Chapter 6: Study Of The First Element — Hydrogen
In the summer of 1783, a crowd gathered on the Champ de Mars in Paris to watch a bag of silk rise into the sky. The bag had been built by Jacques Charles, a French physicist, and it was filled with a gas that almost nobody had a proper name for yet.
Making that gas was the hard part. Charles and his helpers packed barrels with scrap iron and poured dilute sulphuric acid over it. The mixture frothed, grew hot, and gave off a stream of bubbles, and those bubbles were led along pipes into the balloon. It took days. The gas came out damp and dirty, and the silk had to be sealed with rubber to stop it leaking away.
When the ropes were cut, the balloon climbed until it disappeared into the cloud. It came down some twenty kilometres away, near the village of Gonesse. The villagers had never seen such a thing fall out of the sky. They attacked it with pitchforks, tied it to a horse, and dragged it across the fields until it was in pieces.
The gas in that balloon was hydrogen — the lightest substance known, the first element in the periodic table, and the one you are about to study. It came out of a barrel of iron and acid, and that is very nearly how you will make it yourself.
Diagram
Fig. 6.0-1 — Charles’s Hydrogen Balloon Of 1783 Two-part illustration. Left: barrels of scrap iron with acid being poured in from a jug, bubbles rising inside, and pipes leading the gas away from the barrels; label ‘scrap iron and dilute sulphuric acid’ and ‘gas led to the balloon’. Right: the rubberised silk balloon, partly filled and held down by ropes, rising above a watching crowd on open ground; label ‘hydrogen-filled balloon’ and ‘holding ropes’. Period costume, plain line-art style.
6.1 Position Of Hydrogen In The Periodic Table
In the periodic table the elements are set out in order of increasing atomic number. Each vertical column is called a group, and each horizontal row is called a period. Elements in the same group behave in similar ways, because their atoms have the same number of electrons in the outermost shell.
Hydrogen comes first in that arrangement. Its atomic number is 1, so an atom of hydrogen has one proton in its nucleus and one electron outside it. That single electron lies in the K shell, which is the first and innermost shell.
The Hydrogen Atom
The outermost shell of an atom is called its valence shell, and the electrons in it are its valence electrons. Hydrogen has just one shell, so the K shell is also its valence shell, and it holds one valence electron. Since the atom has one electron available for bonding, the valency of hydrogen is 1 — it combines with one atom of hydrogen, or one atom of chlorine, at a time.
A single hydrogen atom is not stable on its own. Two of them join and share their electrons, so the gas exists as \(\\mathrm{H\_2}\) molecules. A molecule made of two atoms is described as diatomic.
Because the first shell of an atom is complete with two electrons, hydrogen sits one electron short of a full K shell. Keep that fact in mind — it explains a good deal of what follows.
Diagram
Fig. 6.1-1 — The Hydrogen Atom And The Hydrogen Molecule Two panels side by side. Left: a single hydrogen atom drawn as a small nucleus labelled ‘1 proton’ with one circular shell around it carrying one dot, shell labelled ‘K shell (valence shell)’. Right: two such atoms overlapping, with a shared pair of two dots between the nuclei, labelled ‘shared pair of electrons’ and captioned ‘hydrogen molecule, H₂’. Simple line-art with dots for electrons.
The basic facts about hydrogen are gathered together in Table 6.1. The relative atomic mass of an element tells you how heavy one of its atoms is compared with a standard atom, and the relative molecular mass does the same for one molecule. Both are simple ratios, so neither carries a unit.
Table 6.1 Hydrogen at a glance
| Property | Value |
|---|---|
| Symbol | \(\\mathrm{H}\) |
| Atomic number | 1 |
| Electronic configuration | 1 (K shell) |
| Valency | 1 |
| Molecular formula | \(\\mathrm{H\_2}\) |
| Relative atomic mass | 1 |
| Relative molecular mass | 2 |
| Position in the periodic table | Period 1, Group 1 (IA) |
Since hydrogen has one electron in its valence shell, and one shell in all, it is placed in the first group and the first period of the periodic table.
A Position Still Argued Over
You would expect an element placed in Group 1 to behave like the other members of that group — sodium, potassium and the rest, which are known as the alkali metals. Hydrogen does not. It is a colourless gas, while every other element in that column is a soft, shiny metal.
Ever since Mendeleev drew up his periodic table, the position of hydrogen in it has been a matter of debate. Some of its properties match those of the Group 1 alkali metals, while others match those of the Group 17 halogens — fluorine, chlorine, bromine and iodine. Group 17 was written as Group VIIA in the older style of numbering, and you will meet both forms.
Thomson suggested that hydrogen should be given a separate position of its own, at the top of the periodic table. Placed there, it belongs to no single group, and the symmetry of the modern table is left undisturbed.
Diagram
Fig. 6.1-2 — Where Hydrogen Sits Outline grid of the periodic table showing periods 1 to 7 and groups 1 to 18, with only the symbols of the first two groups and groups 13 to 18 filled in and the transition block left as a shaded rectangle labelled ‘transition elements’. Highlight the box for H at the head of group 1; shade the alkali metals Li, Na, K, Rb, Cs, Fr down group 1 in one tint and the halogens F, Cl, Br, I, At down group 17 in a second tint. Add two arrows from H, one pointing down group 1 and one pointing across to group 17, labelled ‘resembles the alkali metals’ and ‘resembles the halogens’.
This double resemblance is described by saying that hydrogen shows a dual nature. The two sections that follow set out each side of it in turn.
6.2 Similarities Between Hydrogen And The Alkali Metals
The elements of Group 1 below hydrogen — lithium, sodium, potassium, rubidium, caesium and francium — are called the alkali metals. Hydrogen is not a metal, yet in five or six clear ways it behaves as they do. Those resemblances are what keep it at the head of their column.
6.2.1 Electronic Configuration And Valence Electrons
Every one of these elements has exactly one electron in its outermost shell. The shells are named K, L, M, N, O and P as you move outwards from the nucleus, and Table 6.2 sets out how the electrons of each element are shared among them.
Table 6.2 Electron arrangement of hydrogen and the alkali metals
| Element (atomic number) | K | L | M | N | O | P |
|---|---|---|---|---|---|---|
| H (1) | 1 | |||||
| Li (3) | 2 | 1 | ||||
| Na (11) | 2 | 8 | 1 | |||
| K (19) | 2 | 8 | 8 | 1 | ||
| Rb (37) | 2 | 8 | 18 | 8 | 1 | |
| Cs (55) | 2 | 8 | 18 | 18 | 8 | 1 |
| Fr (87) | 2 | 8 | 18 | 32 | 18 | 8 |
Read down the last filled column of the table and you will see the same figure every time: 1. Whatever the size of the atom, the valence shell holds a single electron. That one electron is what the chemistry of all these elements turns on.
Diagram
Fig. 6.2-1 — One Valence Electron In Every Case Four shell diagrams in a row, drawn to increasing size: H (1), Li (2, 1), Na (2, 8, 1) and K (2, 8, 8, 1). Show the nucleus as a small circle with the symbol inside and concentric shells carrying dots for electrons. Colour or ring the single outermost electron in each atom in a contrasting tint and add one label beneath the row reading ‘one valence electron in every case’.
6.2.2 Valency And Ion Formation
Because each of these atoms has one electron to offer, each has a valency of 1. Hydrogen and every alkali metal therefore combine atom for atom with chlorine, giving \(\\mathrm{HCl}\), \(\\mathrm{LiCl}\), \(\\mathrm{NaCl}\) and so on.
An atom that loses an electron is left with one positive charge more than negative, and the particle formed is called a cation, a positively charged ion. Hydrogen and the alkali metals all form cations in this way.
hydrogen atom → hydrogen ion + electron
\[\\mathrm{H} \\rightarrow \\mathrm{H^{+}} + e^{-}\]
\[\\mathrm{Li} \\rightarrow \\mathrm{Li^{+}} + e^{-}\]
\[\\mathrm{Na} \\rightarrow \\mathrm{Na^{+}} + e^{-}\]
An element whose atoms give up electrons readily to form positive ions is said to be electropositive. Hydrogen and the alkali metals all show electropositive character.
6.2.3 Reducing Power
A substance that removes oxygen from another substance is acting as a reducing agent. Both hydrogen and the alkali metals do this. You will meet the idea in full in section 6.13; for the moment, watch where the oxygen goes.
When hydrogen is passed over heated copper(II) oxide, the black solid turns to reddish-brown copper and water is formed. The oxygen has been pulled away from the copper by the hydrogen.
copper(II) oxide + hydrogen → copper + water
\[\\mathrm{CuO} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Cu} + \\mathrm{H\_2O}\]
Sodium does the same job, taking the oxygen away as sodium oxide.
copper(II) oxide + sodium → copper + sodium oxide
\[\\mathrm{CuO} + 2,\\mathrm{Na} \\rightarrow \\mathrm{Cu} + \\mathrm{Na\_2O}\]
6.2.4 Burning In Oxygen
Hydrogen burns in oxygen and forms its oxide, which is water.
hydrogen + oxygen → water
\[2,\\mathrm{H\_2} + \\mathrm{O\_2} \\rightarrow 2,\\mathrm{H\_2O}\]
A small quantity of hydrogen mixed with air burns with a sharp pop sound. That pop is the sound of the whole sample burning at once, and it is the test by which hydrogen is identified.
The alkali metals also burn vigorously when heated in oxygen, but what they form depends on which metal is used. Lithium forms the simple oxide, called a monoxide.
lithium + oxygen → lithium oxide
\[4,\\mathrm{Li} + \\mathrm{O\_2} \\rightarrow 2,\\mathrm{Li\_2O}\]
Sodium burns to give a peroxide, an oxide that contains more oxygen than the ordinary one. Hydrogen has a peroxide of its own too — hydrogen peroxide, \(\\mathrm{H\_2O\_2}\) — which is one more point of resemblance between them.
sodium + oxygen → sodium peroxide
\[2,\\mathrm{Na} + \\mathrm{O\_2} \\rightarrow \\mathrm{Na\_2O\_2}\]
Potassium, rubidium and caesium go one step further and form superoxides, of the general formula \(\\mathrm{MO\_2}\), where M stands for the metal.
potassium + oxygen → potassium superoxide
\[\\mathrm{K} + \\mathrm{O\_2} \\rightarrow \\mathrm{KO\_2}\]
Tarnishing In Air
The alkali metals are so reactive that they do not even need to be heated. Left exposed to air, their bright surfaces quickly go dull, and this dulling is called tarnishing. Oxides form first, then hydroxides as the oxide meets moisture, and finally carbonates as the hydroxide meets carbon dioxide.
sodium + oxygen → sodium oxide
\[4,\\mathrm{Na} + \\mathrm{O\_2} \\rightarrow 2,\\mathrm{Na\_2O}\]
sodium oxide + water → sodium hydroxide
\[\\mathrm{Na\_2O} + \\mathrm{H\_2O} \\rightarrow 2,\\mathrm{NaOH}\]
sodium hydroxide + carbon dioxide → sodium carbonate + water
\[2,\\mathrm{NaOH} + \\mathrm{CO\_2} \\rightarrow \\mathrm{Na\_2CO\_3} + \\mathrm{H\_2O}\]
For this reason the alkali metals are never left standing in the open. They are stored under an inert organic solvent such as kerosene oil, which keeps air and moisture away from the metal surface. A jar of sodium in the laboratory is always a jar of sodium under liquid, as Fig. 6.2-2 shows.
Diagram
Fig. 6.2-2 — Sodium Stored Under Kerosene Oil Sectional view of a stoppered glass jar standing on a bench. The jar is filled with a clear liquid labelled ‘kerosene oil’; two dull grey lumps labelled ‘sodium’ rest at the bottom, fully covered by the liquid. Label the stopper ‘tight-fitting stopper’ and add a pair of tongs lying beside the jar labelled ‘handled only with tongs, never with the fingers’.
6.2.5 Combination With Non-Metals
Hydrogen and the alkali metals both combine with non-metals such as oxygen, sulphur and chlorine, and the formulae of the compounds they form run closely in parallel. Compare the two rows of Table 6.3.
