Chapter 8: Elements of Groups 1 and Group 2
In the first decade of the 1800s, a young English chemist named Humphry Davy was working at the Royal Institution in London with a tool that had only recently become available — the voltaic pile, an early form of battery. Chemists of that time strongly suspected that potash and soda, the alkalis used for centuries in soap-making, had unknown metals hidden inside them. But no furnace was hot enough, and no known reagent strong enough, to pull those metals out.
Davy tried something no furnace could do. He melted the alkali and passed an electric current through it.
At the negative electrode, tiny bright globules of a soft, silvery metal began to appear. Some of them caught fire the moment they met the moist air of the room. In this way potassium and sodium were separated as metals for the first time, and within about a year Davy had used related electrical methods to obtain calcium, magnesium, strontium and barium as well.
Those small burning globules settled a long-standing puzzle. These metals had never been dug out of the ground as metals because they cannot survive there. They hold their outermost electrons so loosely that air, water, or almost any other substance takes those electrons away almost at once — which is exactly why electricity, and not heat, was needed to set them free.
Diagram
Fig. 8.0 – Opening illustration: a voltaic pile wired to a dish of molten alkali, bright metal globules forming at the negative electrode with one globule flaring on contact with moist air; alongside, a small periodic-table strip with group 1 and group 2 (the s-block) shaded and hydrogen shown standing apart above group 1.
As you work through this chapter, you will collect the ideas needed to reason out each of these questions for yourself.
8.1 Hydrogen
In chapter 7 you saw that the elements of group 1 and group 2 belong to the s-block of the modern periodic table. The name comes from the subshell that receives the last electron: in every one of these elements, the outermost electron enters an ns subshell.
Hydrogen is the first element of the periodic table, and it is printed at the top of group 1, above the alkali metals. Yet it differs from those metals in so many ways that chemists study it on its own before turning to the s-block metals as a family. That is exactly the order this chapter follows.
Hydrogen has the simplest atomic structure of all the elements. A hydrogen atom is built from a nucleus carrying a charge of \(+1\) and a single extranuclear electron — that is, one electron outside the nucleus. There is nothing else: no inner shell, no core of electrons shielding the outer one.
This single electron is not lost easily. Hydrogen has only a little tendency to give it up, but it pairs readily with an electron from another atom to form a covalent bond — a bond made by sharing a pair of electrons. The simplest partner available to a hydrogen atom is another hydrogen atom, so in the free state the element exists not as separate H atoms but as \(\mathrm{H_2}\) molecules. Because of this two-atom form, free hydrogen is often called dihydrogen.
8.1.1 Occurrence
In the free state hydrogen exists as dihydrogen gas. It is the most abundant element in the universe; it makes up 70 % of the total mass of the universe. Hydrogen is also the principal element in the solar system, which is why the Sun and the stars are, in bulk, mostly hydrogen.
On the earth the picture is different. Here hydrogen is only the tenth most abundant element on a mass basis, but it is the third most abundant element on an atom basis.
Why Two Different Rankings
Those two rankings sound contradictory until you notice what each one counts. A “mass basis” ranking asks what fraction of the earth’s mass a given element contributes; an “atom basis” ranking asks what fraction of the total number of atoms it contributes. A hydrogen atom is the lightest atom there is, so a very large number of hydrogen atoms still adds up to a very small mass. Counting heads and weighing them give different answers.
8.1.2 Position of hydrogen in the periodic table
The position of hydrogen in the periodic table has been a subject of discussion for a long time, and the argument is worth following because it shows how a single electronic configuration can point in two directions at once.
The electronic configuration of hydrogen is \(1s^1\). This is similar to \(ns^1\), which is the outer electronic configuration of the alkali metals belonging to group 1. On that evidence alone, hydrogen sits comfortably at the head of group 1.
But \(1s^1\) also resembles the outer electronic configuration of the group 17 elements, which is \(ns^2np^5\). The resemblance is of a different kind: both hydrogen and a halogen atom are one electron short of the configuration of the nearest inert gas. By adding one electron to H we get the electronic configuration of the inert gas He, which is \(1s^2\); and by adding one electron to \(ns^2np^5\) we get \(ns^2np^6\), which is the outer electronic configuration of the remaining inert gases.
So hydrogen can be read as an element with one electron to lose, like an alkali metal, or as an element with one vacancy to fill, like a halogen. Experiment agrees with both readings: some chemical properties of hydrogen are similar to those of the alkali metals, while some resemble the halogens. You will meet the actual reactions that support each side in section 8.1.5.
The Uniqueness of the Hydrogen Ion
There is one property that neither group can match, and it is the strongest argument for treating hydrogen separately. The uniqueness of hydrogen is that \(\mathrm{H^{+}}\), formed by loss of the electron from a hydrogen atom, does not exist freely. It is always associated with other molecules. For example, in water it attaches itself at once to a water molecule:
\[\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\]
This is because \(\mathrm{H^{+}}\) is nothing else but a proton. Remove the one electron from a hydrogen atom and no electron cloud is left at all — only the bare nucleus. Every other cation, including \(\mathrm{Na^{+}}\), still carries a full set of inner electrons around it. A bare proton is so small and its charge so concentrated that it cannot survive on its own; it immediately attaches itself to any species that can offer it a lone pair of electrons.
Hydrogen is, therefore, placed separately above the group 1. It may be noted, here, that metastable metallic hydrogen was discovered at Harvard university, USA, in January 2017.
8.1.3 Isotopes of Hydrogen
If different atoms of the same element have different mass numbers they are called isotopes of each other. You met this idea in chapter 4; recall that the mass number \(A\) counts protons and neutrons together, while the atomic number \(Z\) counts protons only. Since \(Z\) fixes the identity of the element, isotopes differ only in the number of neutrons, \(N = A - Z\).
Hydrogen has three isotopes with mass numbers 1, 2 and 3. They all contain one proton and one electron but a different number of neutrons in the nucleus. Their structures are compared in Fig. 8.1, and their characteristics are collected in Table 8.1.
Diagram
Fig. 8.1 – Isotopes of hydrogen: three circles showing (a) protium ¹₁H, nucleus with 1 proton and 1 orbiting electron; (b) deuterium ²₁H, nucleus with 1 proton + 1 neutron; (c) tritium ³₁H, nucleus with 1 proton + 2 neutrons; label each with symbol and name.
Because all three isotopes have one electron each, their chemistry is essentially the same — chemical behaviour is decided by electrons, not by neutrons. What differs is mass, and therefore density, rate of diffusion, and nuclear stability.
Table 8.1 : Characteristics of isotopes of hydrogen
| Name of the Isotope | Symbol | Atomic number \(Z\) | Atomic mass number \(A\) | Neutron number \(N\) | Abundance | Stability |
|---|---|---|---|---|---|---|
| Hydrogen | H or \(^{1}_{1}\mathrm{H}\) or H-1 | 1 | 1 | 0 | 99.98 % | Stable |
| Deuterium | D or \(^{2}_{1}\mathrm{H}\) or H-2 | 1 | 2 | 1 | 0.015 % | Stable |
| Tritium | T or \(^{3}_{1}\mathrm{H}\) or H-3 | 1 | 3 | 2 | Trace | Radioactive |
Tritium, \(^{3}_{1}\mathrm{H}\), is a radioactive nuclide with a half life period of 12.4 years and emits low energy \(\beta\) particles. Half life period means the time in which half of a given quantity of the radioactive sample decays away; a value of 12.4 years tells you that tritium disappears steadily and therefore is never found on the earth in more than trace amounts.
8.1.4 Preparation of dihydrogen
Hydrogen can be prepared using many methods. The choice depends on the scale: a small quantity for a school laboratory is made in a completely different way from the thousands of tonnes needed by a fertiliser plant.
Laboratory Methods
i. Dihydrogen is prepared in the laboratory by the action of dilute hydrochloric acid on zinc granules.
\[\mathrm{Zn}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
Granules are used rather than powder because they give a steady, controllable stream of gas instead of a sudden rush.
ii. Dihydrogen can be prepared by the action of an aqueous solution of sodium hydroxide on zinc.
\[\mathrm{Zn}\text{(s)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\]
Notice that zinc releases dihydrogen with an acid and also with an alkali. Very few metals behave in both ways, and this double behaviour is characteristic of zinc.
Industrial Methods
i. By electrolysis of pure water: Pure water is a poor conductor of electricity. Therefore a dilute aqueous solution of acid or alkali is used to prepare dihydrogen by electrolysis. The added acid or alkali supplies mobile ions and so allows the current to pass; it is not itself used up in the overall change.
For example, electrolysis of a dilute aqueous solution of sulphuric acid yields two volumes of hydrogen at the cathode and one volume of oxygen at the anode, as shown in Fig. 8.2.
\[2\,\mathrm{H_2O}\text{(l)} \xrightarrow[\ \text{trace of acid or alkali}\ ]{\ \text{electrolysis}\ } 2\,\mathrm{H_2}\uparrow + \mathrm{O_2}\uparrow\]
Diagram
Fig. 8.2 – Electrolysis of acidified water: a cell with two inverted graduated tubes over the electrodes, gas collected at the cathode filling twice the volume of the gas at the anode; label cathode (H₂), anode (O₂), dilute H₂SO₄ solution and battery.
The 2 : 1 volume ratio is not an accident — it is the 2 : 1 mole ratio of the balanced equation showing itself directly, because equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
Pure dihydrogen (> 99.5 % purity) gas is obtained by electrolysis of a warm solution of barium hydroxide between nickel electrodes.
ii. From carbon or hydrocarbon: Three stages are involved in this industrial process of preparation of dihydrogen. The overall route is set out in FC 8.1.
Flowchart
FC 8.1 – Industrial preparation of dihydrogen from carbon or hydrocarbon: hydrocarbon (CH₄) or coke (C) + steam → Stage 1: reaction at 1270 K over Ni catalyst → water-gas (CO + H₂), also called syngas → Stage 2: water-gas shift reaction with steam at 673 K over iron chromate catalyst → CO₂ + H₂ → Stage 3: CO₂ removed by scrubbing with sodium arsenite solution → dihydrogen
Stage 1 : Reaction of steam on a hydrocarbon or coke (C) at 1270 K temperature in the presence of a nickel catalyst gives water-gas, which is a mixture of carbon monoxide and hydrogen.
\[\mathrm{CH_4}\text{(g)} + \mathrm{H_2O}\text{(g)} \xrightarrow[\ \mathrm{Ni}\ ]{\ 1270\ \text{K}\ } \underbrace{\mathrm{CO}\text{(g)} + 3\,\mathrm{H_2}\text{(g)}}_{\text{water-gas}}\]
As water-gas is used for the synthesis of \(\mathrm{CH_3OH}\) and a number of hydrocarbons, it is also called ‘syngas’. Production of syngas is also the first stage of gasification of coal.
\[\mathrm{C}\text{(s)} + \mathrm{H_2O}\text{(g)} \xrightarrow{\ 1270\ \text{K}\ } \mathrm{CO}\text{(g)} + \mathrm{H_2}\text{(g)}\]
Sawdust, scrapwood, etc. can also be used in place of carbon, which makes the process flexible about its raw material.
