrm{TeO_2}\) acts as a reducing agent and which as an oxidising agent.
Write the structural difference between peroxy monosulfuric acid and pyrosulfuric acid, and explain how that difference changes the way the oxidation state of sulfur must be calculated in each.
Write the order of the thermal stability of the hydrides of group 16 elements.
Name two gases which deplete the ozone layer.
Give two uses of \(\mathrm{ClO_2}\).
What is the action of bromine on magnesium metal?
Write the names of the allotropic forms of selenium.
The \(\mathrm{p}K_a\) value of HCl is −7.0 and that of HI is −10.0. Which is the stronger acid?
Give one example showing the reducing property of ozone.
Write the reaction of conc. \(\mathrm{H_2SO_4}\) with sugar.
Give two uses of chlorine.
Complete the following.
- \(\mathrm{ICl_3} + \mathrm{H_2O} \rightarrow\) …… + …… + ICl
- \(\mathrm{I_2} + \mathrm{KClO_3} \rightarrow\) …… + \(\mathrm{KIO_2}\)
- \(\mathrm{BrCl} + \mathrm{H_2O} \rightarrow\) …… + HCl
- \(\mathrm{Cl_2} + \mathrm{ClF_3} \rightarrow\) ……
- \(\mathrm{H_2C}\!=\!\mathrm{CH_2} + \mathrm{ICl} \rightarrow\) ……
- \(\mathrm{XeF_4} + \mathrm{SiO_2} \rightarrow\) …… + \(\mathrm{SiF_4}\)
- \(\mathrm{XeF_6} + 6\,\mathrm{H_2O} \rightarrow\) …… + HF
- \(\mathrm{XeOF_4} + \mathrm{H_2O} \rightarrow\) …… + HF
- Match the following.
| A | B |
|---|---|
| a. \(\mathrm{XeOF_2}\) | i. Xenon trioxydifluoride |
| b. \(\mathrm{XeO_2F_2}\) | ii. Xenon monooxydifluoride |
| c. \(\mathrm{XeO_3F_2}\) | iii. Xenon dioxytetrafluoride |
| d. \(\mathrm{XeO_2F_4}\) | iv. Xenon dioxydifluoride |
The first ionisation enthalpies of S, Cl and Ar are 1000, 1256 and 1520 \(\text{kJ mol}^{-1}\) respectively. Explain the observed trend.
“Acidic character of hydrides of group 16 elements increases from \(\mathrm{H_2O}\) to \(\mathrm{H_2Te}\).” Explain.
How is dioxygen prepared in the laboratory from \(\mathrm{KClO_3}\)?
What happens when (a) lead sulfide reacts with ozone (\(\mathrm{O_3}\)), (b) nitric oxide reacts with ozone?
Give two chemical reactions to explain the oxidizing property of concentrated \(\mathrm{H_2SO_4}\).
Discuss the structure of sulfur dioxide.
Fluorine shows only −1 oxidation state while other halogens show −1, +1, +3, +5 and +7 oxidation states. Explain.
What is the action of chlorine on the following: (a) Fe (b) excess of \(\mathrm{NH_3}\)?
How is hydrogen chloride prepared from sodium chloride?
Draw the structures of \(\mathrm{XeF_6}\), \(\mathrm{XeO_3}\), \(\mathrm{XeOF_4}\) and \(\mathrm{XeF_2}\).
What are inter-halogen compounds? Give two examples.
What is the action of hydrochloric acid on the following: (a) \(\mathrm{NH_3}\) (b) \(\mathrm{Na_2CO_3}\)?
Give two uses of HCl.
Write the names and structural formulae of the oxoacids of chlorine.
What happens when (a) \(\mathrm{Cl_2}\) reacts with \(\mathrm{F_2}\) in equal volume at 437 K, (b) \(\mathrm{Br_2}\) reacts with excess of \(\mathrm{F_2}\)?
How are xenon fluorides \(\mathrm{XeF_2}\), \(\mathrm{XeF_4}\) and \(\mathrm{XeF_6}\) obtained? Give suitable reactions.
How are \(\mathrm{XeO_3}\) and \(\mathrm{XeOF_4}\) prepared?
Give two uses of neon and argon.
Describe the structure of ozone. Give two uses of ozone.
Explain the trend in the following atomic properties of group 16 elements: (i) atomic radii (ii) ionisation enthalpy (iii) electronegativity.
Distinguish between rhombic sulfur and monoclinic sulfur.
How is \(\mathrm{SO_2}\) prepared in the laboratory from sodium sulfite? Give two physical properties of \(\mathrm{SO_2}\).
Describe the manufacturing of \(\mathrm{H_2SO_4}\) by the contact process.