Two atoms of hydrogen are needed for one of oxygen, and two atoms of sodium for one of oxygen, because both have a valency of 1 while oxygen has a valency of 2. The same reasoning gives the sulphides, and a one-to-one formula with chlorine. The likeness in the formulae is not a coincidence — it comes straight from that shared single valence electron.
6.3 Similarities Between Hydrogen And The Halogens
Group 17 holds fluorine, chlorine, bromine and iodine, and these elements are called the halogens. They sit at the far side of the periodic table from the alkali metals, and they are non-metals. Hydrogen resembles them almost as closely as it resembles the metals of Group 1.
6.3.1 Electronic Configuration And Valency
The elements of Group 18 are the inert gases — helium, neon, argon and the rest — whose atoms have complete outer shells and which therefore take almost no part in chemical change. A complete outer shell is the arrangement every other atom is working towards.
Hydrogen and the halogens are all one electron short of the nearest inert gas. Table 6.4 sets the arrangements side by side.
Table 6.4 Hydrogen and the halogens compared with the nearest inert gas
| Element | Electron arrangement | Nearest inert gas | Its arrangement |
|---|---|---|---|
| H | 1 | He | 2 |
| F | 2, 7 | Ne | 2, 8 |
| Cl | 2, 8, 7 | Ar | 2, 8, 8 |
One electron short means one electron wanted, so each of these atoms has a valency of 1. They complete their outer shells by accepting a single electron, and once they have done so their arrangement matches that of the inert gas next to them.
Diagram
Fig. 6.3-1 — One Electron Short Of A Full Shell Two pairs of shell diagrams. Upper pair: a hydrogen atom (1) with an arrow labelled ‘+ 1 electron’ leading to a hydride ion drawn with two electrons in the K shell and the charge shown as H⁻; note beneath, ‘now matches helium, 2’. Lower pair: a chlorine atom (2, 8, 7) with an arrow labelled ‘+ 1 electron’ leading to a chloride ion (2, 8, 8) marked Cl⁻; note beneath, ‘now matches argon, 2, 8, 8’. Show the gained electron in a contrasting tint in each product.
6.3.2 Formation Of Anions And Electronegative Character
An atom that gains an electron carries one negative charge, and such a particle is called an anion. Since hydrogen and the halogens are each one electron short of a full shell, both show a tendency to form anions.
\[\\mathrm{H} + e^{-} \\rightarrow \\mathrm{H^{-}}\]
\[\\mathrm{Cl} + e^{-} \\rightarrow \\mathrm{Cl^{-}}\]
\[\\mathrm{F} + e^{-} \\rightarrow \\mathrm{F^{-}}\]
An element whose atoms attract and accept electrons in this way is said to be electronegative. Both hydrogen and the halogens are non-metals and both show electronegative character — the exact opposite of the electropositive behaviour you met in the previous section.
6.3.3 Physical State And Atomicity
Fluorine and chlorine are gases at room temperature, and so is hydrogen. In physical state it belongs with the top of Group 17 rather than with the soft solid metals of Group 1.
The number of atoms present in one molecule of an element is called its atomicity. Hydrogen and all the halogens exist as molecules built from two atoms each — \(\\mathrm{H\_2}\), \(\\mathrm{F\_2}\), \(\\mathrm{Cl\_2}\), \(\\mathrm{Br\_2}\) and \(\\mathrm{I\_2}\) — so every one of them has an atomicity of 2.
Taken together, sections 6.2 and 6.3 explain why hydrogen cannot be settled into one group. It has the single valence electron and the positive ion of an alkali metal, and the single electron vacancy, the negative ion, the gaseous state and the diatomic molecule of a halogen.
6.4 How Hydrogen Differs From The Alkali Metals And The Halogens
Two resemblances do not make hydrogen a member of either family. In two important respects it stands apart from both, and these differences are as much a part of its character as the similarities are.
The first difference is in the atom itself. A hydrogen atom has only one shell, the K shell. Every alkali metal and every halogen has two shells or more, so hydrogen is by far the smallest and simplest atom of the three sets.
The second difference lies in the oxides. When an oxide dissolves in water, the solution formed may turn blue litmus red, or red litmus blue, or leave both unchanged. An oxide whose solution turns red litmus blue is a basic oxide; one whose solution turns blue litmus red is an acidic oxide; and one that changes neither is a neutral oxide.
The oxide of hydrogen is water, \(\\mathrm{H\_2O}\), and water is neutral. The oxides of the alkali metals, such as \(\\mathrm{Na\_2O}\) and \(\\mathrm{K\_2O}\), are basic. The oxides of the halogens, such as \(\\mathrm{Cl\_2O}\) and \(\\mathrm{Cl\_2O\_7}\), are acidic. On this test hydrogen sits squarely between the two groups, belonging to neither. Fig. 6.4-1 shows how the three kinds behave with litmus.
Diagram
Fig. 6.4-1 — Three Oxides Tested With Litmus Three test tubes in a rack, each half full of a colourless solution, with a strip of litmus paper dipping into each. Left tube labelled ‘solution of Na₂O’, red litmus strip shown turning blue, caption ‘basic oxide’. Centre tube labelled ‘water, H₂O’, with two strips, one red and one blue, both unchanged, caption ‘neutral oxide’. Right tube labelled ‘solution of Cl₂O’, blue litmus strip shown turning red, caption ‘acidic oxide’.
The two differences are set beside the two families in Table 6.5.
6.5 Discovery Of Hydrogen
Hydrogen was made in the laboratory long before anyone understood what it was. In 1672 Robert Boyle prepared the gas by the action of dilute acid on a metal, but he did not follow the matter up, and the gas was treated as one more curious kind of air.
The credit for the discovery goes instead to Henry Cavendish, in 1766. Cavendish prepared the gas from iron and dilute acids — the same chemistry that later filled Charles’s balloon — and he did what Boyle had not. He studied the gas carefully, showed that it was a distinct substance in its own right, and found that when it burns in air, water is formed.
The gas caught fire so readily that Cavendish called it inflammable air, which simply means air that burns. The name described its most obvious behaviour but said nothing about what the gas was made of.
That came from Antoine Lavoisier. In 1783, working from Cavendish’s finding that the gas burns to give water, Lavoisier renamed it hydrogen — a name built from Greek words meaning water and forming. Hydrogen is, quite literally, the water former, and the name has stayed with it ever since. The three steps are set out in Fig. 6.5-1.
Diagram
Fig. 6.5-1 — Three Steps To A Name A simple horizontal timeline with three marked points. 1672, labelled ‘Robert Boyle prepares the gas from a metal and dilute acid’. 1766, labelled ‘Henry Cavendish studies the gas, shows it burns to form water, and calls it inflammable air’. 1783, labelled ‘Antoine Lavoisier renames it hydrogen, the water former’. Draw the line left to right with equal spacing and a small flask icon at each point.
6.6 Occurrence Of Hydrogen
An element may occur in nature in two ways. When it is found on its own, uncombined with anything else, it is said to be in the free state. When it is found locked into compounds with other elements, it is in the combined state. Hydrogen occurs in both, but not in equal measure.
6.6.1 Hydrogen In The Free State
On the earth, free hydrogen is scarce. It is so light that it drifts away from the atmosphere rather than settling in it, and only traces are found. Table 6.6 gives the figures.
Table 6.6 Hydrogen found in the free state
| Source | Hydrogen present |
|---|---|
| Volcanic gases | 0.025% |
| Earth’s crust | 0.98% |
| Earth’s atmosphere | 0.01% |
Away from the earth the picture is completely different. The atmospheres of the Sun and of the stars contain hydrogen in very large amounts. What is a trace here is the main material there.
6.6.2 Hydrogen In The Combined State
In the combined state hydrogen is everywhere, and Fig. 6.6-1 gathers the main places it is found.
Plant and animal tissues are built from compounds of hydrogen with carbon, oxygen and nitrogen. Sugar, starch, proteins, carbohydrates and fats all contain hydrogen, and so do petroleum products.
Hydrogen is also the characteristic constituent of three whole classes of compound. Every acid contains hydrogen; so does every alkali; and so does every hydrocarbon, which is a compound of hydrogen with carbon alone.
The largest store of all is water. Hydrogen makes up 11.1% of water by weight. You can see where that figure comes from by looking at the formula: in every 18 g of water there are 2 g of hydrogen, and 2 out of 18 is a little over one part in nine.
Diagram
Fig. 6.6-1 — Where Hydrogen Is Found Central circle labelled ‘hydrogen’, with six arrows radiating out to small labelled illustrations arranged around it. Clockwise from the top: the Sun and stars, labelled ‘free state, in large amounts’; a beaker of acid labelled ‘acids’; a bottle labelled ‘alkalis’; a fuel drum labelled ‘hydrocarbons and petroleum products’; a leaf and a human outline labelled ‘plant and animal tissues — sugar, starch, proteins, fats’; and a water droplet labelled ‘water, 11.1% hydrogen by weight’. Mark the first arrow ‘free state’ and the remaining five ‘combined state’.
Because water, acids and alkalis are all cheap and all contain hydrogen, they are the obvious starting materials for making the gas. That is where the next section begins.
6.7 Preparation Of Hydrogen From Water
Water, acids and alkalis all contain hydrogen, and all three are cheap and easily had. They are therefore the raw materials from which hydrogen is prepared, both in the laboratory and on the industrial scale. This section deals with the first of them, water.
To get the hydrogen out, something must take the oxygen away from it. A metal will do this if it is reactive enough — it seizes the oxygen and leaves the hydrogen free. But metals differ enormously in how eager they are, and that is why some need only cold water while others must be made red hot and blasted with steam.
6.7.1 Action Of Cold Water On Metals
The most reactive metals — potassium, sodium and calcium — attack cold water at once. In every case the metal takes the place of one hydrogen atom in each water molecule, so a hydroxide of the metal is left in solution and hydrogen gas is set free.
metal + cold water → metal hydroxide + hydrogen
All three reactions are exothermic, which means heat is given out as they proceed. You will see this in the descriptions that follow: the metal melts in its own heat, and in two of the three cases the escaping gas catches fire.
6.7.1.1 Potassium
Potassium is the most violent of the three. Its density is 0.86 g/cc, which is less than that of water, so a piece dropped into a trough floats on the surface instead of sinking.
The reaction is so strongly exothermic that the metal melts almost immediately — its melting point is only about 63 °C — and forms a silver-grey globule that darts about on the surface of the water.
potassium + water → potassium hydroxide + hydrogen
\[2,\\mathrm{K} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{KOH} + \\mathrm{H\_2}\]
Bubbles of hydrogen are seen rising all round the globule. The gas catches fire and burns with a lilac-coloured flame. The solution left behind is colourless, feels soapy between the fingers, and is alkaline — that is, it turns red litmus blue.
Pure hydrogen burns with a pale blue flame, not a lilac one. The lilac colour comes from traces of potassium vapour carried up in the burning gas, in the same way that a potassium salt colours a bunsen flame.
Potassium will catch fire in air alone, because it reacts even with the water vapour present there. It is kept under kerosene oil, and it is never touched with the hand — a pair of tongs is used to lift it, and only a very small piece is ever taken.
Diagram
Fig. 6.7-1 — Action Of Cold Water On Potassium Sectional view of a shallow trough of water standing on a bench. A small molten globule floats on the surface, labelled ‘potassium’, with short motion lines showing it darting across the water. Draw a flame above the globule and label it ‘hydrogen burning with a lilac flame’. Label the liquid ‘cold water’ and add a few bubbles rising round the globule labelled ‘hydrogen’.
6.7.1.2 Sodium
Sodium behaves in much the same way, but less fiercely. Its density is 0.97 g/cc, so it too floats, and its melting point is 97 °C. The heat of the reaction melts it into a silvery globule which darts about on the surface.
The reaction is less exothermic and less vigorous than that of potassium.
sodium + water → sodium hydroxide + hydrogen
\[2,\\mathrm{Na} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{NaOH} + \\mathrm{H\_2}\]
Bubbles of hydrogen are produced steadily, and the gas burns with a golden-yellow flame — again the colour of sodium vapour rather than of hydrogen itself. The solution formed is colourless, soapy, slightly warm and alkaline.