Stage 2 : Water-gas shift reaction: The carbon monoxide in the water-gas is transformed into carbon dioxide by reacting with steam in the presence of iron chromate as catalyst. This is called the water-gas shift reaction.
\[\mathrm{CO}\text{(g)} + \mathrm{H_2O}\text{(g)} \xrightarrow[\text{catalyst}]{\ 673\ \text{K},\ \text{iron chromate}\ } \mathrm{CO_2}\text{(g)} + \mathrm{H_2}\text{(g)}\]
The point of this stage is worth stating plainly: the unwanted carbon monoxide is not simply thrown away. It is made to take an oxygen atom from steam, and in doing so it releases one more molecule of dihydrogen. The yield of hydrogen therefore rises while the impurity is converted into a gas that is easy to remove.
Stage 3 : Carbon dioxide is removed by scrubbing with sodium arsenite solution. Scrubbing means washing the gas stream through a liquid that absorbs one particular component; what leaves the scrubber is dihydrogen.
Sources of Industrial Dihydrogen Today
Today major industrial production of dihydrogen is from petrochemicals (\(\sim\) 77 %), about 18 % from coal, about 4 % by electrolytic methods and about 1 % by other methods. These figures are shown in Fig. 8.3.
Diagram
Fig. 8.3 – Bar chart of sources of industrial dihydrogen: petrochemicals ~77 %, coal ~18 %, electrolytic methods ~4 %, other methods ~1 %; label the vertical axis “percentage of total production”.
8.1.5 Properties of dihydrogen
Physical Properties
Dihydrogen is a colourless, tasteless and odourless gas. It burns with a pale blue flame. It is a nonpolar, water insoluble gas, lighter than air.
Each of these properties follows from the structure of the molecule. The two atoms in \(\mathrm{H_2}\) are identical, so the shared electron pair is pulled equally by both nuclei and the molecule has no positive or negative end — it is nonpolar. Since water is a polar solvent, and like dissolves like, dihydrogen is almost insoluble in water. And because \(\mathrm{H_2}\) has the smallest molar mass of any gas, it is lighter than air.
Chemical Properties
Before looking at individual reactions, note one fact that shapes them all. The bond dissociation energy of the H–H bond is high, 436 kJ mol\(^{-1}\) — this is the energy needed to break the bond in one mole of \(\mathrm{H_2}\) molecules into free atoms. Therefore reactions of dihydrogen take place at high temperature and/or in the presence of a catalyst. Whenever you see “\(\Delta\)” or a catalyst written over an arrow in the reactions below, this is the reason.
i. Reaction with metals : Dihydrogen combines with all the reactive metals such as alkali metals, calcium, strontium and barium at high temperature, to form metal hydrides. For example :
\[2\,\mathrm{Na}\text{(s)} + \mathrm{H_2}\text{(g)} \xrightarrow{\ \Delta\ } 2\,\mathrm{NaH}\text{(s)}\]
In this respect dihydrogen is similar to the halogens, which also react with metals and form metal halides. In sodium hydride the hydrogen is present as the hydride ion, \(\mathrm{H^{-}}\) — that is, hydrogen has gained an electron here, exactly as a halogen atom does.
ii. Reaction with dioxygen: Dihydrogen reacts with dioxygen in the presence of a catalyst or by heating to form water. This reaction is highly exothermic.
\[2\,\mathrm{H_2}\text{(g)} + \mathrm{O_2}\text{(g)} \xrightarrow[\text{or heating}]{\ \text{catalyst}\ } 2\,\mathrm{H_2O}\text{(l)} \qquad \Delta H = -235\ \text{kJ mol}^{-1}\]
The large negative value is what makes dihydrogen a fuel: a great deal of energy is released, and the only product is water.
iii. Reaction with halogens: Dihydrogen inflames with fluorine even at \(-250\ ^\circ\mathrm{C}\) in the dark, whereas it requires a catalyst to react with iodine. The vigour of the reaction of dihydrogen decreases with increasing atomic number of the halogen.
\[\mathrm{H_2}\text{(g)} + \mathrm{X_2}\text{(g)} \rightarrow 2\,\mathrm{HX}\text{(g)}\]
In this respect dihydrogen resembles the alkali metals, which also react with halogens to form halides. Here is the double character of hydrogen showing itself again: with sodium it behaves like a halogen, and with chlorine it behaves like an alkali metal.
iv. Reducing nature of dihydrogen :
a. Dihydrogen reduces the oxides and ions of metals that are less reactive than iron, to the corresponding metals, at moderate temperature. For example :
\[\mathrm{CuO}\text{(s)} + \mathrm{H_2}\text{(g)} \rightarrow \mathrm{Cu}\text{(s)} + \mathrm{H_2O}\text{(l)}\]
\[\mathrm{Fe_3O_4}\text{(s)} + 4\,\mathrm{H_2}\text{(g)} \rightarrow 3\,\mathrm{Fe}\text{(s)} + 4\,\mathrm{H_2O}\text{(l)}\]
\[\mathrm{Pd^{2+}}\text{(aq)} + \mathrm{H_2}\text{(g)} \rightarrow \mathrm{Pd}\text{(s)} + 2\,\mathrm{H^{+}}\text{(aq)}\]
In every one of these changes dihydrogen hands over electrons — it takes the oxygen away from the metal oxide, or takes the positive charge away from the metal ion. That is what a reducing agent does. The restriction “less reactive than iron” matters: metals more reactive than iron hold their oxygen too strongly for dihydrogen to remove it.
b. Hydrogenation of unsaturated organic compound
The hydrogenation of unsaturated organic compounds such as oil using a nickel catalyst gives saturated organic compounds such as solid fat (vanaspati ghee). The change taking place at the double bond is drawn in Fig. 8.4.
Diagram
Fig. 8.4 – Hydrogenation of an unsaturated oil: structure showing C=C (labelled “vegetable oil”) + H₂, arrow over which is written “Ni, Δ”, giving the product CH–CH single bond (labelled “saturated fat”); show the two new C–H bonds formed.
The double bond of the oil molecule opens out and one hydrogen atom adds to each of the two carbon atoms. The liquid oil, whose molecules could not pack closely because of the rigid double bonds, becomes a semi-solid fat once those bonds are saturated.
8.1.6 Uses of dihydrogen
i. The largest use of dihydrogen is in the production of ammonia.
ii. Dihydrogen is used in the formation of vanaspati ghee by catalytic hydrogenation of oils.
iii. Liquid dihydrogen is used as a rocket fuel.
iv. Dihydrogen is used in the preparation of important organic compounds like methanol in bulk quantity.
\[2\,\mathrm{H_2}\text{(g)} + \mathrm{CO}\text{(g)} \xrightarrow{\ \text{cobalt catalyst}\ } \mathrm{CH_3OH}\text{(l)}\]
v. Dihydrogen is used for the preparation of hydrogen chloride (HCl) and metal hydrides.
Read this list together with the reactions of section 8.1.5 and the pattern becomes clear. Ammonia synthesis and methanol synthesis both use dihydrogen as a builder of new bonds; hydrogenation and hydride formation both use its ability to add hydrogen to another substance; and rocket fuel uses the large negative \(\Delta H\) of its reaction with dioxygen.
8.2 Alkali metals and alkaline earth metals
8.2.1 Introduction
The elements of the groups 1 and 2 are placed on the left in the periodic table. Here the last electron enters into the ‘ns’ subshell. Thus they belong to the s-block of the periodic table.
Group 1 of the periodic table consists of the elements: hydrogen, lithium, sodium, potassium, rubidium, caesium and francium. These elements except hydrogen are collectively called alkali metals. We have already looked at hydrogen in the section 8.1, and you saw there why it is kept apart from the family.
Two elements of Group 1, namely, sodium and potassium are the sixth and seventh most abundant elements in the earth’s crust. Francium, in contrast, does not occur appreciably in nature because it is radioactive and has a short half-life period — any francium formed decays away long before it can accumulate.
Group 2 of the periodic table consists of the elements : beryllium, magnesium, calcium, strontium, barium and radium. These elements are collectively called alkaline earth metals because they occur as minerals in rocks. The elements magnesium and calcium are found abundantly in the earth’s crust, but radium is not easy to find. Radium is one of the first two radioactive elements discovered by Madame Curie.
8.2.2 Electronic configuration of elements of group 1 and group 2
The general outer electronic configuration of the group 1 elements is \(ns^1\) and that of the group 2 elements is \(ns^2\). The loosely held s-electrons in the valence shell of these elements can be easily removed to form metal ions. Hence these elements are never found in the free state in nature.
That last sentence is the key to the whole chapter. An outer s-electron in these atoms is far from the nucleus and is screened from it by all the filled inner shells, so the pull holding it is weak. Air and water take those electrons away almost immediately, which is why every one of these metals is mined as a compound and never as the metal itself.
Tables 8.2 and 8.3 show the electronic configurations of the elements of group 1 and group 2 elements, respectively.
Table 8.2 : Electronic configuration of group 1 elements
| Name | Symbol | Atomic number | Condensed electronic configuration | Electronic Configuration |
|---|---|---|---|---|
| Hydrogen | H | 1 | \(1s^1\) | \(1s^1\) |
| Lithium | Li | 3 | \([\mathrm{He}]\,2s^1\) | \(1s^2\,2s^1\) |
| Sodium | Na | 11 | \([\mathrm{Ne}]\,3s^1\) | \(1s^2\,2s^2\,2p^6\,3s^1\) |
| Potassium | K | 19 | \([\mathrm{Ar}]\,4s^1\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^1\) |
| Rubidium | Rb | 37 | \([\mathrm{Kr}]\,5s^1\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^1\) |
| Caesium | Cs | 55 | \([\mathrm{Xe}]\,6s^1\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^2\,4d^{10}\,5p^6\,6s^1\) |
| Francium | Fr | 87 | \([\mathrm{Rn}]\,7s^1\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^2\,4d^{10}\,5p^6\,6s^2\,4f^{14}\,5d^{10}\,6p^6\,7s^1\) |
Table 8.3 : Electronic configuration of group 2 elements
| Name | Symbol | Atomic number | Condensed electronic configuration | Electronic Configuration |
|---|---|---|---|---|
| Beryllium | Be | 4 | \([\mathrm{He}]\,2s^2\) | \(1s^2\,2s^2\) |
| Magnesium | Mg | 12 | \([\mathrm{Ne}]\,3s^2\) | \(1s^2\,2s^2\,2p^6\,3s^2\) |
| Calcium | Ca | 20 | \([\mathrm{Ar}]\,4s^2\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\) |
| Strontium | Sr | 38 | \([\mathrm{Kr}]\,5s^2\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^2\) |
| Barium | Ba | 56 | \([\mathrm{Xe}]\,6s^2\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^2\,4d^{10}\,5p^6\,6s^2\) |
| Radium | Ra | 88 | \([\mathrm{Rn}]\,7s^2\) | \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^2\,3d^{10}\,4p^6\,5s^2\,4d^{10}\,5p^6\,6s^2\,4f^{14}\,5d^{10}\,6p^6\,7s^2\) |
Reading the Condensed Notation
The condensed form, such as \([\mathrm{Ne}]\,3s^1\) for sodium, is only a shorthand. The symbol of the previous inert gas in square brackets stands for all the electrons of that inert gas — for Na, the eleven electrons are written as the ten of neon plus one in \(3s\). Writing configurations this way makes the family pattern obvious at a glance: every group 1 element ends in \(ns^1\) and every group 2 element in \(ns^2\), no matter how long the full configuration becomes.