7.17 Practice With Numericals
Calculate the volume of dioxygen, measured at STP, liberated by the complete decomposition of 12.26 g of potassium chlorate: \(2\,\mathrm{KClO_3} \rightarrow 2\,\mathrm{KCl} + 3\,\mathrm{O_2}\). (Molar masses in \(\text{g mol}^{-1}\): K = 39.1, Cl = 35.5, O = 16.0; molar volume at STP = 22.4 L \(\text{mol}^{-1}\))
3.20 kg of sulfur is converted to sulfuric acid by the contact process. Assuming complete conversion at every stage, calculate the mass of \(\mathrm{H_2SO_4}\) obtained. (Molar masses in \(\text{g mol}^{-1}\): S = 32.06, H = 1.008, O = 16.00)
The bond dissociation enthalpy of \(\mathrm{F_2}\) is 158.8 \(\text{kJ mol}^{-1}\). Calculate the energy required to dissociate all the molecules in 3.80 g of fluorine gas into atoms. (Molar mass of F = 19.00 \(\text{g mol}^{-1}\))
Calculate the percentage of sulfur by mass in copper pyrites, \(\mathrm{CuFeS_2}\). (Molar masses in \(\text{g mol}^{-1}\): Cu = 63.5, Fe = 55.85, S = 32.06)
Determine the oxidation state of the central atom in each of the following: \(\mathrm{HIO_3}\), \(\mathrm{KBrO_3}\), \(\mathrm{XeO_2F_2}\), \(\mathrm{H_2S_2O_7}\) and \(\mathrm{ClF_5}\).
15.80 g of potassium permanganate reacts with excess concentrated hydrochloric acid according to \(2\,\mathrm{KMnO_4} + 16\,\mathrm{HCl} \rightarrow 2\,\mathrm{KCl} + 2\,\mathrm{MnCl_2} + 8\,\mathrm{H_2O} + 5\,\mathrm{Cl_2}\). Calculate the volume of chlorine obtained at STP. (Molar masses in \(\text{g mol}^{-1}\): K = 39.1, Mn = 54.94, O = 16.00; molar volume at STP = 22.4 L \(\text{mol}^{-1}\))
0.30 mol of \(\mathrm{XeF_4}\) is hydrolysed completely according to \(3\,\mathrm{XeF_4} + 6\,\mathrm{H_2O} \rightarrow 2\,\mathrm{Xe} + \mathrm{XeO_3} + 12\,\mathrm{HF} + 1\tfrac{1}{2}\,\mathrm{O_2}\). Calculate (i) the moles of HF formed, (ii) the moles of free xenon released and (iii) the volume of \(\mathrm{O_2}\) obtained at STP. (Molar volume at STP = 22.4 L \(\text{mol}^{-1}\))
178.14 g of oleum, \(\mathrm{H_2S_2O_7}\), is diluted with the required quantity of water: \(\mathrm{H_2S_2O_7} + \mathrm{H_2O} \rightarrow 2\,\mathrm{H_2SO_4}\). Calculate the mass of sulfuric acid produced and the mass of water consumed. (Molar masses in \(\text{g mol}^{-1}\): H = 1.008, S = 32.06, O = 16.00)
6.41 g of sulfur dioxide is passed into acidified potassium permanganate solution: \(2\,\mathrm{KMnO_4} + 5\,\mathrm{SO_2} + 2\,\mathrm{H_2O} \rightarrow \mathrm{K_2SO_4} + 2\,\mathrm{MnSO_4} + 2\,\mathrm{H_2SO_4}\). Calculate the mass of \(\mathrm{KMnO_4}\) required for complete reaction. (Molar masses in \(\text{g mol}^{-1}\): S = 32.06, O = 16.00, K = 39.1, Mn = 54.94)
At 293 K, 30.8 cm³ of dioxygen dissolves in 1000 cm³ of water. Calculate the mass of dissolved dioxygen present in a 5.00 L aquarium filled with water saturated at this temperature. (Molar volume of a gas at 293 K and 1 atm = 24.0 L \(\text{mol}^{-1}\); molar mass of O = 16.00 \(\text{g mol}^{-1}\))
A sample of gypsum, \(\mathrm{CaSO_4{\cdot}2H_2O}\), weighing 8.61 g is heated until all the water of crystallisation is driven off. Calculate the mass of the residue obtained. (Molar masses in \(\text{g mol}^{-1}\): Ca = 40.08, S = 32.06, O = 16.00, H = 1.008)
What is the oxidation state of Te in \(\mathrm{TeO_3}\)?
What is the oxidation state of S in \(\mathrm{H_2SO_4}\)?
What is the oxidation state of xenon in the following compounds: \(\mathrm{XeOF_4}\), \(\mathrm{XeO_3}\), \(\mathrm{XeF_6}\), \(\mathrm{XeF_4}\), \(\mathrm{XeF_2}\)?