An amalgam is an alloy of a metal with mercury. Sodium amalgam and potassium amalgam react smoothly with water instead of violently, so hydrogen can be prepared safely from them.
Diagram
Fig. 6.7-2 — Action Of Water On Sodium Beaker of water on a stand, with a small globule labelled ‘sodium’ floating on the surface. An inverted funnel is placed over the globule, its stem leading up into a test tube held mouth-downward and clamped above; label the tube ‘hydrogen collected’. Label the beaker contents ‘water’ and the bubbles under the funnel ‘hydrogen’.
6.7.1.3 Calcium
Calcium is denser than water, so unlike the other two it sinks. Its reaction is less vigorous than that of sodium, and the gas comes off in a steady stream of bubbles rather than in a rush.
calcium + water → calcium hydroxide + hydrogen
\[\\mathrm{Ca} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{Ca(OH)\_2} + \\mathrm{H\_2}\]
Because the calcium sinks, the bubbles rise from the bottom of the trough and are easy to trap. An inverted funnel is placed over the metal and a test tube held mouth-downward over the stem, as in Fig. 6.7-3, and the hydrogen collects in the tube.
Calcium hydroxide is only slightly soluble in water, so most of it stays suspended and the liquid turns milky and turbid. The solution is alkaline: a drop of red litmus added to it turns blue.
Diagram
Fig. 6.7-3 — Action Of Cold Water On Calcium Sectional view of a beaker of water. A grey lump labelled ‘calcium metal’ rests at the bottom, under an inverted glass funnel labelled ‘inverted funnel’. A column of bubbles rises through the funnel stem into a test tube held mouth-downward above it, labelled ‘hydrogen gas’. Label the liquid ‘water’ and show it faintly clouded, with the note ‘solution turns milky’.
The hydroxides of sodium, potassium and calcium all give the same litmus result: red litmus turns blue in each of their solutions. That single test is enough to show that an alkaline hydroxide has been formed.
6.7.2 Action Of Hot Water On Magnesium
Magnesium is less reactive than the three metals above it, and cold water leaves it very nearly alone. With boiling water, however, it reacts slowly, forming the base magnesium hydroxide and setting hydrogen free.
magnesium + boiling water → magnesium hydroxide + hydrogen
\[\\mathrm{Mg} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{Mg(OH)\_2} + \\mathrm{H\_2}\\uparrow\]
Push the temperature higher still, to steam, and the reaction becomes brilliant. Magnesium burns in steam with an intense white light, liberating hydrogen and leaving a white ash of magnesium oxide.
magnesium + steam → magnesium oxide + hydrogen
\[\\mathrm{Mg} + \\mathrm{H\_2O}\\text{(steam)} \\rightarrow \\mathrm{MgO} + \\mathrm{H\_2}\\uparrow\]
Notice that hot water gives the hydroxide while steam gives the oxide. The magnesium oxide formed crumbles apart in the heat instead of forming a tight skin, so fresh metal keeps being exposed to the steam and hydrogen keeps coming off.
Diagram
Fig. 6.7-4 — Action Of Steam On Magnesium Flask of boiling water on a stand over a burner, labelled ‘boiling water’, with a delivery tube carrying steam into a hard-glass tube clamped horizontally. Inside the hard-glass tube show a coil of magnesium labelled ‘magnesium ribbon’, heated from below by a second burner. The far end of the hard-glass tube leads down through a delivery tube into a trough of water, where a gas jar stands inverted collecting gas; label the jar ‘H₂’ and the trough ‘water’.
6.7.3 Action Of Steam On Metals
Below magnesium the metals will not touch water at all, hot or cold. They react only with steam, and only when they themselves are strongly heated. Three of them matter here.
6.7.3.1 Aluminium
Aluminium reacts with steam to liberate hydrogen, forming aluminium oxide.
aluminium + steam → aluminium oxide + hydrogen
\[2,\\mathrm{Al} + 3,\\mathrm{H\_2O} \\xrightarrow{\\ 800^\\circ\\text{C}\\ } \\mathrm{Al\_2O\_3} + 3,\\mathrm{H\_2}\]
The oxide formed does not crumble away. It clings to the surface as a thin unbroken coat, which shuts the steam out from the metal beneath and makes the aluminium inactive. The reaction therefore stops of its own accord. At high enough temperatures the coating breaks up, and aluminium reacts with steam again, liberating hydrogen.
6.7.3.2 Zinc
Zinc is less reactive still. It reacts only when it is heated and steam is passed over it, and the product is white zinc oxide.
zinc + steam → zinc oxide + hydrogen
\[\\mathrm{Zn} + \\mathrm{H\_2O} \\xrightarrow{\\ \\Delta\\ } \\mathrm{ZnO} + \\mathrm{H\_2}\]
Zinc oxide has a useful habit worth remembering: it is yellow while it is hot and white once it has cooled.
6.7.3.3 Red-Hot Iron And A Reversible Reaction
Iron is less reactive than zinc, but red-hot iron does react with steam. The product is triferric tetraoxide, \(\\mathrm{Fe\_3O\_4}\), together with hydrogen.
iron + steam ⇌ triferric tetraoxide + hydrogen
\[3,\\mathrm{Fe} + 4,\\mathrm{H\_2O} \\rightleftharpoons \\mathrm{Fe\_3O\_4} + 4,\\mathrm{H\_2}\]
The double arrow is there for a reason. This change is a reversible reaction — one that runs in both directions, so that the products can react together and give back the original substances. If the hydrogen formed is not removed, it attacks the iron oxide and reduces it to iron again.
Diagram
Fig. 6.7-5 — Action Of Steam On Red-Hot Iron Flask of boiling water on a stand over a burner, labelled ‘boiling water’, with a tube carrying steam into one end of a horizontal silica glass tube labelled ‘silica glass tube’. Inside the silica tube show a plug of steel wool labelled ‘steel wool’, heated strongly from below by a second burner labelled ‘heat’. The other end of the silica tube carries a delivery tube down into a trough of water, with an inverted gas jar collecting the gas; label it ‘hydrogen’.
How Equilibrium Is Reached
Watch the reaction from the moment it starts. At first there is plenty of iron and plenty of steam, and no products at all, so the forward reaction is fast and the backward reaction has nothing to work with. The speed of a reaction depends on how much of the reacting substances is present.
As the iron and steam are used up, the forward reaction slows. As the oxide and hydrogen build up, the backward reaction speeds up. There comes a moment when the two are running at exactly the same speed, and the reaction is then said to have reached chemical equilibrium.
From then on the reaction does not appear to move in either direction. The amounts of reactants and products stop changing, although both changes are still going on underneath. For this reaction that stage is reached at 700 °C.
Comparing The Metals With Steam
The rate at which these metals react with steam is fastest for magnesium, and falls steadily as you go on to aluminium, then zinc, then iron.
With magnesium, aluminium and zinc there is a second effect. The oxide formed sticks to the surface of the metal and blocks the steam from reaching it, so the reaction stops after some time. Table 6.7 collects what each metal needs.
6.7.4 Action Of Steam On Carbon
It is not only metals that will pull the oxygen out of water. Carbon, a non-metal, does the same when the conditions are severe enough.
When steam is passed over red-hot coke, the products are carbon monoxide and hydrogen. The mixture of the two gases is called water gas.
carbon + steam → carbon monoxide + hydrogen
\[\\mathrm{C} + \\mathrm{H\_2O} \\rightarrow \\mathrm{CO} + \\mathrm{H\_2}\]
Water gas is a mixture, not a compound, and to obtain hydrogen from it the carbon monoxide must be taken out. This is done by passing the mixture through ammoniacal cuprous chloride solution, which dissolves the carbon monoxide and leaves hydrogen behind.
6.8 The Activity Series And Its Application
The previous section showed the metals falling into a clear order. Potassium attacked cold water violently, sodium a little less so, calcium more gently still; magnesium needed hot water, aluminium and zinc needed steam, iron needed to be red hot. That order is not an accident, and it is worth setting down properly.
6.8.1 Arrangement Of The Series
When the metals are listed in decreasing order of reactivity, the list is called the activity series — the reactivity series is another name for the same thing.
Potassium, the most reactive metal, stands at the top. Gold, the least reactive, stands at the bottom. Fig. 6.8-1 gives the series as you will use it.
Diagram
Fig. 6.8-1 — The Activity Series Of Metals Vertical ladder of boxes, one symbol to a box, reading from top to bottom: K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au. Show the entry [H] in square brackets and in a contrasting tint, labelled ‘hydrogen, included for comparison’. Down the left-hand side draw a long arrow pointing downwards labelled ‘reactivity decreases’; down the right-hand side a second arrow pointing downwards labelled ‘electropositive character and reducing power decrease’. Mark the top box ‘most reactive’ and the bottom box ‘least reactive’.
Hydrogen is a non-metal, yet it has been slipped into the list between lead and copper. It is included because it can form a positive ion, just as the metals do, and its place in the series is fixed by how readily it does so. Once hydrogen has a position, the series can be used to predict whether a given metal will push hydrogen out of a compound or not.
The ability of metals to reduce water to hydrogen falls away as you go down:
\[\\mathrm{K} > \\mathrm{Na} > \\mathrm{Ca} > \\mathrm{Mg} > \\mathrm{Al} > \\mathrm{Zn} \\ \\ldots\]
Lead and the metals below it in the series show no reaction with water at all, even when the metal is hot and steam is used.
6.8.2 Salient Features Of The Activity Series
The series is more than a ranking. Once you know where a metal stands, a great many of its properties follow.
Electropositive Character And Reducing Power
Electropositive character decreases down the series. So does reducing power, which means potassium is the strongest reducing agent in the list and gold the weakest.
Both of these say the same thing in different words. A metal high in the series gives up its valence electrons easily; a metal low down holds on to them. Since losing electrons is what makes a metal electropositive, and since a reducing agent works by handing electrons over, the two properties must run together.
The tendency of metals to lose valence electrons also decreases down the series. Potassium is therefore the metal most readily oxidised, and gold the least.
Displacement Of Hydrogen
Metals above hydrogen in the series displace hydrogen from water and from dilute acids. Metals below hydrogen do not. This single rule explains a good deal.
Sodium stands well above hydrogen, and you have already seen it drive hydrogen out of cold water. Copper stands below hydrogen, and nothing whatever happens:
\[\\mathrm{Cu} + \\mathrm{H\_2O} \\rightarrow \\text{no reaction}\]
\[\\mathrm{Cu} + \\mathrm{HCl}\\text{ (dil.)} \\rightarrow \\text{no reaction}\]
Displacement Of One Metal By Another
A metal higher in the series displaces a metal below it from a solution of the latter’s salt. Drop a strip of zinc into blue copper(II) sulphate solution and the zinc goes into solution while copper is deposited on it as a reddish-brown coating. The blue colour fades as the copper leaves the solution.
zinc + copper(II) sulphate → zinc sulphate + copper
\[\\mathrm{Zn} + \\mathrm{CuSO\_4}\\text{(aq)} \\rightarrow \\mathrm{ZnSO\_4}\\text{(aq)} + \\mathrm{Cu}\]
The larger the gap between the two metals in the series, the more rapidly the displacement occurs.
Diagram
Fig. 6.8-2 — Zinc Displacing Copper From Copper(II) Sulphate Solution Two beakers side by side. Left beaker, labelled ‘at the start’: deep blue solution labelled ‘copper(II) sulphate solution’ with a clean grey strip labelled ‘zinc’ dipping into it. Right beaker, labelled ‘after some time’: the solution drawn much paler, labelled ‘colour fades’, and the strip now coated with a dark deposit labelled ‘reddish-brown copper deposited on the zinc’. Add an arrow between the two beakers.
Reduction Of Metal Oxides
The oxides of the metals at the top of the series — potassium, sodium, calcium, magnesium and aluminium — cannot be reduced by the common reducing agents such as hydrogen, carbon monoxide or carbon. These metals hold their oxygen too firmly. Their oxides can be reduced only by electrolysis.