8.2.3 Trends in atomic and physical properties of elements of group 1 and group 2
All the alkali metals are silvery white and soft. Due to their large atomic size these elements have low density. They are the most electropositive elements — that is, they give up electrons more readily than any other elements.
The alkaline earth metals are also, in general, silvery white lustrous and soft, but harder than the alkali metals. They are also strongly electropositive in nature, but comparatively less electropositive than the alkali metals. Two valence electrons and a higher nuclear charge make each atom hold on a little more tightly, and this shows up as greater hardness, higher melting point and higher density.
Some atomic and physical properties of the alkali metals and the alkaline earth metals are listed in the tables 8.4 and 8.5 respectively.
Table 8.4 : Physical properties of group 1 elements (except hydrogen)
| Symbol | Atomic radius (pm) | Ionic radius (pm) | Density (g cm\(^{-3}\)) | Ionization enthalpy (kJ mol\(^{-1}\)) | Electronegativity | Melting point (K) | Abundance in the lithosphere | Standard reduction potential \(E^\circ\) (V) |
|---|---|---|---|---|---|---|---|---|
| Li | 152 | 76 | 0.54 | 520 | 1.0 | 454 | 18 ppm | –3.04 |
| Na | 186 | 102 | 0.97 | 496 | 0.9 | 371 | 2.27 % | –2.714 |
| K | 227 | 138 | 0.86 | 419 | 0.8 | 336 | 1.84 % | –2.925 |
| Rb | 248 | 152 | 1.53 | 403 | 0.8 | 312 | 78.12 ppm | –2.930 |
| Cs | 265 | 167 | 1.90 | 376 | 0.7 | 302 | 2.6 ppm | –2.927 |
| Fr | – | (180) | – | $$375 | – | – | \(\sim 10^{-18}\) ppm | – |
Table 8.5 : Physical properties of group 2 elements
| Symbol | Atomic radius (pm) | Ionic radius (pm) | Density (g cm\(^{-3}\)) | Ionization enthalpy 1st (kJ mol\(^{-1}\)) | Ionization enthalpy 2nd (kJ mol\(^{-1}\)) | Electronegativity | Melting point (K) | Abundance in the lithosphere | Standard reduction potential \(E^\circ\) (V) |
|---|---|---|---|---|---|---|---|---|---|
| Be | 111 | 31 | 1.84 | 899 | 1757 | 1.5 | 1560 | 2 ppm | –1.97 |
| Mg | 160 | 72 | 1.74 | 737 | 1450 | 1.2 | 924 | 2.76 % | –2.36 |
| Ca | 197 | 100 | 1.55 | 590 | 1145 | 1.0 | 1124 | 4.6 % | –2.84 |
| Sr | 215 | 118 | 2.63 | 549 | 1064 | 1.0 | 1062 | 384 ppm | –2.89 |
| Ba | 222 | 135 | 3.59 | 503 | 965 | 0.9 | 1002 | 390 ppm | –2.92 |
| Ra | – | 148 | (5.5) | 509 | 979 | – | 973 | \(10^{-6}\) ppm | –2.92 |
Why the Ions Are Colourless
Unipositive ions of all the elements of group 1 have inert gas configuration. Thus they have no unpaired electron and their compounds are diamagnetic and colourless. The divalent ions of group 2 elements also have inert gas configuration with no unpaired electron, and therefore their compounds are also diamagnetic and colourless.
Two words there need unpacking. Diamagnetic means the substance is weakly repelled by a magnetic field, which is what happens when every electron in it is paired; a species with unpaired electrons would instead be paramagnetic and be drawn into the field. Colour, likewise, usually requires electrons that can be lifted into a nearby empty level by visible light. An ion with a completed inert gas shell offers neither unpaired electrons nor a low-lying vacancy, so it is both diamagnetic and colourless.
Lithium and beryllium differ from the rest of the elements of the groups 1 and 2, respectively, because of their extremely small size and comparatively high electronegativity. Keep this in mind — it returns in section 8.2.5 as the reason behind the diagonal relationship.
Regularity of the Periodic Trends
The physical properties of group 1 and group 2 elements show reasonable regularity in the periodic trends. Thus the atomic and ionic radii and densities increase down both the groups. Ionization enthalpies and electronegativities decrease down both the groups. The elements of both these groups, in general, have high negative values of standard reduction potentials.
Read those trends together and one cause runs through all of them. Going down a group, a new shell is added at each step, so the valence electron sits further from the nucleus and is screened by more inner electrons. The atom grows; the outer electron is held more weakly, so less energy is needed to remove it and the ionization enthalpy falls; and the atom pulls less strongly on a shared electron pair, so the electronegativity falls too.
Two small departures from the trend are worth noticing in Table 8.4, because Boards and entrance papers like to ask about them: the density of potassium is lower than that of sodium, and the standard reduction potential of lithium is the most negative of the group even though lithium has the highest ionization enthalpy in it. The second of these is explained in section 8.2.4 under the reducing nature of these metals.
The variation of the two most important of these properties is plotted in Fig. 8.5.
Diagram
Fig. 8.5 – Two line graphs sharing an x-axis of the group 1 elements Li → Cs: upper graph atomic radius (pm) rising steadily, lower graph first ionization enthalpy (kJ mol⁻¹) falling steadily; label each point with the element symbol.
8.2.4 Chemical properties of elements of group 1 and group 2
The alkali metals and alkaline earth metals are very reactive in nature. As a result of this they are always found in the combined state. Their reactivity is due to their low ionization enthalpy values in general. The reactivity of these metals increases with increasing atomic radius and the corresponding lowering of ionization enthalpy down the groups 1 and 2 can be noticed.
i. Reaction with oxygen/air
Group 1 - All the elements of group 1 rapidly lose their lustre in air due to the formation of a layer of oxide, or peroxide, and in some cases superoxide, by reaction with oxygen in air.
\[4\,\mathrm{Li} + \mathrm{O_2} \rightarrow 2\,\mathrm{Li_2O} \qquad \text{(Lithium oxide)}\]
\[2\,\mathrm{Na} + \mathrm{O_2} \rightarrow \mathrm{Na_2O_2} \qquad \text{(Sodium peroxide)}\]
\[\mathrm{K} + \mathrm{O_2} \rightarrow \mathrm{KO_2} \qquad \text{(Potassium superoxide)}\]
Notice the pattern down the group: the small lithium ion gives the simple oxide with \(\mathrm{O^{2-}}\), sodium gives the peroxide with \(\mathrm{O_2^{2-}}\), and the still larger potassium ion gives the superoxide with \(\mathrm{O_2^{-}}\). A large, weakly polarising cation can support a large anion without pulling it apart, and that is why the bigger metals form the bigger oxygen species.
The oxides of group 1 metals are strongly basic in nature. They dissolve in water forming aqueous solutions of strong alkali. For example
\[\mathrm{Li_2O}\text{(s)} + \mathrm{H_2O}\text{(l)} \rightarrow 2\,\mathrm{LiOH}\text{(aq)}\]
Group 2 : These metals are protected from air oxidation by an oxide film formed on their surface. The film is thin, unbroken and sticks to the metal, so it seals the surface instead of flaking off — which is exactly why a magnesium ribbon can be stored in air while sodium cannot.
All the elements of group 2 burn when ignited in air forming MO type oxides. The product is a mixture of oxide and nitride, because these metals are reactive enough to attack even the nitrogen of the air.
\[2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\]
\[3\,\mathrm{Mg} + \mathrm{N_2} \rightarrow \mathrm{Mg_3N_2}\]
Further heating of the oxide in air results in the formation of peroxide.
ii. Reaction with water
Group 1 : Lithium, sodium and potassium all float on water due to hydrogen bubbles released on reaction with water. Lithium reacts slowly but sodium and potassium react vigorously with water. Due to the highly exothermic reaction, sodium and potassium catch fire when put in water.
\[2\,\mathrm{Na} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + \mathrm{H_2}\uparrow\]
There are two separate reasons for floating here, and it is worth separating them. These metals are less dense than water to begin with — look up the density of Li, Na and K in Table 8.4 — and the stream of hydrogen bubbles clinging to the metal lowers the effective density further.
iii. Reaction with Hydrogen
Group 1 : Alkali metals react with hydrogen at high temperature to form the corresponding metal hydrides.
\[2\,\mathrm{M} + \mathrm{H_2} \xrightarrow{\ 673\ \text{K}\ } 2\,\mathrm{M^{+}H^{-}}\]
Group 2 : All the metals of group 2, except beryllium, when heated with hydrogen form \(\mathrm{MH_2}\) type hydrides.
\[\mathrm{M} + \mathrm{H_2} \xrightarrow{\ \Delta\ } \mathrm{MH_2}\]
The formulae written with charges, \(\mathrm{M^{+}H^{-}}\), make an important point: in these hydrides hydrogen is present as the hydride ion, carrying a negative charge. This is the ionic hydride formation you met from hydrogen’s side in section 8.1.5, seen now from the metal’s side.
iv. Reaction with Halogens
Group 1 : All the alkali metals react vigorously with halogens to produce their ionic halide salts.
\[2\,\mathrm{M} + \mathrm{X_2} \rightarrow 2\,\mathrm{M^{+}X^{-}}\]
Group 2: All the alkaline earth metals combine with halogens at high temperature to form halides.
\[\mathrm{M} + \mathrm{X_2} \rightarrow \mathrm{MX_2}\]
v. Reducing nature
Group 1 : The reducing power of an element is measured in terms of the standard electrode potential (\(E^\circ\)) corresponding to the transformation
\[\mathrm{M^{+}}\text{(aq)} + e^{-} \rightarrow \mathrm{M}\text{(s)}\]
All the alkali metals have high negative values of \(E^\circ\), indicative of their strong reducing nature. Lithium is the most powerful and sodium is the least powerful in the group (see Table 8.4).
A negative \(E^\circ\) says that the reverse change — the metal giving up its electron — is the favoured direction, and giving up electrons to something else is exactly what a reducing agent does. The more negative the value, the stronger the reducing agent.
Group 2 : All the alkaline earth metals have high negative values of standard reduction potential (\(E^\circ\)) (see Table 8.5), and are strong reducing agents. However their reducing power is less than those of the alkali metals.
vi. Solution in liquid ammonia
Group 1 : The alkali metals are soluble in liquid ammonia giving deep blue coloured solutions which show electrical conductivity.
\[\mathrm{M} + (x + y)\,\mathrm{NH_3} \rightarrow [\mathrm{M(NH_3)}_x]^{+} + [e(\mathrm{NH_3})_y]^{-}\]
What this equation describes is unusual and worth pausing over. The metal atom hands its valence electron over to the solvent itself: the cation is surrounded by ammonia molecules, and so is the released electron. An electron held in a cage of ammonia molecules in this way is called an ammoniated electron.