Metal oxides below aluminium in the series can be reduced simply by heating them with a reducing agent.
\[\\mathrm{ZnO} + \\mathrm{C} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Zn} + \\mathrm{CO}\]
\[\\mathrm{PbO} + \\mathrm{C} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Pb} + \\mathrm{CO}\]
\[\\mathrm{PbO} + \\mathrm{CO} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Pb} + \\mathrm{CO\_2}\]
\[\\mathrm{PbO} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Pb} + \\mathrm{H\_2O}\]
Go lower still and even a reducing agent becomes unnecessary. The oxides and nitrates of the least reactive metals — mercury, silver and gold — simply fall apart on strong heating and give the metal itself.
\[2,\\mathrm{HgO} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{Hg} + \\mathrm{O\_2}\]
\[2,\\mathrm{AgNO\_3} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{Ag} + 2,\\mathrm{NO\_2} + \\mathrm{O\_2}\]
Rusting And Occurrence In Nature
Metals below copper in the series do not rust easily. Their reluctance to combine with oxygen and moisture is exactly what makes silver and gold useful for ornaments that must last.
The same reluctance decides how a metal is found in the earth. Since the series lists metals in decreasing order of electropositivity, those at the top are expected to be combined with the most electronegative elements. And so they are: sodium and potassium occur mainly as their chlorides, aluminium as its oxide, while zinc, lead, copper, mercury and silver occur as their sulphides.
6.8.3 Displacement Of Hydrogen From Dilute Acids
Every acid contains hydrogen. If a metal is reactive enough to take the place of that hydrogen in the acid, the hydrogen is set free as a gas and a salt of the metal is left in solution.
metal + dilute acid → salt + hydrogen
Hydrogen is displaced in this way from dilute hydrochloric acid and from dilute sulphuric acid, but only by metals more reactive than hydrogen itself. How violently it happens depends on how far above hydrogen the metal stands, as Table 6.8 shows.
Magnesium, aluminium, zinc and iron are the four metals of the middle range, and each gives hydrogen together with its own salt. With hydrochloric acid the salt is a chloride; with sulphuric acid it is a sulphate.
magnesium + dilute hydrochloric acid → magnesium chloride + hydrogen
\[\\mathrm{Mg} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{MgCl\_2} + \\mathrm{H\_2}\]
\[\\mathrm{Mg} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{MgSO\_4} + \\mathrm{H\_2}\]
aluminium + dilute hydrochloric acid → aluminium chloride + hydrogen
\[2,\\mathrm{Al} + 6,\\mathrm{HCl} \\rightarrow 2,\\mathrm{AlCl\_3} + 3,\\mathrm{H\_2}\]
\[2,\\mathrm{Al} + 3,\\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{Al\_2(SO\_4)\_3} + 3,\\mathrm{H\_2}\]
zinc + dilute hydrochloric acid → zinc chloride + hydrogen
\[\\mathrm{Zn} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2}\]
\[\\mathrm{Zn} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{ZnSO\_4} + \\mathrm{H\_2}\]
iron + dilute hydrochloric acid → iron(II) chloride + hydrogen
\[\\mathrm{Fe} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{FeCl\_2} + \\mathrm{H\_2}\]
\[\\mathrm{Fe} + \\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{FeSO\_4} + \\mathrm{H\_2}\]
Aluminium takes three atoms of chlorine to its one, and three sulphate groups to its two atoms, because its valency is 3 while that of the other three metals here is 2. The coefficients in those equations are simply the consequence.
6.8.4 Displacement Of Hydrogen From Alkalis
Three metals — zinc, lead and aluminium — do something none of the others can. They react with acids, in the ordinary way you have just seen, and they also react with hot concentrated alkalis, again giving hydrogen and again forming a soluble salt. This double ability is what is meant by saying these three have a unique nature.
metal + hot concentrated alkali → soluble salt + hydrogen
Zinc gives zincates:
\[\\mathrm{Zn} + 2,\\mathrm{NaOH} \\rightarrow \\mathrm{Na\_2ZnO\_2} + \\mathrm{H\_2}\]
\[\\mathrm{Zn} + 2,\\mathrm{KOH} \\rightarrow \\mathrm{K\_2ZnO\_2} + \\mathrm{H\_2}\]
The salts are sodium zincate and potassium zincate. Lead gives the corresponding plumbites:
\[\\mathrm{Pb} + 2,\\mathrm{NaOH} \\rightarrow \\mathrm{Na\_2PbO\_2} + \\mathrm{H\_2}\]
\[\\mathrm{Pb} + 2,\\mathrm{KOH} \\rightarrow \\mathrm{K\_2PbO\_2} + \\mathrm{H\_2}\]
Aluminium gives aluminates, and here the product depends on the conditions. With fused sodium hydroxide — solid alkali melted by heat, with no water present — sodium aluminate is formed:
\[2,\\mathrm{Al} + 6,\\mathrm{NaOH}\\text{ (fused)} \\rightarrow 2,\\mathrm{Na\_3AlO\_3} + 3,\\mathrm{H\_2}\]
With alkali solution, where water is present, the meta aluminate is formed instead:
\[2,\\mathrm{Al} + 2,\\mathrm{NaOH} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{NaAlO\_2} + 3,\\mathrm{H\_2}\]
\[2,\\mathrm{Al} + 2,\\mathrm{KOH} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{KAlO\_2} + 3,\\mathrm{H\_2}\]
The last two salts are sodium meta aluminate and potassium meta aluminate.
Why These Three Are Different
The reason lies in their oxides and hydroxides. The oxides and hydroxides of zinc, lead and aluminium are amphoteric: they react with acids and with bases alike, and in both cases they give a salt and water.
\[\\mathrm{ZnO} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2O}\]
\[\\mathrm{ZnO} + 2,\\mathrm{NaOH} \\rightarrow \\mathrm{Na\_2ZnO\_2} + \\mathrm{H\_2O}\]
In the first reaction zinc oxide behaves as a base would, neutralising an acid. In the second it behaves as an acidic oxide would, neutralising an alkali. A substance that can face both ways is amphoteric, and it is this two-faced character that lets the metals themselves give hydrogen with acid and with alkali.
6.9 Laboratory Preparation Of Hydrogen
Everything in the last two sections now has to be put to work. To prepare hydrogen on the bench you must choose a metal, choose an acid, choose a way of collecting the gas, and then clean it up — and every one of those choices has a reason behind it.
6.9.1 Reactants, Apparatus And Reaction
The reactants are granulated zinc, together with dilute hydrochloric acid or dilute sulphuric acid.
The apparatus is shown in Fig. 6.9-1. Some pieces of granulated zinc are placed in a flat-bottomed flask, which is closed with an airtight cork carrying two holes. Through one hole passes a thistle funnel with a long stem and a stopper; through the other passes a long delivery tube. The acid is poured in through the funnel.
The reaction is the one you met among the metals and dilute acids in section 6.8. Either acid may be used.
zinc + dilute hydrochloric acid → zinc chloride + hydrogen
\[\\mathrm{Zn} + 2,\\mathrm{HCl}\\text{(dil.)} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2}\\uparrow\]
zinc + dilute sulphuric acid → zinc sulphate + hydrogen
\[\\mathrm{Zn} + \\mathrm{H\_2SO\_4}\\text{(dil.)} \\rightarrow \\mathrm{ZnSO\_4} + \\mathrm{H\_2}\\uparrow\]
The reaction begins gradually. Bubbles form on the surface of the metal and the liquid begins to fizz — this fizzing is called effervescence — and the gas passes along the delivery tube.
Diagram
Fig. 6.9-1 — Laboratory Preparation Of Hydrogen Flat-bottomed flask clamped on a stand, containing grey lumps labelled ‘zinc granules’ under a layer of liquid. A two-holed rubber cork closes the flask, labelled ‘airtight cork’. Through one hole a thistle funnel passes, its long stem reaching below the surface of the liquid; label it ‘thistle funnel (or dropping funnel)’ and mark the liquid being poured in as ‘dilute sulphuric acid or dilute hydrochloric acid’. Through the other hole a long delivery tube leads out and bends down into a trough of water; label the tube ‘hydrogen’. In the trough, a gas jar stands inverted on a beehive shelf, filling with gas and pushing the water down; label the trough ‘water’ and the jar ‘hydrogen collected over water’.
Why Zinc, And Not Another Metal
Zinc is the preferred metal for this preparation. Every other candidate fails on some practical ground.
Sodium and potassium react with acid violently, and are far too dangerous to use.
Calcium and magnesium are expensive.
Aluminium has so great an affinity for oxygen that it carries a protective coating of \(\\mathrm{Al\_2O\_3}\), and because of that coating it does not give hydrogen with the acid.
Iron has to be heated, and the hydrogen then obtained carries impurities such as hydrogen sulphide and sulphur dioxide.
Lead does react with dilute sulphuric or hydrochloric acid, but the lead sulphate or lead chloride formed is insoluble. It settles as a coating over the metal, and further reaction is prevented.
Copper, mercury and the other metals below hydrogen cannot be used at all, since only a metal more reactive than hydrogen can displace it from an acid.
Zinc escapes all six objections. It is cheap, it lies above hydrogen, it needs no heating, its salts are soluble, and it reacts briskly without reacting dangerously.
6.9.2 Why Granulated Zinc Is Chosen
Granulated zinc is preferred to pure zinc, which sounds like the wrong way round until you see the reason. Granulated zinc contains traces of copper as an impurity, and that copper has a catalysing effect on the reaction, speeding up the rate at which the gas comes off. Pure zinc reacts disappointingly slowly.
6.9.3 Collection Of The Gas
The gas is not collected until all the air has been driven out of the apparatus, since a mixture of hydrogen and air is dangerous. Once the air has been expelled, hydrogen is collected over water by downward displacement of water — the jar stands full of water, mouth downward, and the gas pushes the water out as it enters.
Two properties make this the right method.
Hydrogen is virtually insoluble in water. Only about 20 mL of the gas dissolves in one litre of water under normal conditions, so almost none of the sample is lost.
Hydrogen forms an explosive mixture with air. It therefore cannot be collected by downward displacement of air, even though it is lighter than air and that method would otherwise work.
6.9.4 Impurities Present In The Gas
The hydrogen collected is not pure. It carries with it:
hydrogen sulphide, \(\\mathrm{H\_2S}\)
sulphur dioxide, \(\\mathrm{SO\_2}\)
oxides of nitrogen
phosphine, \(\\mathrm{PH\_3}\)
arsine, \(\\mathrm{AsH\_3}\)
carbon dioxide
water vapour
Most of these come from small amounts of sulphur, phosphorus and arsenic present in the zinc or the acid.
6.9.5 Removal Of The Impurities
The impurities are taken out by passing the gas through a series of solutions, each of which traps a particular set. The order of passage is given in FC 6.1 and the apparatus in Fig. 6.9-2.
Flowchart
FC 6.1 — Purification Of Hydrogen Impure hydrogen → silver nitrate solution (removes arsine and phosphine) → lead nitrate solution (removes hydrogen sulphide) → caustic potash solution (removes sulphur dioxide, carbon dioxide and oxides of nitrogen) → anhydrous calcium chloride (removes water vapour) → pure, dry hydrogen
- Silver nitrate solution, which removes arsine and phosphine.
\[\\mathrm{AsH\_3} + 6,\\mathrm{AgNO\_3} \\rightarrow \\mathrm{Ag\_3As!\\cdot!3AgNO\_3} + 3,\\mathrm{HNO\_3}\]
\[\\mathrm{PH\_3} + 6,\\mathrm{AgNO\_3} \\rightarrow \\mathrm{Ag\_3P!\\cdot!3AgNO\_3} + 3,\\mathrm{HNO\_3}\]
- Lead nitrate solution, which removes hydrogen sulphide. The black lead sulphide formed is insoluble and drops out of solution.
\[\\mathrm{Pb(NO\_3)\_2} + \\mathrm{H\_2S} \\rightarrow \\mathrm{PbS} + 2,\\mathrm{HNO\_3}\]
- Caustic potash solution, which removes sulphur dioxide, carbon dioxide and the oxides of nitrogen. All three are acidic gases, and an alkali absorbs them.
\[\\mathrm{SO\_2} + 2,\\mathrm{KOH} \\rightarrow \\mathrm{K\_2SO\_3} + \\mathrm{H\_2O}\]
\[\\mathrm{CO\_2} + 2,\\mathrm{KOH} \\rightarrow \\mathrm{K\_2CO\_3} + \\mathrm{H\_2O}\]
\[2,\\mathrm{NO\_2} + 2,\\mathrm{KOH} \\rightarrow \\mathrm{KNO\_2} + \\mathrm{KNO\_3} + \\mathrm{H\_2O}\]
Diagram
Fig. 6.9-2 — Purification Of Hydrogen Four vessels in a row on a bench, connected in series by glass tubing that dips below the liquid in each. Impure gas enters at the left, labelled ‘impure H₂’. First wash bottle labelled ‘AgNO₃ solution’, with the outgoing tube marked ‘H₂S, CO₂, SO₂, H₂O still present’. Second wash bottle labelled ‘Pb(NO₃)₂ solution’, outgoing tube marked ‘CO₂, SO₂, H₂O still present’. Third wash bottle labelled ‘KOH solution’, outgoing tube marked ‘H₂O still present’. Fourth vessel a U-tube packed with granules, labelled ‘anhydrous CaCl₂’. The final tube is labelled ‘pure, dry hydrogen’.
6.9.6 Drying The Gas
Water vapour is the last impurity, and it is removed by a drying agent — a substance that absorbs moisture without reacting with the gas itself. Fused calcium chloride, sticks of caustic potash, or phosphorus pentoxide may all be used.
Concentrated sulphuric acid is an excellent drying agent, and you will meet it as one elsewhere, but it must never be used here. It reacts with hydrogen, which defeats the whole purpose of the operation.
\[\\mathrm{H\_2SO\_4} + \\mathrm{H\_2} \\rightarrow 2,\\mathrm{H\_2O} + \\mathrm{SO\_2}\]
Once purified and dried, the gas is collected over mercury rather than over water, since mercury has no reaction with it and adds no moisture back.
6.9.7 Precautions During The Preparation
A mixture of hydrogen and air explodes violently if it is brought near a flame. The precautions below all follow from that one fact, and none of them is optional.
The apparatus must be airtight, with no leakage of gas at any joint.
No flame may be burning anywhere near the apparatus.
The gas must be collected only after all the air in the apparatus has escaped. To check, collect a little of the gas in a test tube and take it to a flame away from the bench. If it burns quietly, no air is left in the flask.
The end of the thistle funnel must dip below the surface of the acid. If it does not, the gas escapes up the funnel instead of going down the delivery tube.
6.9.8 Acids That Are Not Used
Nitric acid is not used in this preparation, even in its dilute form. Nitric acid is a powerful oxidising agent: it decomposes to release oxygen, and that oxygen oxidises the hydrogen to water as fast as it is formed. Instead of collecting a gas you are left with nothing but a solution.
\[3,\\mathrm{Zn} + 8,\\mathrm{HNO\_3} \\rightarrow 3,\\mathrm{Zn(NO\_3)\_2} + 4,\\mathrm{H\_2O} + 2,\\mathrm{NO}\\uparrow\]
There is one narrow exception. Magnesium and manganese are the only metals that react with very dilute nitric acid — about 1% acid — to liberate hydrogen. At that dilution the oxidising action of the acid is much reduced.
\[\\mathrm{Mg} + 2,\\mathrm{HNO\_3} \\rightarrow \\mathrm{Mg(NO\_3)\_2} + \\mathrm{H\_2}\\uparrow\]
Concentrated sulphuric acid is ruled out for the same kind of reason. It is a strong oxidiser, and with zinc it gives sulphur dioxide instead of hydrogen.
\[\\mathrm{Zn} + 2,\\mathrm{H\_2SO\_4} \\rightarrow \\mathrm{ZnSO\_4} + \\mathrm{SO\_2} + 2,\\mathrm{H\_2O}\]
Dilute sulphuric acid and dilute hydrochloric acid are therefore the only two acids used.
6.9.9 Test For Hydrogen
Hydrogen is combustible — it burns — but it does not support combustion. A burning splint held at the mouth of a jar of hydrogen sets the gas alight, but pushed down into the jar the splint goes out.
Burnt quietly in air or oxygen, hydrogen gives a pale blue flame and forms water, as you saw in section 6.2. But when it has been premixed with air or with oxygen, the whole sample burns at once and it explodes with a characteristic pop sound. That pop is the standard test for hydrogen.
6.10 Manufacture Of Hydrogen
A flask of zinc and acid will fill a gas jar. Industry needs hydrogen by the tonne — for making ammonia, for hardening oils, for refining fuels — and at that scale zinc is far too costly. The raw materials for manufacture are the cheapest things available: coke, steam and electricity.
6.10.1 The Bosch Process
The Bosch process is used to prepare hydrogen on a large scale, and it works in two stages.
Stage One — Making Water Gas
Steam is passed over hot coke at 1000 °C in furnaces of a special design, called converters. The product is the mixture of carbon monoxide and hydrogen that you met in section 6.7 under the name water gas.
\[\\mathrm{C} + \\mathrm{H\_2O} \\xrightarrow{\\ 1000^\\circ\\text{C}\\ } \\mathrm{CO} + \\mathrm{H\_2}\]
This reaction is endothermic — it absorbs heat rather than giving it out — which is why the coke must be kept hot from outside.
Stage Two — Converting The Carbon Monoxide
Only half the mixture is hydrogen, and the carbon monoxide in it is a waste of good raw material. So the water gas is mixed with excess steam and passed over heated ferric oxide, which acts as the catalyst, and chromic oxide, \(\\mathrm{Cr\_2O\_3}\), which acts as the promoter. The temperature is 450 °C.
\[\\mathrm{CO} + \\mathrm{H\_2} + \\mathrm{H\_2O} \\xrightarrow{\\ \\mathrm{Fe\_2O\_3},\\ 450^\\circ\\text{C}\\ } \\mathrm{CO\_2} + 2,\\mathrm{H\_2}\]
This reaction is exothermic. Notice what it has achieved: the carbon monoxide has taken oxygen from the steam and become carbon dioxide, and in doing so it has released a second molecule of hydrogen. The mixture leaving the converter is carbon dioxide and hydrogen.
6.10.2 Separation Of Carbon Dioxide And Carbon Monoxide
Two gases must now be got rid of. FC 6.2 sets out the whole process from coke to hydrogen.
Flowchart
FC 6.2 — Manufacture Of Hydrogen By The Bosch Process Coke heated to 1000 °C in a converter + steam → water gas (CO + H₂) → mixed with excess steam → passed over Fe₂O₃ catalyst with Cr₂O₃ promoter at 450 °C → mixture of CO₂ and H₂ → CO₂ removed, either by cold water under 30 atm pressure or by caustic potash solution → any uncombined CO removed by ammoniacal cuprous chloride solution → hydrogen
Removing The Carbon Dioxide
Either of two methods may be used.
The mixture of carbon dioxide and hydrogen is forced through cold water under a pressure of 30 atmospheres. Carbon dioxide is much the more soluble of the two, so it dissolves and hydrogen is left behind.
The mixture is passed through caustic potash solution, which absorbs the carbon dioxide by reacting with it — the same absorption you saw in the purification train in section 6.9 — and leaves hydrogen.
Removing The Carbon Monoxide
A little carbon monoxide always escapes conversion in stage two. The gas mixture is therefore passed through ammoniacal cuprous chloride solution, which dissolves any uncombined carbon monoxide.
\[\\mathrm{CuCl} + \\mathrm{CO} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{CuCl}!\\cdot!\\mathrm{CO}!\\cdot!2\\mathrm{H\_2O}\]
Hydrogen gas is left behind, and the manufacture is complete.
6.10.3 Manufacture By Electrolysis Of Acidulated Water
Hydrogen is also obtained commercially by the electrolysis of acidulated water — the decomposition of water by passing an electric current through it.
Pure water is a poor conductor of electricity, so a very small amount of a less volatile acid, such as sulphuric acid, is added to make it conduct. Water treated this way is described as acidulated. The acid is not used up; it only makes conduction possible.
On passing the current, water dissociates into ions.
\[\\mathrm{H\_2O} \\rightleftharpoons \\mathrm{H^{+}} + \\mathrm{OH^{-}}\]
The cathode is the negatively charged electrode and the anode is the positively charged one. Opposite charges attract, so each ion travels to the electrode of opposite sign.
Diagram
Fig. 6.10-1 — Electrolysis Of Acidulated Water Glass vessel of liquid labelled ‘acidulated water’, with a funnel-shaped reservoir at the top for topping up. Two inverted graduated tubes, A and B, stand in the liquid, each over one electrode. The left electrode is connected to the negative terminal of a battery and labelled ‘cathode’; the tube above it is labelled ‘hydrogen’ and drawn with twice the volume of collected gas. The right electrode is connected to the positive terminal and labelled ‘anode’; the tube above it is labelled ‘oxygen’. Show the battery with its + and − terminals clearly marked.
At the cathode, the positively charged hydrogen ions collect. Each takes an electron and becomes an atom, and the atoms pair up into molecules.
\[\\mathrm{H^{+}} + e^{-} \\rightarrow \\mathrm{H}\]
\[\\mathrm{H} + \\mathrm{H} \\rightarrow \\mathrm{H\_2}\]
Hydrogen gas is therefore evolved at the cathode.
At the anode, the negatively charged hydroxide ions give up their electrons, and the resulting groups break down to water and oxygen.
\[\\mathrm{OH^{-}} - e^{-} \\rightarrow \\mathrm{OH}\]
\[\\mathrm{OH} + \\mathrm{OH} \\rightarrow \\mathrm{H\_2O} + \\mathrm{O}\]
\[\\mathrm{O} + \\mathrm{O} \\rightarrow \\mathrm{O\_2}\]
Taken together, those three steps amount to one balanced change at the anode, and it is worth writing it out in that form as well, since the electrons then account for exactly:
\[4,\\mathrm{OH^{-}} - 4e^{-} \\rightarrow 2,\\mathrm{H\_2O} + \\mathrm{O\_2}\]
Oxygen is therefore evolved at the anode. Adding these steps together gives the overall change:
\[2,\\mathrm{H\_2O} \\xrightarrow{\\ \\text{electric current}\\ } 2,\\mathrm{H\_2} + \\mathrm{O\_2}\]
The main advantage of this method is that oxygen is produced at the anode at the same time, and both gases come off pure. Look again at Fig. 6.10-1 and you will see one more thing: exactly twice the volume of gas collects over the cathode as over the anode. Water is being taken apart into two volumes of hydrogen for every one of oxygen, which is direct proof that water is a compound of the two.
6.11 Properties Of Hydrogen
You have prepared the gas, purified it and dried it. What is it actually like? The properties fall into two sets — those you can observe without changing the substance, and those that show up only when it reacts.
6.11.1 Physical Properties
Colour: colourless.
Odour: odourless.
Taste: tasteless.
Physiological nature: non-toxic.
Density: \(0.0899\\ \\text{g L}^{-1}\). It is the lightest substance known, and it is 14.4 times less dense than air.
Solubility: very slightly soluble in water — the one litre of water dissolving about 20 mL of gas that decided how the gas is collected.
Liquefaction: liquefies below \(-240\\ ^\\circ\\text{C}\) under a pressure of about 20 atmospheres.
Showing That Hydrogen Is Light
The density figure can be demonstrated rather than merely quoted. Take some soap solution in a beaker and allow hydrogen to pass through it. Soap bubbles filled with hydrogen form, come out of the solution, and rise up through the air instead of sinking — which ordinary bubbles blown with breath do not do. The behaviour shown in Fig. 6.11-1 proves that hydrogen is lighter than air.