The ammoniated electron is responsible for the deep blue colour of these solutions. These solutions are paramagnetic — the freed electrons are unpaired — and on standing they liberate hydrogen slowly, resulting in the formation of the metal amide. The blue colour changes to bronze and the solution becomes diamagnetic.
\[\mathrm{M^{+}}\text{(am)} + e^{-}\text{(am)} + \mathrm{NH_3}\text{(l)} \rightarrow \mathrm{MNH_2}\text{(am)} + \tfrac{1}{2}\,\mathrm{H_2}\text{(g)}\]
(Here (am) denotes solution in ammonia.)
Group 2 : Similar to the alkali metals, the alkaline earth metals are also soluble in liquid ammonia, which gives deep blue black coloured solutions.
\[\mathrm{M} + (x + 2y)\,\mathrm{NH_3} \rightarrow [\mathrm{M(NH_3)}_x]^{2+} + 2\,[e(\mathrm{NH_3})_y]^{-}\]
Note the coefficient 2 in front of the ammoniated electron: a group 2 atom releases two electrons, not one, which is why twice as many ammoniated electrons appear.
8.2.5 Diagonal Relationship
It is expected that elements belonging to the same group exhibit similarity and gradation in their properties. The first alkali metal lithium and the first alkaline earth metal beryllium do not fulfil this expectation.
Thus, lithium shows many differences when compared with the remaining alkali metals and resembles magnesium, the second alkaline earth metal. Likewise beryllium shows many differences with the remaining alkaline earth metals and shows similarity with aluminium, the second element of the next main group (group 13).
The relative placement of these elements with similar properties in the periodic table appears to be across a diagonal (see Table 8.6) and is called diagonal relationship.
Table 8.6 : Diagonal relationship
| Period Main Group | 1 | 2 | 13 |
|---|---|---|---|
| 2 | Li | Be | B |
| 3 | Na | Mg | Al |
Diagram
Fig. 8.6 – The diagonal relationship: a 2 × 3 grid of the elements of Table 8.6 with arrows drawn diagonally from Li to Mg and from Be to Al, each arrow labelled “similar properties”.
The table 8.7 shows some properties of lithium and magnesium which elucidate their diagonal relationship.
Table 8.7 : Resemblance between Li and Mg
| Element Criterion | Products of reaction with air, \(\mathrm{M} \xrightarrow{\text{air}} ?\) | Products of thermal decomposition of carbonate, \(\mathrm{M_2CO_3/MCO_3} \xrightarrow{\Delta} ?\) | Property of chloride | Formula of crystalline chloride |
|---|---|---|---|---|
| Li | \(\mathrm{Li_2O} + \mathrm{Li_3N}\) | \(\mathrm{Li_2O} + \mathrm{CO_2}\) | deliquescent | \(\mathrm{LiCl \cdot 2H_2O}\) |
| Mg | \(\mathrm{MgO} + \mathrm{Mg_3N_2}\) | \(\mathrm{MgO} + \mathrm{CO_2}\) | deliquescent | \(\mathrm{MgCl_2 \cdot 8H_2O}\) |
| Group 1 (except Li) | \(\mathrm{M_2O}/\mathrm{M_2O_2}/\mathrm{MO_2}\) | No reaction | Not deliquescent | MCl |
| Group 2 | \(\mathrm{MO} + \mathrm{M_3N_2}\) | \(\mathrm{MO} + \mathrm{CO_2}\) | deliquescent | \(\mathrm{MCl_2 \cdot xH_2O}\) |
Read the first two rows against the third. Lithium behaves like a group 2 metal and not like its own family: it forms a nitride in air, its carbonate decomposes on heating, and its chloride is deliquescent (it absorbs so much moisture from the air that it dissolves in the water it has taken up) and crystallises with water of crystallisation. The other alkali metals do none of these things.
In table 8.8 some properties of Be and Al are shown which indicate the diagonal relationship.
Table 8.8 : Resemblance between Be and Al
| Element Criterion | Nature of bonding in the chloride | Whether Lewis acid | Solubility in organic solvent | Properties of oxide — Acidic/Basic/Amphoteric |
|---|---|---|---|---|
| Be | Covalent, chain structure with Cl bridges | \(\mathrm{BeCl_2}\) is a strong Lewis acid | Soluble | Amphoteric: \(\mathrm{BeO} + 2\,\mathrm{HCl} \rightarrow \mathrm{BeCl_2} + \mathrm{H_2O}\); \(\mathrm{BeO} + 2\,\mathrm{NaOH} \rightarrow \mathrm{Na_2BeO_2} + \mathrm{H_2O}\) |
| Al | Covalent, dimer with Cl bridges | \(\mathrm{AlCl_3}\) is a strong Lewis acid | Soluble | Amphoteric: \(\mathrm{Al_2O_3} + 6\,\mathrm{HCl} \rightarrow 2\,\mathrm{AlCl_3} + 3\,\mathrm{H_2O}\); \(\mathrm{Al_2O_3} + 2\,\mathrm{NaOH} \rightarrow 2\,\mathrm{NaAlO_2} + \mathrm{H_2O}\) |
| Group 2 | Ionic | Not a Lewis acid | Insoluble | Basic: \(\mathrm{MO} + 2\,\mathrm{HCl} \rightarrow \mathrm{MCl_2} + \mathrm{H_2O}\); \(\mathrm{MO} + \mathrm{NaOH} \rightarrow\) No reaction |
| Group 13 | Covalent | Lewis acid | Soluble | Amphoteric |
The bridged structures named in the first column of Table 8.8 are drawn in Fig. 8.7.
Diagram
Fig. 8.7 – (a) Chain structure of solid BeCl₂: repeating …Cl–Be–Cl–Be–Cl… units with each pair of Be atoms joined through two bridging Cl atoms; (b) dimeric Al₂Cl₆: two Al atoms joined by two bridging Cl atoms, each Al carrying two terminal Cl atoms; label the bridging and terminal chlorine atoms in both.
An oxide described as amphoteric reacts with an acid and with an alkali, as the two equations in each of the first two rows show. A basic oxide, such as MgO or CaO, reacts only with the acid. This single difference is one of the sharpest tests of the Be–Al likeness.
Why the Diagonal Likeness Exists
The diagonal relationship between the elemental pairs belonging to different groups and periods is due to the similarity in some of their atomic properties. Thus the atomic and ionic radii of Li and Mg are very similar (see the table 8.4). In the case of Be and Al the charge to radius ratio of their ions is very similar (Be : \(\dfrac{2}{31}\) and Al : \(\dfrac{3}{53.55}\)).
The reason this happens at all is that two opposite trends cancel along a diagonal. Moving one step to the right across a period makes atoms smaller and more electronegative; moving one step down a group makes them larger and less electronegative. Take both steps together — which is what a diagonal move is — and the two changes largely cancel, leaving the second element with much the same size, charge density and polarising power as the first.
8.2.6 Uses of elements of group 1 and group 2
Group 1
i. Lithium metal is used in long-life batteries used in digital watches, calculators and computers.
ii. Liquid sodium has been used for heat transfer in nuclear power stations.
iii. Potassium chloride is used as a fertilizer.
iv. Potassium is used in manufacturing potassium superoxide (\(\mathrm{KO_2}\)) for oxygen generation. It is a good absorbent of carbon dioxide.
v. Caesium is used in photoelectric cells.
Group 2
i. Beryllium is used as a moderator in nuclear reactors.
ii. An alloy of magnesium and aluminium is widely used as a structural material and in aircraft.
iii. Calcium ions are an important ingredient in the biological system, essential for healthy growth of bones and teeth.
iv. Barium sulphate is used in medicine as a barium meal for intestinal X-ray.
v. Radium is used in radiotherapy for cancer treatment.
Several of these uses are simply a property of the element being put to work, and it is worth naming the link each time. Lithium’s low density and strongly negative \(E^\circ\) make it ideal for a light, high-voltage cell; caesium’s very low ionization enthalpy — the lowest in Table 8.4 — means visible light is energetic enough to knock an electron out of it, which is what a photoelectric cell needs; and barium sulphate is safe as a “barium meal” only because it is so insoluble that the toxic \(\mathrm{Ba^{2+}}\) ion is never released, while its heavy atoms block X-rays and outline the gut clearly.
8.2.7 Biological importance of elements of group 1 and group 2
Group 1
i. Sodium ion is present as the largest supply in all extracellular fluids. These fluids provide the medium for transporting nutrients to the cells.
ii. The concentration of sodium ion in extracellular fluids regulates the flow of water across the membrane.
iii. Sodium ions participate in the transmission of nerve signals.
iv. Potassium ions are the most abundant ions within cells. These are required for maximum efficiency in the synthesis of proteins and also in the oxidation of glucose.
Group 2
i. \(\mathrm{Mg^{2+}}\) ions are an important part of chlorophyll in green plants.
ii. \(\mathrm{Mg^{2+}}\) ions play an important role in the breakage of glucose and fat molecules, in the synthesis of proteins with enzymes, and in the regulation of cholesterol level.
iii. \(\mathrm{Ca^{2+}}\) ions are important for bones and teeth in the form of apatite, \([\mathrm{Ca_3(PO_4)_2}]\).
iv. \(\mathrm{Ca^{2+}}\) ions play an important role in blood clotting.
v. \(\mathrm{Ca^{2+}}\) ions are required for contraction and stretching of muscles.
vi. \(\mathrm{Ca^{2+}}\) ions are also required to maintain the regular beating of the heart.
8.3 Some important compounds of elements of s-block
In this section we consider five important compounds of s-block elements with reference to their preparation, properties and uses. Each of them is a substance you can meet outside a laboratory — in a washing tub, a soap factory, a cement kiln, a first-aid box or an organic chemistry bench — and in each case the properties explain the use.
8.3.1 Sodium Carbonate (washing soda) \(\mathrm{Na_2CO_3 \cdot 10H_2O}\)
Preparation
Sodium carbonate is commercially prepared by the Solvay process. The raw materials are cheap and plentiful — common salt, limestone and ammonia — which is what makes the process economical on an industrial scale. The route is summarised in FC 8.2.
Flowchart
FC 8.2 – Solvay process for sodium carbonate: brine (concentrated NaCl solution) saturated with NH₃ → CO₂ passed in → (NH₄)₂CO₃ → NH₄HCO₃ → double decomposition with NaCl → NaHCO₃ precipitates (low solubility) + NH₄Cl in solution → Stage 2: NaHCO₃ heated (Δ) → Na₂CO₃ + H₂O + CO₂ → CO₂ returned to the first stage; branch from NH₄Cl solution → treated with slaked lime Ca(OH)₂ → NH₃ recovered and returned to the brine + CaCl₂ as by-product
In the first stage of the process \(\mathrm{CO_2}\) is passed into a concentrated solution of NaCl which is saturated with \(\mathrm{NH_3}\). Crystals of sodium bicarbonate separate as a result of the following reactions.