Diagram
Fig. 6.11-1 — Hydrogen-Filled Soap Bubbles Rise In The Air Flask on a bench containing grey lumps labelled ‘zinc granules’ with liquid poured in through a funnel labelled ‘dilute hydrochloric acid’. A delivery tube leads from the flask and dips into a beaker of liquid labelled ‘soap solution’. Show a stream of bubbles forming at the surface of the soap solution and a group of them drifting upwards above the beaker; label them ‘hydrogen filled soap bubbles’ and add an upward arrow beside them.
6.11.2 Occlusion Of Hydrogen By Metals
Some metals soak hydrogen up. Palladium, platinum and nickel all take the gas into themselves at room temperature, holding it on and within their surfaces. One volume of finely divided palladium takes up as much as 900 volumes of hydrogen at room temperature.
This taking up of a gas by a solid is called occlusion, or adsorption.
6.11.3 Chemical Properties
Hydrogen is a reactive element. Under the right conditions it reacts with non-metals, with metals, with metallic oxides and even with organic compounds. One thing it will not do is support combustion.
6.11.3.1 Reaction With Non-Metals
Hydrogen combines with oxygen, chlorine, nitrogen and sulphur, but each reaction has its own conditions, set out in Table 6.9.
Table 6.9 Reaction of hydrogen with non-metals
| Reactants | Equation | Conditions and observations |
|---|---|---|
| Hydrogen and oxygen | \(2,\\mathrm{H\_2} + \\mathrm{O\_2} \\rightarrow 2,\\mathrm{H\_2O}\) | In oxygen, in the presence of an electric spark, hydrogen burns with a pop sound. Pure hydrogen burns quietly in air with a pale blue flame, forming water (Fig. 6.11-2). |
| Hydrogen and chlorine | \(\\mathrm{H\_2} + \\mathrm{Cl\_2} \\rightarrow 2,\\mathrm{HCl}\) | Equal volumes react slowly in diffused sunlight but explosively in direct sunlight, since a large amount of energy is released. Hydrogen continues to burn in chlorine. The product is hydrogen chloride. |
| Hydrogen and nitrogen (Haber process) | \(\\mathrm{N\_2} + 3,\\mathrm{H\_2} \\rightleftharpoons 2,\\mathrm{NH\_3}\) | Three volumes of hydrogen to one volume of nitrogen, at 450–500 °C and 200–900 atm, over finely divided iron as catalyst with molybdenum as promoter. The product is ammonia. |
| Hydrogen and sulphur | \(\\mathrm{H\_2} + \\mathrm{S} \\rightarrow \\mathrm{H\_2S}\) | Hydrogen is passed through molten sulphur. The product is hydrogen sulphide. |
Diagram
Fig. 6.11-2 — Hydrogen Burns In Air To Form Water Delivery tube entering at the left labelled ‘moist pure H₂’, leading into a U-tube packed with granules labelled ‘anhydrous CaCl₂’. From the U-tube the dried gas passes up a narrow jet, where it is shown alight; label the flame ‘burning H₂ with blue flame’. Over the jet an inverted boiling tube is clamped, labelled ‘boiling tube’, with droplets condensing inside it and running down into a small vessel below labelled ‘water’.
6.11.3.2 Reaction With Metals
Dry hydrogen passed over heated potassium, sodium or calcium reacts with them to give their hydrides — compounds of a metal with hydrogen, in which hydrogen is present as the negative ion \(\\mathrm{H^{-}}\) that you met in section 6.3.
Table 6.10 Reaction of hydrogen with metals
| Metal | Equation | Product |
|---|---|---|
| Potassium | \(2,\\mathrm{K} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{KH}\) | Potassium hydride |
| Sodium | \(2,\\mathrm{Na} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{NaH}\) | Sodium hydride |
| Calcium | \(\\mathrm{Ca} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } \\mathrm{CaH\_2}\) | Calcium hydride |
The hydrides so formed react further with water, and in doing so they give hydrogen back again.
calcium hydride + water → calcium hydroxide + hydrogen
\[\\mathrm{CaH\_2} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{Ca(OH)\_2} + 2,\\mathrm{H\_2}\]
6.11.3.3 Reaction With Metallic Oxides
Hydrogen gas is a strong reducing agent. It removes oxygen from strongly heated oxides of the less active metals, leaving the metal behind, and is itself oxidised to water in the process.
The change is easy to follow with the eye, because in every case the coloured oxide gives way to the colour of the metal. Table 6.11 gives the four standard cases.
Table 6.11 Reduction of metallic oxides by hydrogen
| Equation | Colour change as the oxide becomes the metal |
|---|---|
| \(\\mathrm{Fe\_2O\_3} + 3,\\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } 2,\\mathrm{Fe} + 3,\\mathrm{H\_2O}\) | Brown to grey |
| \(\\mathrm{Fe\_3O\_4} + 4,\\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } 3,\\mathrm{Fe} + 4,\\mathrm{H\_2O}\) | Brown to grey |
| \(\\mathrm{PbO} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Pb} + \\mathrm{H\_2O}\) | Yellow to silvery grey |
| \(\\mathrm{CuO} + \\mathrm{H\_2} \\xrightarrow{\\ \\Delta\\ } \\mathrm{Cu} + \\mathrm{H\_2O}\) | Black to reddish brown |
The last of these is the one usually done on the bench, and Fig. 6.11-3 shows the arrangement. Pure dry hydrogen is passed over heated copper(II) oxide in a tube; the excess hydrogen leaving the far end is burnt off at a jet, so that none escapes into the room.
Diagram
Fig. 6.11-3 — Reduction Of Copper Oxide To Copper Metal Horizontal hard-glass tube clamped on stands, with a boat of black powder inside labelled ‘CuO’, heated from below by a bunsen burner. Gas enters from the left through a tube labelled ‘pure and dry H₂’. The right-hand end of the tube leads to a jet where the surplus gas is alight; label the flame ‘bunsen flame of unreacted H₂’ and the gas stream ‘unreacted H₂’. Below the exit, show a small vessel collecting condensed droplets, labelled ‘water’.
Not every oxide can be treated this way. The oxides of the active metals — sodium, potassium, aluminium, calcium and magnesium — have too great an affinity for oxygen, and hydrogen cannot take it from them, as you saw when the activity series was set out in section 6.8.
6.11.3.4 Reaction With Unsaturated Organic Compounds
In the presence of a catalyst such as nickel, hydrogen combines directly with organic compounds that contain a double or a triple bond between two carbon atoms. The process is called hydrogenation.
ethylene + hydrogen → ethane
\[\\mathrm{H\_2C{=}CH\_2} + \\mathrm{H\_2} \\rightarrow \\mathrm{C\_2H\_6}\]
acetylene + hydrogen → ethane
\[\\mathrm{HC{\\equiv}CH} + 2,\\mathrm{H\_2} \\rightarrow \\mathrm{C\_2H\_6}\]
Look at the second equation and you will see why two molecules of hydrogen are needed. A triple bond has two bonds to break open where a double bond has only one, so acetylene takes up twice as much hydrogen as ethylene before it reaches the same saturated product.
6.12 Uses Of Hydrogen
Almost every use of hydrogen follows from one of two facts about it: it releases a great deal of heat when it burns, and it is a strong reducing agent. Keep those two in mind and the list below almost writes itself.
Hydrogen As A Fuel
Hydrogen has a high heat of combustion, which means a given mass of it gives out an unusually large amount of heat when burnt. It is used as a fuel in three forms:
coal gas;
water gas, the mixture of carbon monoxide and hydrogen you met in the Bosch process;
liquid hydrogen, which is non-polluting and easy to store.
Hydrogen Flames For Cutting And Welding
Burn hydrogen in pure oxygen rather than in air and the flame becomes far hotter than an ordinary one. A mixture of the two gases is burnt in a specially designed apparatus called an oxy-hydrogen torch, which produces temperatures touching 2500 °C.
That flame is used for cutting and welding metals, for melting platinum and quartz, and for fusing alumina to produce the synthetic rubies and sapphires used as jewels in watches.
Hotter still is the atomic hydrogen torch, which reaches about 2800 °C. It is used for welding alloys containing metals such as tungsten, manganese and chromium — metals with melting points so high that an ordinary flame cannot touch them.
Diagram
Fig. 6.12-1 — The Oxy-Hydrogen Torch Two gas cylinders standing side by side, the left labelled ‘hydrogen’ and the right labelled ‘oxygen’, each with a regulator valve and a flexible hose. The two hoses run to a hand-held torch; label the point where they join inside it ‘mixing chamber’ and the outlet ‘nozzle’. Show a short, sharply pointed flame at the nozzle directed at a metal plate, labelled ‘flame, temperature up to 2500 °C’ and ‘metal being cut’. Add small arrows on each hose showing the direction of gas flow.
Hydrogen is also used in self-lighting gas jets and in automatic lighters.
Hydrogen As A Raw Material
Great quantities of hydrogen go into the manufacture of compounds, above all ammonia by the Haber process, whose conditions you met in section 6.11.
Hydrogen is used for the hydrogenation of vegetable oil. Liquid vegetable fats such as groundnut oil and coconut oil are treated with hydrogen in the presence of finely divided nickel, platinum or palladium as catalyst, at high pressure and a temperature of about 200 °C. The product is the solid fat sold as vanaspati ghee. The oils are liquid because their molecules contain double bonds; adding hydrogen across those bonds is what turns the liquid into a solid.
Hydrogen also serves in producing artificial petrol from coal. Hydrogen is passed under high pressure over powdered coal in the presence of a catalyst, and the product resembles petroleum but contains a higher proportion of hydrogen.
Hydrogen In The Extraction Of Metals
Because it is a reducing agent, hydrogen is used to extract certain metals from their oxides. Heated oxides of the less active metals — iron and copper among them — are reduced to the metal, as the equations and colour changes in Table 6.11 showed. The oxides of the active metals resist it, and those metals must be won by electrolysis instead.
6.13 Oxidation And Reduction
Twice already in this chapter a reaction has been described by saying that oxygen was taken away from one substance and given to another. It is time to give that idea its proper name, and then to widen it, because the idea turns out to be much bigger than oxygen.
Oxidation is a reaction in which oxygen combines with another substance, either wholly or in part. Reduction is a reaction in which a compound is deprived of all, or part, of the oxygen it contains.
Take the reaction you met when lead(II) oxide was heated in a stream of hydrogen.
\[\\mathrm{PbO}\\text{(s)} + \\mathrm{H\_2}\\text{(g)} \\rightarrow \\mathrm{Pb}\\text{(s)} + \\mathrm{H\_2O}\\text{(l)}\]
Oxygen has been added to the hydrogen, so the hydrogen has been oxidised to water. Oxygen has been removed from the lead(II) oxide, so it has been reduced to lead. Both changes happened in the same reaction, and neither could have happened without the other.
The substance that supplied the oxygen is called the oxidising agent — here, lead(II) oxide. The substance that took it away is the reducing agent — here, hydrogen.
Widening The Idea Beyond Oxygen
Hydrogen combines with oxygen so readily that chemists came to think of hydrogen as a kind of chemical opposite of oxygen. So the term oxidation was widened. A substance is also said to be oxidised when it gives up some or all of its hydrogen.
\[\\mathrm{H\_2S}\\text{(g)} + \\mathrm{Cl\_2}\\text{(g)} \\rightarrow 2,\\mathrm{HCl}\\text{(g)} + \\mathrm{S}\\text{(s)}\]
The chlorine has stripped the hydrogen out of hydrogen sulphide, so hydrogen sulphide has been oxidised to sulphur and chlorine is the oxidising agent. Looked at the other way, hydrogen sulphide has handed hydrogen to the chlorine, so hydrogen sulphide is acting as a reducing agent.
Oxidation cannot take place in a reaction without reduction occurring at the same time. Consider the steam and red-hot iron of section 6.7.
\[3,\\mathrm{Fe}\\text{(s)} + 4,\\mathrm{H\_2O}\\text{(g)} \\rightarrow \\mathrm{Fe\_3O\_4}\\text{(s)} + 4,\\mathrm{H\_2}\\text{(g)}\]
Oxygen has been added to iron, so the iron has been oxidised, and water vapour is therefore the oxidising agent in this reaction. Iron has removed oxygen from the water vapour, so iron is the reducing agent.
Now look at water in a different reaction.
\[2,\\mathrm{H\_2O}\\text{(l)} + 2,\\mathrm{F\_2}\\text{(g)} \\rightarrow 4,\\mathrm{HF}\\text{(aq)} + \\mathrm{O\_2}\\text{(g)}\]
Here water transfers hydrogen to fluorine, and so it is behaving as a reducing agent — the exact opposite of its role a moment ago. We must not regard being an oxidising agent or a reducing agent as a fixed property of a substance. Each reaction has to be examined on its own to see which part a given substance is playing.
6.13.1 Oxidation
Behind every one of these changes lies a movement of electrons, and that gives the most general definition of all.
\[\\mathrm{Zn} \\rightarrow \\mathrm{Zn^{2+}} + 2e^{-}\]
Oxidation is also defined as a chemical process that involves any one of the following four changes.
- Addition of oxygen. An element combines with oxygen to give its oxide.
\[\\mathrm{C} + \\mathrm{O\_2} \\rightarrow \\mathrm{CO\_2}\]
\[2,\\mathrm{Mg} + \\mathrm{O\_2} \\rightarrow 2,\\mathrm{MgO}\]
Carbon and magnesium, by gaining oxygen, are oxidised to carbon dioxide and magnesium oxide respectively.
- Addition of an electronegative ion.
\[2,\\mathrm{FeCl\_2} + \\mathrm{Cl\_2} \\rightarrow 2,\\mathrm{FeCl\_3}\]
\[\\mathrm{Zn} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2}\]
Ferrous chloride and zinc are oxidised to ferric chloride and zinc chloride by the gain of the electronegative ion \(\\mathrm{Cl^{-}}\).
- Removal of hydrogen.
\[\\mathrm{H\_2S} + \\mathrm{Cl\_2} \\rightarrow 2,\\mathrm{HCl} + \\mathrm{S}\]
Hydrogen sulphide is oxidised to sulphur by loss of hydrogen.
- Removal of an electropositive ion.
\[2,\\mathrm{KI} + \\mathrm{H\_2O\_2} \\rightarrow 2,\\mathrm{KOH} + \\mathrm{I\_2}\]
Potassium iodide is oxidised to iodine by loss of the electropositive ion \(\\mathrm{K^{+}}\).
6.13.2 Reduction
Reduction is the mirror image of oxidation at every point.
\[\\mathrm{Cu^{2+}} + 2e^{-} \\rightarrow \\mathrm{Cu}\]
Reduction is also defined as a chemical process that involves any one of the following.
- Removal of oxygen.
\[\\mathrm{CuO} + \\mathrm{H\_2} \\rightarrow \\mathrm{Cu} + \\mathrm{H\_2O}\]
\[\\mathrm{ZnO} + \\mathrm{C} \\rightarrow \\mathrm{Zn} + \\mathrm{CO}\]
Copper(II) oxide and zinc(II) oxide are reduced to copper and zinc respectively, by losing oxygen.
- Addition of an electropositive ion.
\[2,\\mathrm{HgCl\_2} + \\mathrm{SnCl\_2} \\rightarrow \\mathrm{Hg\_2Cl\_2} + \\mathrm{SnCl\_4}\]
Mercuric chloride is reduced to mercurous chloride by gain of the electropositive ion \(\\mathrm{Hg^{+}}\).
- Addition of hydrogen.
\[\\mathrm{Cl\_2} + \\mathrm{H\_2S} \\rightarrow 2,\\mathrm{HCl} + \\mathrm{S}\]
\[2,\\mathrm{NH\_3} + 3,\\mathrm{Cl\_2} \\rightarrow \\mathrm{N\_2} + 6,\\mathrm{HCl}\]
Chlorine is reduced to hydrogen chloride by gain of hydrogen.
- Removal of an electronegative ion.
\[2,\\mathrm{FeCl\_3} + \\mathrm{H\_2S} \\rightarrow 2,\\mathrm{FeCl\_2} + 2,\\mathrm{HCl} + \\mathrm{S}\]
Ferric chloride is reduced to ferrous chloride by loss of the electronegative ion \(\\mathrm{Cl^{-}}\).
Set the two lists side by side, as in Table 6.12, and every line of one is the reverse of the corresponding line of the other.
6.13.3 Oxidising Agents And Reducing Agents
An oxidising agent is one that oxidises other substances, either by accepting electrons, or by providing oxygen or an electronegative ion, or by removing hydrogen or an electropositive ion. In the reactions listed under oxidation above, the oxidising agents were oxygen, chlorine, hydrochloric acid and hydrogen peroxide.
A reducing agent is one that reduces other substances, either by providing electrons, or by providing hydrogen or an electropositive ion, or by removing oxygen or an electronegative ion. In the reactions listed under reduction, the reducing agents were hydrogen, carbon, stannous chloride, hydrogen sulphide and ammonia.
Some common oxidising agents:
Solids: manganese dioxide, red lead, lead dioxide, potassium permanganate (\(\\mathrm{KMnO\_4}\)), potassium dichromate (\(\\mathrm{K\_2Cr\_2O\_7}\)), bleaching powder.
Liquids: hydrogen peroxide, concentrated nitric acid, concentrated sulphuric acid, bromine.
Gases: oxygen, ozone, chlorine, sulphur dioxide.
Some common reducing agents:
Solids: carbon, metals such as zinc, aluminium, copper and sodium, stannous chloride, glucose.
Liquids: hydrogen peroxide, hydrogen iodide, hydrogen bromide.
Gases: hydrogen sulphide, carbon monoxide, sulphur dioxide, ammonia.
Two names appear in both lists — hydrogen peroxide and sulphur dioxide. That is not a misprint. It is the point made earlier, in a different form: what a substance does depends on what it is put with.
6.13.4 Redox Reactions
If one substance in a reaction is oxidised, some other substance must necessarily be reduced. The reason is now clear from the electronic definitions: electrons do not vanish. The electrons lost during oxidation are the very electrons gained during reduction, at the same moment.
Take zinc dropped into copper sulphate solution, the displacement you met in section 6.8.
\[\\mathrm{CuSO\_4} + \\mathrm{Zn} \\rightarrow \\mathrm{ZnSO\_4} + \\mathrm{Cu}\]
Written in ionic form, with the sulphate shown as the ion it really is:
\[\\mathrm{Cu^{2+}SO\_4^{2-}} + \\mathrm{Zn} \\rightarrow \\mathrm{Zn^{2+}SO\_4^{2-}} + \\mathrm{Cu}\]
The sulphate ion is unchanged throughout. Strip it away and the real change stands out, and it can be split into two half reactions — one showing the loss of electrons, the other the gain.
\[\\mathrm{Zn} \\rightarrow \\mathrm{Zn^{2+}} + 2e^{-} \\qquad \\text{(oxidation)}\]
\[\\mathrm{Cu^{2+}} + 2e^{-} \\rightarrow \\mathrm{Cu} \\qquad \\text{(reduction)}\]
The two electrons released by the zinc atom are the two taken up by the copper ion. Add the half reactions together and the electrons cancel, leaving the overall change.
\[\\mathrm{Cu^{2+}} + \\mathrm{Zn} \\rightarrow \\mathrm{Zn^{2+}} + \\mathrm{Cu}\]
Oxidation and reduction always go side by side, and a reaction of this kind is called a redox reaction.
Diagram
Fig. 6.13-1 — Electron Transfer In A Redox Reaction Write the equation Cu²⁺ + Zn → Zn²⁺ + Cu across the centre of the figure in large type. Draw a curved arrow leaving Zn on the left, labelled ‘2e⁻ lost — oxidation’, and curving over to Cu²⁺, where it ends with the label ‘2e⁻ gained — reduction’. Beneath the equation draw a bracket under Zn and Zn²⁺ labelled ‘oxidised’, and a second bracket under Cu²⁺ and Cu labelled ‘reduced’.
Two Worked Examples Of Redox
In the reduction of copper(II) oxide by hydrogen, the hydrogen is the reducing agent and reduces copper(II) oxide to copper. At the same time the copper(II) oxide, acting as the oxidising agent, oxidises the hydrogen to water. The net reaction is therefore a redox reaction.
In the reaction between hydrogen sulphide and chlorine, hydrogen sulphide loses its hydrogen to chlorine and is thereby oxidised to sulphur; chlorine, which gains that hydrogen, is reduced to hydrochloric acid. Fig. 6.13-2 marks up both.
Diagram
Fig. 6.13-2 — Reading A Redox Equation Two labelled equations, one above the other. Upper: CuO + H₂ → Cu + H₂O, with ‘oxidising agent’ printed under CuO and ‘reducing agent’ under H₂; draw a bracket linking H₂ to H₂O above the equation labelled ‘oxidised’, and a bracket linking CuO to Cu below the equation labelled ‘reduced’. Lower: H₂S + Cl₂ → 2HCl + S, with ‘reducing agent’ printed under H₂S and ‘oxidising agent’ under Cl₂; bracket linking H₂S to S above labelled ‘oxidised’, bracket linking Cl₂ to 2HCl below labelled ‘reduced’.
6.14 Summary
Hydrogen is the first element of the periodic table. Its atomic number is 1, its electronic configuration is 1 (K shell), its valency is 1, and it is placed in Period 1, Group 1 (IA).
The gas exists as diatomic molecules, \(\\mathrm{H\_2}\), with a relative atomic mass of 1 and a relative molecular mass of 2.
Hydrogen shows a dual nature: it resembles the alkali metals of Group 1 in some properties and the halogens of Group 17 in others. Its position in the table has been debated ever since Mendeleev, and Thomson suggested a separate place for it at the top.
Like the alkali metals, hydrogen has one valence electron, a valency of 1, forms a cation by loss of that electron, acts as a reducing agent, burns in oxygen, and combines with non-metals such as oxygen, sulphur and chlorine.
Like the halogens, hydrogen is one electron short of the nearest inert gas, has a valency of 1, forms an anion, is electronegative, is a gas, and has an atomicity of 2.
Hydrogen differs from both families in two ways: its atom has only one shell, and its oxide, water, is neutral, whereas the oxides of the alkali metals are basic and those of the halogens acidic.
Robert Boyle prepared the gas in 1672. Henry Cavendish, in 1766, showed that it was a distinct substance and that it burns to form water, and called it inflammable air. Lavoisier named it hydrogen, the water former, in 1783.
In the free state hydrogen occurs only in traces on the earth, but in large amounts in the Sun and the stars. In the combined state it occurs in water, acids, alkalis, hydrocarbons and all plant and animal tissues. Water is 11.1% hydrogen by weight.
Potassium, sodium and calcium react with cold water, forming their hydroxides and liberating hydrogen. All three reactions are exothermic and all three solutions are alkaline.
\[2,\\mathrm{K} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{KOH} + \\mathrm{H\_2}\]
\[2,\\mathrm{Na} + 2,\\mathrm{H\_2O} \\rightarrow 2,\\mathrm{NaOH} + \\mathrm{H\_2}\]
\[\\mathrm{Ca} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{Ca(OH)\_2} + \\mathrm{H\_2}\]
- Magnesium reacts slowly with boiling water to give the hydroxide, but burns in steam with an intense white light to give the oxide.
\[\\mathrm{Mg} + 2,\\mathrm{H\_2O} \\rightarrow \\mathrm{Mg(OH)\_2} + \\mathrm{H\_2}\\uparrow \\qquad \\mathrm{Mg} + \\mathrm{H\_2O} \\rightarrow \\mathrm{MgO} + \\mathrm{H\_2}\\uparrow\]
- Aluminium, zinc and iron react only with steam, and only when heated. Aluminium and zinc stop reacting because their oxides coat the metal.
\[2,\\mathrm{Al} + 3,\\mathrm{H\_2O} \\xrightarrow{\\ 800^\\circ\\text{C}\\ } \\mathrm{Al\_2O\_3} + 3,\\mathrm{H\_2} \\qquad \\mathrm{Zn} + \\mathrm{H\_2O} \\xrightarrow{\\ \\Delta\\ } \\mathrm{ZnO} + \\mathrm{H\_2}\]
- The reaction of steam with red-hot iron is reversible, and reaches chemical equilibrium at 700 °C if the hydrogen is not removed.
\[3,\\mathrm{Fe} + 4,\\mathrm{H\_2O} \\rightleftharpoons \\mathrm{Fe\_3O\_4} + 4,\\mathrm{H\_2}\]
Steam passed over red-hot coke gives water gas, a mixture of carbon monoxide and hydrogen. The carbon monoxide is dissolved out in ammoniacal cuprous chloride solution.