Reactions in the first stage:
\[2\,\mathrm{NH_3}\text{(aq)} + \mathrm{H_2O} + \mathrm{CO_2}\text{(g)} \rightarrow (\mathrm{NH_4})_2\mathrm{CO_3}\text{(aq)}\]
\[(\mathrm{NH_4})_2\mathrm{CO_3}\text{(aq)} + \mathrm{H_2O} + \mathrm{CO_2}\text{(aq)} \rightarrow 2\,\mathrm{NH_4HCO_3}\text{(aq)}\]
\[\mathrm{NH_4HCO_3}\text{(aq)} + \mathrm{NaCl}\text{(aq)} \rightarrow \mathrm{NH_4Cl}\text{(aq)} + \mathrm{NaHCO_3}\text{(s)}\]
Sodium bicarbonate has low solubility, and therefore its crystals precipitate out, which are formed as a result of the double decomposition reaction between ammonium bicarbonate and sodium chloride.
That low solubility is the pivot of the whole process. In a double decomposition reaction the two salts simply exchange partners, and the exchange is driven forward only because one of the products — sodium bicarbonate — leaves the solution as a solid. All the other species stay dissolved and can be dealt with separately.
In the second stage the separated crystals of sodium bicarbonate are heated to obtain sodium carbonate.
\[2\,\mathrm{NaHCO_3}\text{(s)} \xrightarrow{\ \Delta\ } \mathrm{Na_2CO_3}\text{(s)} + \mathrm{H_2O}\text{(g)} + \mathrm{CO_2}\text{(g)}\]
Recovery of Ammonia
In this process the recovery of ammonia is done by treating the solution of \(\mathrm{NH_4Cl}\) obtained with slaked lime, \(\mathrm{Ca(OH)_2}\). The byproduct of this reaction is calcium chloride.
\[2\,\mathrm{NH_4Cl}\text{(aq)} + \mathrm{Ca(OH)_2}\text{(s)} \rightarrow 2\,\mathrm{NH_3}\text{(g)} + \mathrm{CaCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)}\]
The recovered ammonia goes straight back into the brine at the start of the process, and the carbon dioxide released in the second stage is returned as well. Ammonia is the one costly ingredient, so a process that used it up would never have been economical.
Why the Process Does Not Work for Potassium
Potassium carbonate cannot be prepared by the Solvay process because potassium hydrogen carbonate is highly water soluble and cannot be precipitated by reaction with potassium chloride. Since the entire process depends on one product leaving the solution as a solid, a soluble bicarbonate stops the sequence before it starts.
Properties
Sodium carbonate (washing soda) is a white crystalline solid having the formula \(\mathrm{Na_2CO_3 \cdot 10H_2O}\). It is highly soluble in water. On heating, the decahydrate loses water molecules to form the monohydrate. On heating above 373 K temperature the monohydrate further loses water and changes into a white anhydrous powder called soda-ash.
\[\mathrm{Na_2CO_3 \cdot 10H_2O}\text{(s)} \xrightarrow{\ 373\ \text{K}\ } \mathrm{Na_2CO_3 \cdot H_2O}\text{(s)} + 9\,\mathrm{H_2O}\text{(g)}\]
\[\mathrm{Na_2CO_3 \cdot H_2O}\text{(s)} \xrightarrow{\ >\,373\ \text{K}\ } \mathrm{Na_2CO_3}\text{(s)} + \mathrm{H_2O}\text{(g)}\]
The ten water molecules written after the dot are water of crystallisation — water built into the crystal lattice in a fixed proportion, not water that the solid has merely absorbed. This is why the loss happens in definite steps, nine molecules first and the tenth afterwards, at a higher temperature.
The aqueous solution of sodium carbonate is alkaline because of its hydrolysis by the following reaction:
\[\mathrm{Na_2CO_3} + \mathrm{H_2O} \rightarrow \mathrm{NaHCO_3} + \mathrm{NaOH}\]
Hydrolysis here means a reaction with water in which the salt of a strong base and a weak acid gives back some free alkali. The carbonate ion takes a hydrogen from water, leaving hydroxide ions behind — and free \(\mathrm{OH^{-}}\) in solution is exactly what makes a solution alkaline.
Uses
i. The alkaline properties of sodium carbonate are responsible for the emulsifying effect on grease and dirt. It is used as a cleaning material.
ii. It is used to make hard water soft (as a water softener), as it precipitates out the soluble calcium and magnesium salts in hard water as carbonates. For example :
\[\mathrm{Ca(HCO_3)_2}\text{(aq)} + \mathrm{Na_2CO_3}\text{(aq)} \rightarrow \mathrm{CaCO_3}\text{(s)} + 2\,\mathrm{NaHCO_3}\text{(aq)}\]
iii. It is used for the commercial production of soap and caustic soda.
iv. It is an important laboratory reagent.
8.3.2 Sodium hydroxide (caustic soda) NaOH
Preparation
Sodium hydroxide is commercially obtained by the electrolysis of a saturated aqueous solution of sodium chloride. Brine solution is subjected to electrolysis in a Castner-Kellner cell. Mercury is used as the cathode and a carbon rod as the anode. Metallic sodium liberated at the cathode forms sodium amalgam. Chlorine gas is evolved at the anode. The arrangement is shown in Fig. 8.8.
Diagram
Fig. 8.8 – Castner-Kellner cell: a trough of brine with a flowing mercury layer at the base acting as cathode, carbon rod anodes dipping into the brine from above, chlorine gas escaping at the anode, sodium amalgam flowing out of the cell into a separate chamber where it meets water to give NaOH and H₂; label brine inlet, mercury cathode, carbon anode, Cl₂ outlet, amalgam flow, water inlet, NaOH outlet.
Cathode reaction:
\[\mathrm{Na^{+}} + e^{-} \xrightarrow{\ \mathrm{Hg}\ } \text{Na-amalgam}\]
Anode reaction:
\[\mathrm{Cl^{-}} \rightarrow \tfrac{1}{2}\,\mathrm{Cl_2} + e^{-}\]
The mercury cathode is not a detail to skip over. Sodium metal set free in contact with water would react with it at once, so the cell gives the sodium somewhere else to go: it dissolves in the mercury as an amalgam and is carried out of the cell before it can meet the brine.
Sodium hydroxide is obtained by treating the sodium amalgam with water, when hydrogen gas is liberated.
\[2\,\text{Na-Hg} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + 2\,\mathrm{Hg} + \mathrm{H_2}\]
The mercury released here is returned to the cell and used again, so the same mercury circulates continuously.
Flowchart
FC 8.3 – Manufacture of caustic soda: saturated brine (NaCl solution) → Castner-Kellner cell → branch A: at carbon anode, Cl⁻ → ½Cl₂ collected as chlorine gas; branch B: at mercury cathode, Na⁺ + e⁻ → sodium amalgam → amalgam run into water → NaOH solution + H₂ gas + mercury → mercury returned to the cell
Properties
Sodium hydroxide is a white deliquescent solid, having a melting point of 591 K. It is highly water soluble and gives a strongly alkaline solution. The surface of the solution absorbs atmospheric \(\mathrm{CO_2}\) to form \(\mathrm{Na_2CO_3}\).
Deliquescent, as you saw in section 8.2.5 with lithium chloride, means the solid takes up so much moisture from the air that it eventually dissolves in it. A pellet of caustic soda left on a watch glass turns into a wet film within minutes, which is why it is always weighed quickly and stored in a tightly closed bottle.
Uses
i. Sodium hydroxide is used in the purification of bauxite (the aluminium ore).
ii. It is used in the commercial production of soap, paper, artificial silk and many chemicals.
iii. It is used for mercerising cotton fabrics.
iv. It is used in petroleum refining.
v. It is an important laboratory reagent.
8.3.3 Calcium Carbonate (\(\mathrm{CaCO_3}\))
Calcium carbonate is found in nature as chalk, limestone and marble. All three are the same compound; they differ in how the crystals were formed and packed, which is why one crumbles on a blackboard and another can be polished into a monument.
Preparation
i. When carbon dioxide is bubbled through a solution of calcium hydroxide (slaked lime), water insoluble solid calcium carbonate is formed.
\[\mathrm{Ca(OH)_2}\text{(aq)} + \mathrm{CO_2}\text{(g)} \rightarrow \mathrm{CaCO_3}\text{(s)} + \mathrm{H_2O}\text{(l)}\]
Excess carbon dioxide transforms the precipitate of \(\mathrm{CaCO_3}\) into water soluble calcium bicarbonate and therefore has to be avoided. This is the familiar limewater test seen in reverse: the milkiness that appears when \(\mathrm{CO_2}\) is passed into limewater disappears again if the gas is passed for too long.
ii. When a solution of calcium chloride is added to a solution of sodium carbonate, calcium carbonate is formed as a precipitate.
\[\mathrm{CaCl_2}\text{(aq)} + \mathrm{Na_2CO_3}\text{(aq)} \rightarrow \mathrm{CaCO_3}\text{(s)} + 2\,\mathrm{NaCl}\text{(aq)}\]
Properties
Calcium carbonate is a soft, light, white powder. It is practically water insoluble. On heating to 1200 K calcium carbonate decomposes into calcium oxide and carbon dioxide.
\[\mathrm{CaCO_3}\text{(s)} \xrightarrow{\ 1200\ \text{K}\ } \mathrm{CaO}\text{(s)} + \mathrm{CO_2}\uparrow\]
It reacts with dilute acids to give the corresponding calcium salt and carbon dioxide.
\[\mathrm{CaCO_3} + 2\,\mathrm{HCl} \rightarrow \mathrm{CaCl_2} + \mathrm{CO_2}\uparrow + \mathrm{H_2O}\]
\[\mathrm{CaCO_3} + \mathrm{H_2SO_4} \rightarrow \mathrm{CaSO_4} + \mathrm{CO_2}\uparrow + \mathrm{H_2O}\]
The brisk effervescence of carbon dioxide in these two reactions is the standard laboratory test for a carbonate, and it is also the reason marble and limestone buildings are slowly eaten away by acidic rain.
Uses
i. Calcium carbonate in the form of marble is used as building material.
ii. Calcium carbonate is used in the manufacture of quicklime (CaO), which is the major ingredient of cement.
iii. A mixture of \(\mathrm{CaCO_3}\) and \(\mathrm{MgCO_3}\) is used as a flux in the extraction of metals from ores.
iv. It is required for the manufacture of high quality paper.
v. It is an important ingredient in toothpaste, chewing gum, dietary supplements of calcium and a filler in cosmetics.
A flux, in the third use, is a substance added to a furnace to combine with the earthy impurities of the ore and carry them off as an easily removed slag — the same role you will meet again in the extraction of metals.
8.3.4 Hydrogen peroxide (\(\mathrm{H_2O_2}\))
Hydrogen peroxide is a low cost, clean and mild oxidising agent. A 30 % aqueous solution of hydrogen peroxide is commercially available. “Clean” here has a precise meaning: when \(\mathrm{H_2O_2}\) acts as an oxidising agent the by-product it leaves behind is water, so nothing has to be filtered out or disposed of afterwards.
Preparation
i. Hydrated barium peroxide is treated with ice cold dilute sulfuric acid. The precipitate of barium sulphate formed is filtered off to get the hydrogen peroxide solution.
\[\mathrm{BaO_2 \cdot 8H_2O}\text{(s)} + \mathrm{H_2SO_4}\text{(aq)} \xrightarrow[\ 273\ \text{K}\ ]{\ \text{low temp}\ } \mathrm{BaSO_4}\downarrow + \mathrm{H_2O_2}\text{(aq)} + 8\,\mathrm{H_2O}\text{(l)}\]
The low temperature is essential, because hydrogen peroxide decomposes readily when warm; and the insolubility of barium sulphate is what allows the product to be separated by simple filtration.
ii. A small quantity of sodium peroxide is added to an ice-cold solution of dilute sulfuric acid with stirring, which gives hydrogen peroxide (Merck process).
\[\mathrm{Na_2O_2}\text{(aq)} + \mathrm{H_2SO_4}\text{(aq)} \xrightarrow{\ 273\ \text{K}\ } \mathrm{H_2O_2}\text{(aq)} + \mathrm{Na_2SO_4}\text{(aq)}\]
iii. A 50 % solution of sulfuric acid is subjected to electrolytic oxidation to form peroxydisulfuric acid at the anode.
\[2\,\mathrm{HSO_4^{-}} \xrightarrow{\ \text{Electrolysis}\ } \mathrm{H_2S_2O_8} + 2\,e^{-}\]
Hydrolysis of the peroxydisulfuric acid yields hydrogen peroxide.
\[\mathrm{HO}\text{-}\mathrm{SO_2}\text{-}\mathrm{O}\text{-}\mathrm{O}\text{-}\mathrm{SO_2}\text{-}\mathrm{OH} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{H_2SO_4} + \mathrm{H_2O_2}\]
Notice the O–O linkage written out in the middle of the peroxydisulfuric acid formula. That oxygen–oxygen bond is carried through unbroken into the product; it is the defining feature of every peroxide. This method can be extended to the laboratory preparation of \(\mathrm{D_2O_2}\), simply by using heavy water in place of ordinary water in the hydrolysis step.
iv. Industrially hydrogen peroxide is prepared by air-oxidation of 2-ethylanthraquinol. The 2-ethylanthraquinol is regenerated by catalytic hydrogenation of 2-ethylanthraquinone.
\[\text{2-ethylanthraquinol} \xrightarrow[\ \text{air}\ ]{\ \mathrm{O_2}\ } \mathrm{H_2O_2} + \text{2-ethylanthraquinone}\]
The two steps together form a closed cycle, drawn in Fig. 8.9 and set out as a sequence in FC 8.4.
Diagram
Fig. 8.9 – The anthraquinone cycle: structure of 2-ethylanthraquinol (fused three-ring system with two OH groups and a C₂H₅ side chain) with a curved arrow labelled “O₂, air” to the structure of 2-ethylanthraquinone (same ring system with two C=O groups), releasing H₂O₂; a second curved arrow labelled “H₂/Pd, Δ” returning the quinone to the quinol, closing the loop.
Flowchart
FC 8.4 – Industrial preparation of H₂O₂ by the anthraquinone process: 2-ethylanthraquinol + O₂ from air → H₂O₂ + 2-ethylanthraquinone → H₂O₂ separated out as product; 2-ethylanthraquinone + H₂ over Pd catalyst → 2-ethylanthraquinol → returns to the start of the cycle
The elegance of this industrial route is that the organic compound is never consumed. It picks up hydrogen in one step, hands it to oxygen as hydrogen peroxide in the next, and returns to where it began — so the true raw materials are only hydrogen and air.
Properties
i. Pure \(\mathrm{H_2O_2}\) is a very pale blue coloured liquid, having a boiling point of 272.4 K.
ii. \(\mathrm{H_2O_2}\) is miscible with water and forms a hydrate, \(\mathrm{H_2O_2 \cdot H_2O}\).
iii. The strength of an aqueous solution of \(\mathrm{H_2O_2}\) is expressed in ‘volume’ units. The commercially marketed 30 % (by mass) solution of \(\mathrm{H_2O_2}\) has a volume strength of 100 volume. It means that 1 mL of the 30 % solution of \(\mathrm{H_2O_2}\) will give 100 mL of oxygen at STP.
iv. \(\mathrm{H_2O_2}\) acts as a mild oxidising as well as reducing agent.
a. Oxidising action of \(\mathrm{H_2O_2}\) in acidic medium
\[2\,\mathrm{Fe^{2+}}\text{(aq)} + 2\,\mathrm{H^{+}}\text{(aq)} + \mathrm{H_2O_2}\text{(aq)} \rightarrow 2\,\mathrm{Fe^{3+}}\text{(aq)} + 2\,\mathrm{H_2O}\text{(l)}\]
b. Reducing action of \(\mathrm{H_2O_2}\) in acidic medium
\[2\,\mathrm{MnO_4^{-}} + 6\,\mathrm{H^{+}} + 5\,\mathrm{H_2O_2} \rightarrow 2\,\mathrm{Mn^{2+}} + 8\,\mathrm{H_2O} + 5\,\mathrm{O_2}\]
Uses
i. Hydrogen peroxide is used as a mouthwash, germicide, mild antiseptic, preservative for milk and wine, and bleaching agent for soft materials, due to its mild oxidising property.
ii. Hydrogen peroxide, due to its reducing property, is used as an antichlor to remove excess chlorine from fabrics which have been bleached by chlorine.
iii. Nowadays it is also used in environmental chemistry for pollution control and restoration of aerobic conditions in sewage water.
8.3.5 Lithium aluminium hydride (\(\mathrm{LiAlH_4}\))
Lithium aluminium hydride is commonly abbreviated as LAH. It has the chemical formula \(\mathrm{LiAlH_4}\).
Preparation
Lithium hydride is treated with aluminium chloride to give lithium aluminium hydride.
\[4\,\mathrm{LiH} + \mathrm{AlCl_3} \rightarrow \mathrm{LiAlH_4} + 3\,\mathrm{LiCl}\]
Properties
Lithium aluminium hydride is a colourless solid. It reacts violently with water and even with atmospheric moisture. The reason lies in the hydride ion it carries: \(\mathrm{H^{-}}\) and the hydrogen of water are at opposite extremes, \(-1\) and \(+1\), and when they meet they combine at once to give hydrogen gas. Every reaction using LAH is therefore carried out in a dry solvent, and dry ether is the usual choice.
Uses
(i) LAH is a source of hydride and is therefore used as a reducing agent in organic synthesis. It reduces an ester to a primary alcohol, and likewise reduces a carboxylic acid to a primary alcohol, as shown in Fig. 8.10.
\[\mathrm{RCOOR'} \xrightarrow[\ \text{dry ether}\ ]{\ \mathrm{LiAlH_4}\ } \mathrm{RCH_2OH} + \mathrm{R'OH}\]
\[\mathrm{RCOOH} \xrightarrow[\ \text{dry ether}\ ]{\ \mathrm{LiAlH_4}\ } \mathrm{RCH_2OH} + \mathrm{H_2O}\]
Diagram
Fig. 8.10 – Two reduction schemes drawn with structural formulae: (a) an ester R–CO–OR′ with LiAlH₄/dry ether over the arrow giving R–CH₂–OH plus R′–OH; (b) a carboxylic acid R–CO–OH with LiAlH₄/dry ether over the arrow giving R–CH₂–OH plus H₂O; mark the C=O group in each starting material and the new –CH₂–OH group in each product.
In both cases the carbonyl carbon, which was joined to oxygen by a double bond, ends up as a \(\mathrm{-CH_2OH}\) group. That is what “reduction” means for an organic compound: hydrogen has been added and oxygen removed.
(ii) LAH is useful to prepare \(\mathrm{PH_3}\) (phosphine).
\[4\,\mathrm{PCl_3} + 3\,\mathrm{LiAlH_4} \rightarrow 4\,\mathrm{PH_3} + 3\,\mathrm{AlCl_3} + 3\,\mathrm{LiCl}\]
Here the same hydride-donating ability is put to a different use: each chlorine on phosphorus is replaced by a hydrogen, and the chlorine leaves with the aluminium and lithium.
8.4 Summary
Hydrogen: structure and position
- Hydrogen has the simplest atomic structure of all elements: a nucleus of charge \(+1\) and one extranuclear electron. It pairs its electron easily to form a covalent bond and exists as \(\mathrm{H_2}\), called dihydrogen.
- Its configuration \(1s^1\) resembles \(ns^1\) of group 1 and also \(ns^2np^5\) of group 17, since both hydrogen and a halogen are one electron short of an inert gas configuration. So some properties match the alkali metals and some match the halogens.
- \(\mathrm{H^{+}}\) is a bare proton and cannot exist freely; it always attaches to another molecule, as in \(\mathrm{H^{+}} + \mathrm{H_2O} \rightarrow \mathrm{H_3O^{+}}\). Hydrogen is therefore placed separately above group 1.
- Hydrogen is the most abundant element in the universe (70 % of its total mass) and the principal element of the solar system. On the earth it is tenth most abundant by mass but third by number of atoms.
Isotopes of hydrogen
- Isotopes are atoms of the same element with the same atomic number but different mass numbers. Hydrogen has three: protium (\(^{1}_{1}\mathrm{H}\), 99.98 %, stable), deuterium (\(^{2}_{1}\mathrm{H}\), 0.015 %, stable) and tritium (\(^{3}_{1}\mathrm{H}\), trace, radioactive).
- All three have one proton and one electron and differ only in neutron number, so their chemistry is essentially identical. Tritium has a half life period of 12.4 years and emits low energy \(\beta\) particles.
Preparation of dihydrogen
- Laboratory: \(\mathrm{Zn}\text{(s)} + 2\,\mathrm{HCl}\text{(aq)} \rightarrow \mathrm{ZnCl_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\), and also \(\mathrm{Zn}\text{(s)} + 2\,\mathrm{NaOH}\text{(aq)} \rightarrow \mathrm{Na_2ZnO_2}\text{(aq)} + \mathrm{H_2}\text{(g)}\).
- Electrolysis of water needs a trace of acid or alkali because pure water conducts poorly; it gives two volumes of hydrogen at the cathode to one volume of oxygen at the anode. Dihydrogen of over 99.5 % purity comes from electrolysis of warm barium hydroxide solution between nickel electrodes.
- Industrial production today is about 77 % from petrochemicals, 18 % from coal, 4 % electrolytic and 1 % by other methods.
Flowchart
FC 8.5 – Industrial dihydrogen from carbon or hydrocarbon, condensed: CH₄ or coke + steam → Stage 1 at 1270 K over Ni → water-gas (CO + H₂), also called syngas → Stage 2, water-gas shift with steam at 673 K over iron chromate → CO₂ + H₂ → Stage 3, CO₂ scrubbed out with sodium arsenite solution → dihydrogen
Properties and uses of dihydrogen
- Physical: colourless, tasteless, odourless, nonpolar, water insoluble, lighter than air, burns with a pale blue flame.
- The H–H bond dissociation energy is high (436 kJ mol\(^{-1}\)), so its reactions need high temperature and/or a catalyst.
- With metals it forms hydrides, \(2\,\mathrm{Na}\text{(s)} + \mathrm{H_2}\text{(g)} \xrightarrow{\Delta} 2\,\mathrm{NaH}\text{(s)}\) — behaving like a halogen. With halogens it forms \(\mathrm{H_2} + \mathrm{X_2} \rightarrow 2\,\mathrm{HX}\) — behaving like an alkali metal; the vigour falls from fluorine to iodine.
- With dioxygen: \(2\,\mathrm{H_2}\text{(g)} + \mathrm{O_2}\text{(g)} \rightarrow 2\,\mathrm{H_2O}\text{(l)}\), \(\Delta H = -235\ \text{kJ mol}^{-1}\), highly exothermic.
- It is a reducing agent for oxides and ions of metals less reactive than iron, and hydrogenates unsaturated oils over nickel to give vanaspati ghee.
- Uses: ammonia production (largest), vanaspati ghee, rocket fuel as liquid dihydrogen, bulk methanol, and preparation of HCl and metal hydrides.
Elements of groups 1 and 2
- Both groups are s-block: the last electron enters the \(ns\) subshell, giving \(ns^1\) for group 1 and \(ns^2\) for group 2. The loosely held valence s-electrons are lost easily, so these metals are never found free in nature.
- Group 1 (except hydrogen) are the alkali metals: Li, Na, K, Rb, Cs, Fr. Group 2 are the alkaline earth metals: Be, Mg, Ca, Sr, Ba, Ra, named because they occur as minerals in rocks.
- \(\mathrm{M^{+}}\) of group 1 and \(\mathrm{M^{2+}}\) of group 2 have inert gas configurations with no unpaired electrons, so their compounds are diamagnetic and colourless.
- Down both groups: atomic radius, ionic radius and density increase; ionization enthalpy and electronegativity decrease; standard reduction potentials are highly negative.
- Alkali metals are silvery white, soft and the most electropositive elements; alkaline earth metals are harder and slightly less electropositive. Li and Be stand apart from their families because of very small size and comparatively high electronegativity.
Chemical properties of groups 1 and 2
- With oxygen, group 1 gives oxide, peroxide or superoxide depending on cation size: \(4\,\mathrm{Li} + \mathrm{O_2} \rightarrow 2\,\mathrm{Li_2O}\); \(2\,\mathrm{Na} + \mathrm{O_2} \rightarrow \mathrm{Na_2O_2}\); \(\mathrm{K} + \mathrm{O_2} \rightarrow \mathrm{KO_2}\). These oxides are strongly basic. \(\mathrm{KO_2}\) absorbs \(\mathrm{CO_2}\) and releases \(\mathrm{O_2}\): \(4\,\mathrm{KO_2} + 2\,\mathrm{CO_2} \rightarrow 2\,\mathrm{K_2CO_3} + 3\,\mathrm{O_2}\).
- Group 2 metals are protected by a surface oxide film but burn on ignition to MO type oxides mixed with nitride: \(2\,\mathrm{Mg} + \mathrm{O_2} \rightarrow 2\,\mathrm{MgO}\) and \(3\,\mathrm{Mg} + \mathrm{N_2} \rightarrow \mathrm{Mg_3N_2}\).
- With water: \(2\,\mathrm{Na} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + \mathrm{H_2}\); lithium reacts slowly, sodium and potassium vigorously. Beryllium does not react, magnesium decomposes hot water, the rest react with cold water: \(\mathrm{Ca} + 2\,\mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2} + \mathrm{H_2}\).
- With hydrogen: \(2\,\mathrm{M} + \mathrm{H_2} \xrightarrow{673\ \text{K}} 2\,\mathrm{M^{+}H^{-}}\) for group 1; \(\mathrm{M} + \mathrm{H_2} \xrightarrow{\Delta} \mathrm{MH_2}\) for group 2 except beryllium.
- With halogens: \(2\,\mathrm{M} + \mathrm{X_2} \rightarrow 2\,\mathrm{M^{+}X^{-}}\) and \(\mathrm{M} + \mathrm{X_2} \rightarrow \mathrm{MX_2}\).
- Reducing nature is measured by \(E^\circ\) for \(\mathrm{M^{+}}\text{(aq)} + e^{-} \rightarrow \mathrm{M}\text{(s)}\); all these metals have high negative values, lithium being the most powerful reducing agent in group 1 and sodium the least. Group 2 metals are strong reducing agents but weaker than the alkali metals.
- Both groups dissolve in liquid ammonia to give deep blue conducting solutions containing ammoniated electrons, which are paramagnetic and on standing give the amide with liberation of hydrogen, turning bronze and diamagnetic.
Diagonal relationship
- Lithium resembles magnesium, and beryllium resembles aluminium — pairs placed diagonally in the periodic table.
- Li and Mg have very similar atomic and ionic radii; Be and Al have very similar charge to radius ratios of their ions (Be \(\frac{2}{31}\), Al \(\frac{3}{53.55}\)).
- Evidence: Li and Mg both form a nitride in air, both have carbonates that decompose on heating to the oxide and \(\mathrm{CO_2}\), and both have deliquescent, hydrated chlorides. \(\mathrm{BeCl_2}\) and \(\mathrm{AlCl_3}\) are both covalent, chlorine-bridged, strong Lewis acids soluble in organic solvents, and BeO and \(\mathrm{Al_2O_3}\) are both amphoteric.
Uses and biological importance
- Uses include lithium in long-life batteries, liquid sodium for heat transfer in nuclear power stations, KCl as fertilizer, \(\mathrm{KO_2}\) for oxygen generation, caesium in photoelectric cells, beryllium as a nuclear moderator, Mg–Al alloys in aircraft, calcium for bones and teeth, barium sulphate as a barium meal and radium in radiotherapy.
- \(\mathrm{Na^{+}}\) dominates extracellular fluid, regulating water flow across membranes and carrying nerve signals; \(\mathrm{K^{+}}\) is the most abundant ion inside cells, needed for protein synthesis and glucose oxidation.
- \(\mathrm{Mg^{2+}}\) is part of chlorophyll and works in glucose and fat breakdown, enzyme-assisted protein synthesis and cholesterol regulation; \(\mathrm{Ca^{2+}}\) builds bone and teeth as apatite \([\mathrm{Ca_3(PO_4)_2}]\) and is needed for blood clotting, muscle action and regular heartbeat.
Sodium carbonate, \(\mathrm{Na_2CO_3 \cdot 10H_2O}\)
- Made by the Solvay process; the key step is the low solubility of \(\mathrm{NaHCO_3}\), which precipitates and is then heated: \(2\,\mathrm{NaHCO_3}\text{(s)} \xrightarrow{\Delta} \mathrm{Na_2CO_3}\text{(s)} + \mathrm{H_2O}\text{(g)} + \mathrm{CO_2}\text{(g)}\).
- Ammonia is recovered with slaked lime, \(2\,\mathrm{NH_4Cl}\text{(aq)} + \mathrm{Ca(OH)_2}\text{(s)} \rightarrow 2\,\mathrm{NH_3}\text{(g)} + \mathrm{CaCl_2}\text{(aq)} + \mathrm{H_2O}\text{(l)}\). Potassium carbonate cannot be made this way because \(\mathrm{KHCO_3}\) is highly soluble.
- On heating, the decahydrate goes to the monohydrate at 373 K and then to anhydrous soda-ash above 373 K. Its solution is alkaline through hydrolysis: \(\mathrm{Na_2CO_3} + \mathrm{H_2O} \rightarrow \mathrm{NaHCO_3} + \mathrm{NaOH}\).
- Used as a cleaning material, a water softener, in soap and caustic soda manufacture and as a laboratory reagent.
Flowchart
FC 8.6 – Solvay process condensed: brine saturated with NH₃ + CO₂ → NH₄HCO₃ → double decomposition with NaCl → NaHCO₃ precipitate + NH₄Cl solution → NaHCO₃ heated → Na₂CO₃ + H₂O + CO₂ (CO₂ recycled); NH₄Cl + Ca(OH)₂ → NH₃ recycled + CaCl₂ by-product
Sodium hydroxide, NaOH
- Made by electrolysis of saturated brine in a Castner-Kellner cell with a mercury cathode and carbon anode: \(\mathrm{Na^{+}} + e^{-} \xrightarrow{\mathrm{Hg}}\) Na-amalgam at the cathode, \(\mathrm{Cl^{-}} \rightarrow \tfrac{1}{2}\mathrm{Cl_2} + e^{-}\) at the anode; then \(2\,\text{Na-Hg} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH} + 2\,\mathrm{Hg} + \mathrm{H_2}\).
- A white deliquescent solid of melting point 591 K, highly soluble, strongly alkaline, and its surface absorbs atmospheric \(\mathrm{CO_2}\) to form \(\mathrm{Na_2CO_3}\).
- Used in bauxite purification, soap, paper and artificial silk manufacture, mercerising cotton, petroleum refining and as a laboratory reagent.
Calcium carbonate, \(\mathrm{CaCO_3}\)
- Occurs as chalk, limestone and marble; prepared by passing \(\mathrm{CO_2}\) into slaked lime (avoiding excess, which gives soluble bicarbonate) or by adding \(\mathrm{CaCl_2}\) solution to \(\mathrm{Na_2CO_3}\) solution.
- Practically insoluble in water; decomposes on strong heating, \(\mathrm{CaCO_3}\text{(s)} \xrightarrow{1200\ \text{K}} \mathrm{CaO}\text{(s)} + \mathrm{CO_2}\), and reacts with dilute acids to give the calcium salt, \(\mathrm{CO_2}\) and water.
- Used as building material, for quicklime and cement, as a flux in metal extraction, in high quality paper, and in toothpaste, chewing gum, calcium supplements and cosmetics.
Hydrogen peroxide, \(\mathrm{H_2O_2}\)
- A low cost, clean, mild oxidising agent, sold as a 30 % aqueous solution.
- Prepared from \(\mathrm{BaO_2 \cdot 8H_2O}\) with ice-cold dilute \(\mathrm{H_2SO_4}\), from \(\mathrm{Na_2O_2}\) with ice-cold dilute \(\mathrm{H_2SO_4}\) (Merck process), by electrolytic oxidation of 50 % \(\mathrm{H_2SO_4}\) to peroxydisulfuric acid followed by hydrolysis, and industrially by air-oxidation of 2-ethylanthraquinol, which is regenerated by catalytic hydrogenation.
- A very pale blue liquid of boiling point 272.4 K, miscible with water, forming the hydrate \(\mathrm{H_2O_2 \cdot H_2O}\). Its strength is expressed in volume units: the 30 % solution is 100 volume, meaning 1 mL gives 100 mL of oxygen at STP.
- Acts as an oxidising agent, \(2\,\mathrm{Fe^{2+}} + 2\,\mathrm{H^{+}} + \mathrm{H_2O_2} \rightarrow 2\,\mathrm{Fe^{3+}} + 2\,\mathrm{H_2O}\), and as a reducing agent, \(2\,\mathrm{MnO_4^{-}} + 6\,\mathrm{H^{+}} + 5\,\mathrm{H_2O_2} \rightarrow 2\,\mathrm{Mn^{2+}} + 8\,\mathrm{H_2O} + 5\,\mathrm{O_2}\).
- Used as mouthwash, germicide, antiseptic, preservative and bleach; as an antichlor; and in pollution control and restoring aerobic conditions in sewage water.
Lithium aluminium hydride, \(\mathrm{LiAlH_4}\) (LAH)
- Prepared as \(4\,\mathrm{LiH} + \mathrm{AlCl_3} \rightarrow \mathrm{LiAlH_4} + 3\,\mathrm{LiCl}\); a colourless solid that reacts violently with water and even atmospheric moisture.
- A source of hydride, so it is used as a reducing agent in organic synthesis, converting esters and carboxylic acids into alcohols in dry ether, and it is used to prepare phosphine from \(\mathrm{PCl_3}\).
8.5 Check Your Concepts
Hydrogen is printed at the top of group 1, yet it is drawn slightly apart from the alkali metals. Give two properties of hydrogen that support its placement in group 1 and two that support its placement in group 17.
Why can \(\mathrm{H^{+}}\) not exist freely in aqueous solution, while \(\mathrm{Na^{+}}\) can? Write the equation for what happens to \(\mathrm{H^{+}}\) in water.
All three isotopes of hydrogen react with chlorine in exactly the same way. Explain why, in terms of what decides chemical behaviour.
In the industrial preparation of dihydrogen, Stage 1 needs 1270 K but Stage 2 is carried out at 673 K. Explain the reason for the large difference, and state the purpose of Stage 2.
Why is a trace of acid or alkali added before water is electrolysed to obtain dihydrogen? Why does the gas collected at the cathode occupy twice the volume of the gas collected at the anode?
Assertion: Reactions of dihydrogen are usually carried out at high temperature or in the presence of a catalyst. Reason: The bond dissociation energy of the H–H bond is high. State whether the assertion and reason are correct, and whether the reason explains the assertion.
In the reaction \(2\,\mathrm{Na}\text{(s)} + \mathrm{H_2}\text{(g)} \rightarrow 2\,\mathrm{NaH}\text{(s)}\), identify the species oxidised and the species reduced, and state which family — alkali metals or halogens — dihydrogen is imitating here.
Arrange \(\mathrm{F_2}\), \(\mathrm{Cl_2}\), \(\mathrm{Br_2}\) and \(\mathrm{I_2}\) in decreasing order of the vigour of their reaction with dihydrogen, and give the reason for the order.
Lithium forms the normal oxide, sodium forms the peroxide and potassium forms the superoxide on reaction with oxygen. Explain this pattern in terms of the size of the cation.
Why are the compounds of group 1 and group 2 metals colourless and diamagnetic?
Lithium has the highest ionization enthalpy among the alkali metals, and yet it is the strongest reducing agent in the group. Resolve this apparent contradiction.
Explain why the alkali metals and alkaline earth metals are never found in the free state in nature, and why electrolysis rather than heating is needed to obtain them as metals.
A solution of sodium in liquid ammonia is deep blue and paramagnetic, but on standing it turns bronze and becomes diamagnetic. Explain both observations and write the equation for the change taking place.
Complete the following reactions and balance them: (a) \(\mathrm{Ca} + 2\,\mathrm{H_2O} \rightarrow\) (b) \(3\,\mathrm{Mg} + \mathrm{N_2} \rightarrow\) (c) \(\mathrm{Li_2O} + \mathrm{H_2O} \rightarrow\) (d) \(4\,\mathrm{KO_2} + 2\,\mathrm{CO_2} \rightarrow\)
Beryllium does not react with water, magnesium reacts only with hot water, and calcium reacts with cold water. Explain the trend using the atomic properties of these three metals.
State what is meant by a diagonal relationship, and explain why a diagonal move in the periodic table leaves two elements with similar atomic properties.
Match the following:
| Column A | Column B |
|---|---|
| a. \(\mathrm{KO_2}\) | i. barium meal for intestinal X-ray |
| b. \(\mathrm{BaSO_4}\) | ii. moderator in nuclear reactors |
| c. Be | iii. photoelectric cells |
| d. Cs | iv. oxygen generation in breathing equipment |
Why is BeO described as amphoteric while MgO is described as basic? Write one equation in support of each description.
The Solvay process depends entirely on one physical property of sodium bicarbonate. Name that property, explain its role, and state why the same process fails for potassium carbonate.
In the Castner-Kellner cell, why is mercury used as the cathode instead of an ordinary metal plate? Write the reaction occurring at each electrode.
Excess carbon dioxide must be avoided when calcium carbonate is prepared by passing \(\mathrm{CO_2}\) into slaked lime. Explain what goes wrong, and connect your answer to the milkiness of limewater disappearing on prolonged passage of \(\mathrm{CO_2}\).
Hydrogen peroxide acts as an oxidising agent in one reaction and as a reducing agent in another. Using the oxidation state of oxygen in \(\mathrm{H_2O_2}\), explain how both are possible, and state what visible difference you would notice in the two cases.
Assertion: Hydrogen peroxide is stored in dark coloured bottles and kept cool. Reason: Hydrogen peroxide decomposes into water and oxygen, and light and heat speed up the decomposition. State whether the assertion and reason are correct, and whether the reason explains the assertion.
Why must reductions with \(\mathrm{LiAlH_4}\) be carried out in dry ether? Explain in terms of the hydride ion it supplies.
Sodium metal catches fire in water, yet sodium ions are essential in the human body. Explain how the same element can be both, referring to electronic configuration.
Arrange Li, Na, K, Rb and Cs in increasing order of (a) atomic radius and (b) first ionization enthalpy, and give the single reason that explains both orders.
Explain the following
- Hydrogen shows similarity with alkali metals as well as halogens.
- Standard reduction potentials of alkali metals have high negative values.
- Alkaline earth metals have low values of electronegativity, which decrease down the group.
- Sodium dissolves in liquid ammonia to form a solution which shows electrical conductivity.
- \(\mathrm{BeCl_2}\) is covalent while \(\mathrm{MgCl_2}\) is ionic.
- Lithium floats on water while sodium floats and catches fire when put in water.
\(\mathrm{CO_2}\) is passed into a concentrated solution of NaCl which is saturated with \(\mathrm{NH_3}\).
A 50 % solution of sulphuric acid is subjected to electrolytic oxidation and the product is hydrolysed.
Magnesium is heated in air.
Beryllium oxide is treated separately with aqueous HCl and aqueous NaOH solutions.
Describe the diagonal relationship between Li and Mg with the help of two illustrative properties.
Describe the industrial production of dihydrogen from steam. Also write the chemical reactions involved.
A water sample, which did not give lather with soap, was found to contain \(\mathrm{Ca(HCO_3)_2}\) and \(\mathrm{Mg(HCO_3)_2}\). Which chemical will make this water give lather with soap? Explain with the help of chemical reactions.
Name the isotopes of hydrogen. Write their atomic composition schematically and explain which of these is radioactive.
Name the following
- Alkali metal with the smallest atom.
- The most abundant element in the universe.
- Radioactive alkali metal.
- Ions having high concentration in cell sap.
- A compound having hydrogen, aluminium and lithium as its constituent elements.
8.6 Practice With Numericals
Calculate the volume of dihydrogen obtained at STP when 13 g of zinc reacts completely with an excess of aqueous sodium hydroxide. (Zn \(= 65\ \text{g mol}^{-1}\); molar volume at STP \(= 22.7\ \text{L mol}^{-1}\))
Calculate the mass of sodium hydride formed when 2.3 g of sodium is heated in an excess of dihydrogen. (Na \(= 23\), H \(= 1\ \text{g mol}^{-1}\))
50 g of calcium carbonate is heated to 1200 K until decomposition is complete. Calculate the mass of calcium oxide formed and the volume of carbon dioxide obtained at STP. (\(\mathrm{CaCO_3} = 100\), CaO \(= 56\ \text{g mol}^{-1}\); molar volume at STP \(= 22.7\ \text{L mol}^{-1}\))
In a breathing apparatus, 28.4 g of potassium superoxide reacts with carbon dioxide according to \(4\,\mathrm{KO_2} + 2\,\mathrm{CO_2} \rightarrow 2\,\mathrm{K_2CO_3} + 3\,\mathrm{O_2}\). Calculate the mass of carbon dioxide absorbed and the volume of oxygen released at STP. (\(\mathrm{KO_2} = 71\), \(\mathrm{CO_2} = 44\ \text{g mol}^{-1}\); molar volume at STP \(= 22.7\ \text{L mol}^{-1}\))
Calculate the percentage by mass of a hydrogen peroxide solution which is 11.35 volume. (\(\mathrm{H_2O_2} = 34\ \text{g mol}^{-1}\); molar volume at STP \(= 22.7\ \text{L mol}^{-1}\))
Calculate the percentage by mass of water of crystallisation in \(\mathrm{MgCl_2 \cdot 8H_2O}\). (Mg \(= 24\), Cl \(= 35.5\), O \(= 16\), H \(= 1\ \text{g mol}^{-1}\))
Determine the oxidation state of hydrogen in NaH and in \(\mathrm{H_2O}\), and the oxidation state of nitrogen in \(\mathrm{Li_3N}\).
Using the spin-only formula \(\mu = \sqrt{n(n+2)}\ \text{BM}\), calculate the magnetic moment of the \(\mathrm{Na^{+}}\) ion and of the \(\mathrm{Mg^{2+}}\) ion, and state what your answers tell you about the magnetic behaviour of their compounds.
One litre of a hard water sample contains 0.01 mol of \(\mathrm{Ca(HCO_3)_2}\). Calculate the mass of anhydrous sodium carbonate required to precipitate all the calcium as \(\mathrm{CaCO_3}\). (\(\mathrm{Na_2CO_3} = 106\ \text{g mol}^{-1}\))
In the electrolysis of acidified water, 5.6 L of dihydrogen is collected at the cathode at STP. Calculate the volume of dioxygen collected at the anode under the same conditions, and state the mole ratio you have used.
Calculate the mass of lithium aluminium hydride obtained when 8 g of lithium hydride reacts completely with aluminium chloride according to \(4\,\mathrm{LiH} + \mathrm{AlCl_3} \rightarrow \mathrm{LiAlH_4} + 3\,\mathrm{LiCl}\). (Li \(= 7\), H \(= 1\), Al \(= 27\ \text{g mol}^{-1}\))
Calculate the number of neutrons present in 1 g of tritium, \(^{3}_{1}\mathrm{H}\). (Atomic mass of tritium \(= 3\ \text{g mol}^{-1}\); \(N_A = 6.022 \times 10^{23}\ \text{mol}^{-1}\))
8.7 Objective Questions
- The unstable isotope of hydrogen is …..
- H-1
- H-2
- H-3
- H-4
- Identify the odd one.
- Rb
- Ra
- Sr
- Be
- Which of the following is a Lewis acid?
- \(\mathrm{BaCl_2}\)
- KCl
- \(\mathrm{BeCl_2}\)
- LiCl
- What happens when crystalline \(\mathrm{Na_2CO_3}\) is heated?
- releases \(\mathrm{CO_2}\)
- loses \(\mathrm{H_2O}\)
- decomposes into \(\mathrm{NaHCO_3}\)
- colour changes