The activity series is the arrangement of metals in decreasing order of reactivity: K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au. Hydrogen is included because it forms a positive ion.
Down the series, electropositive character, reducing power and the tendency to be oxidised all decrease. A metal higher in the series displaces one lower down from a solution of its salt.
Metals above hydrogen displace hydrogen from dilute hydrochloric acid and dilute sulphuric acid, forming their salts; metals below hydrogen do not.
Zinc, lead and aluminium have a unique nature: they give hydrogen with acids and also with hot concentrated alkalis, because their oxides and hydroxides are amphoteric.
\[\\mathrm{Zn} + 2,\\mathrm{NaOH} \\rightarrow \\mathrm{Na\_2ZnO\_2} + \\mathrm{H\_2}\]
- In the laboratory, hydrogen is prepared from granulated zinc and dilute hydrochloric or dilute sulphuric acid. Zinc is chosen because it is cheap, lies above hydrogen, needs no heating and forms soluble salts; granulated zinc is preferred to pure zinc because the copper in it catalyses the reaction.
\[\\mathrm{Zn} + 2,\\mathrm{HCl} \\rightarrow \\mathrm{ZnCl\_2} + \\mathrm{H\_2}\\uparrow\]
The gas is collected by downward displacement of water, because it is almost insoluble in water and because it forms an explosive mixture with air.
The impurities present are hydrogen sulphide, sulphur dioxide, oxides of nitrogen, phosphine, arsine, carbon dioxide and water vapour. They are removed by silver nitrate solution, lead nitrate solution, caustic potash solution, and a drying agent such as anhydrous calcium chloride.
Concentrated sulphuric acid is not used to dry hydrogen, and nitric acid and concentrated sulphuric acid are not used to prepare it, because all of them oxidise the hydrogen. Magnesium and manganese alone give hydrogen with very dilute nitric acid.
Hydrogen is combustible but does not support combustion. It burns with a pale blue flame to form water, and a premixed sample explodes with a pop sound — the test for the gas.
On a large scale hydrogen is manufactured by the Bosch process, from steam and hot coke at 1000 °C followed by conversion over \(\\mathrm{Fe\_2O\_3}\) with \(\\mathrm{Cr\_2O\_3}\) at 450 °C, and by the electrolysis of acidulated water, which gives hydrogen at the cathode and oxygen at the anode.
Hydrogen is colourless, odourless, tasteless and non-toxic, has a density of \(0.0899\\ \\text{g L}^{-1}\), is 14.4 times less dense than air, and is very slightly soluble in water. It is occluded by palladium, platinum and nickel.
Hydrogen reacts with non-metals to give water, hydrogen chloride, ammonia and hydrogen sulphide; with heated potassium, sodium and calcium to give their hydrides; and with the oxides of the less active metals, reducing them to the metal.
In the presence of a nickel catalyst, hydrogen adds across double and triple bonds in organic compounds. This hydrogenation converts liquid vegetable oils into solid vanaspati ghee.
Hydrogen is used as a fuel, in the oxy-hydrogen and atomic hydrogen torches, in the manufacture of ammonia, in hydrogenation of oils, in producing artificial petrol from coal, and in the extraction of metals from their oxides.
Oxidation is the addition of oxygen or an electronegative ion, the removal of hydrogen or an electropositive ion, or the loss of electrons. Reduction is the exact reverse of each of these.
An oxidising agent oxidises another substance and is itself reduced; a reducing agent reduces another substance and is itself oxidised. Whether a substance acts as one or the other depends on the reaction it is in.
Oxidation and reduction always occur together, because the electrons lost in one are the electrons gained in the other. Such a reaction is called a redox reaction.
\[\\mathrm{Zn} \\rightarrow \\mathrm{Zn^{2+}} + 2e^{-} \\qquad \\mathrm{Cu^{2+}} + 2e^{-} \\rightarrow \\mathrm{Cu}\]
Flowchart
FC 6.3 — The Routes To Hydrogen In This Chapter Cold water + K, Na or Ca → hydroxide + hydrogen; branch: hot water + Mg → hydroxide + hydrogen; branch: steam + Al, Zn or red-hot Fe → oxide + hydrogen; branch: steam + red-hot coke → water gas → CO removed by ammoniacal cuprous chloride → hydrogen; branch: dilute HCl or dilute H₂SO₄ + Mg, Al, Zn or Fe → salt + hydrogen (the laboratory route, using granulated zinc); branch: hot concentrated NaOH or KOH + Zn, Pb or Al → soluble salt + hydrogen; branch: Bosch process, steam + coke then conversion and removal of CO₂ and CO → hydrogen; branch: electrolysis of acidulated water → hydrogen at cathode, oxygen at anode
6.15 Key Terms
Acidic oxide — an oxide whose solution in water turns blue litmus red.
Activity series — the arrangement of metals in decreasing order of their reactivity.
Alkaline — describing a solution that turns red litmus blue.
Amalgam — an alloy of a metal with mercury.
Amphoteric — describing an oxide or hydroxide that reacts with both acids and bases to give a salt and water.
Anion — a negatively charged ion, formed when an atom gains one or more electrons.
Anode — the positively charged electrode, at which negative ions give up their electrons.
Atomic hydrogen torch — a torch reaching about 2800 °C, used for welding high-melting alloys.
Atomicity — the number of atoms present in one molecule of an element.
Basic oxide — an oxide whose solution in water turns red litmus blue.
Catalyst — a substance that alters the rate of a reaction without being used up in it.
Cathode — the negatively charged electrode, at which positive ions take up electrons.
Cation — a positively charged ion, formed when an atom loses one or more electrons.
Chemical equilibrium — the stage of a reversible reaction at which the forward and backward reactions proceed at the same speed.
Combined state — an element as it occurs in nature joined to other elements in compounds.
Diatomic — made of two atoms, as a molecule of hydrogen or chlorine is.
Downward displacement of water — collecting a gas in an inverted jar full of water, which the gas pushes out.
Drying agent — a substance that removes water vapour from a gas without reacting with the gas.
Dual nature — the property of hydrogen by which it resembles both the alkali metals and the halogens.
Effervescence — the fizzing of a liquid as bubbles of gas are released in it.
Electrolysis — the decomposition of a substance by passing an electric current through it.
Electronegative — describing an element whose atoms tend to accept electrons and form anions.
Electropositive — describing an element whose atoms tend to lose electrons and form cations.
Endothermic — describing a reaction that absorbs heat as it proceeds.
Exothermic — describing a reaction that gives out heat as it proceeds.
Free state — an element as it occurs in nature uncombined with any other element.
Group — a vertical column of the periodic table, whose elements have the same number of valence electrons.
Half reaction — one of the two parts of a redox reaction, showing either the loss or the gain of electrons.
Hydride — a compound of a metal with hydrogen, containing the hydrogen as \(\\mathrm{H^{-}}\).
Hydrogenation — the addition of hydrogen across a double or triple bond, in the presence of a catalyst.
Inert gas — an element of Group 18, whose atoms have a complete outer shell and take almost no part in chemical change.
Inflammable air — the original name given to hydrogen by Cavendish, because the gas burns readily.
Neutral oxide — an oxide that changes the colour of neither red nor blue litmus.
Occlusion — the taking up of a gas by a solid, as palladium takes up hydrogen.
Oxidation — the addition of oxygen or of an electronegative ion, the removal of hydrogen or of an electropositive ion, or the loss of electrons.
Oxidising agent — a substance that oxidises another substance and is itself reduced.
Oxy-hydrogen torch — an apparatus in which hydrogen is burnt in oxygen to give a flame reaching about 2500 °C.
Period — a horizontal row of the periodic table.
Peroxide — an oxide containing more oxygen than the ordinary oxide of the same element, as in \(\\mathrm{Na\_2O\_2}\).
Pop sound — the small explosion heard when hydrogen premixed with air is lit; the test for the gas.
Promoter — a substance added to make a catalyst work more effectively.
Redox reaction — a reaction in which oxidation and reduction take place simultaneously.
Reducing agent — a substance that reduces another substance and is itself oxidised.
Reduction — the removal of oxygen or of an electronegative ion, the addition of hydrogen or of an electropositive ion, or the gain of electrons.
Relative atomic mass — the mass of one atom of an element compared with a standard atom; it has no unit.
Relative molecular mass — the sum of the relative atomic masses of all the atoms in one molecule.
Reversible reaction — a reaction that proceeds in both directions, so that the products can react to give back the reactants.
Superoxide — an oxide of the general formula \(\\mathrm{MO\_2}\), formed by potassium, rubidium and caesium.
Unique nature — the behaviour of zinc, lead and aluminium in giving hydrogen with acids and with hot concentrated alkalis alike.
Valence shell — the outermost shell of an atom.
Valency — the combining capacity of an atom, measured by the number of hydrogen atoms it can combine with or replace.
Water gas — the mixture of carbon monoxide and hydrogen obtained by passing steam over red-hot coke.
6.16 Points To Ponder
Hydrogen is placed at the head of Group 1, yet you have seen that it behaves as much like a halogen as like an alkali metal. If you were drawing up the periodic table afresh, where would you put it — and what would you be forced to give up by each choice you might make?
The activity series in this chapter was built from how metals behave with water and with dilute acids. Suppose it had been built instead from how the same metals behave with chlorine, or with oxygen. Would the order come out the same? What would it mean if it did?
Six practical objections ruled out every metal except zinc for the laboratory preparation, and not one of them was about the chemistry being wrong — cost, safety, heating and solubility decided it. How far do such practical considerations shape what is actually done in a laboratory or a factory?
A reversible reaction that has reached equilibrium looks, from the outside, exactly like a reaction that has stopped: nothing changes any more. What evidence could you look for to tell the two situations apart?
Hydrogen peroxide appears in this chapter both as an oxidising agent and as a reducing agent, and so does sulphur dioxide. Does this mean the labels are unreliable, or does it tell you something about what kind of property being an oxidising agent really is?
Burning hydrogen produces water and nothing else, yet making that hydrogen by the Bosch process leaves carbon dioxide behind at the factory. When we call a fuel clean, at which point in its history should we be looking?
6.17 Objective Questions
Multiple Choice
- The element that has one valence electron and is at the same time one electron short of the nearest inert gas configuration is [1]
- lithium
- hydrogen
- chlorine
- helium
- Hydrogen is collected over water rather than by downward displacement of air because [1]
- it is heavier than air
- it is highly soluble in water
- it forms an explosive mixture with air
- it reacts with the glass of the jar
- The metal that sinks in cold water and yet liberates hydrogen from it is [1]
- sodium
- potassium
- calcium
- magnesium
- In the second stage of the Bosch process, chromic oxide acts as [1]
- the catalyst
- the promoter
- a drying agent
- a reducing agent
- Zinc oxide is best described as [1]
- basic
- acidic
- neutral
- amphoteric
- In the reaction \(3,\\mathrm{Fe} + 4,\\mathrm{H\_2O} \\rightarrow \\mathrm{Fe\_3O\_4} + 4,\\mathrm{H\_2}\), the oxidising agent is [1]
- iron
- steam
- triferric tetraoxide
- hydrogen
Assertion–Reason
- For each of the following, choose from the codes given below.
- Both A and R are true and R is the correct explanation of A.
- Both A and R are true but R is not the correct explanation of A.
- A is true but R is false.
- A is false but R is true.
Assertion (A): Granulated zinc is used in place of pure zinc in the laboratory preparation of hydrogen. Reason (R): The traces of copper present in granulated zinc have a catalysing effect on the reaction. [1]
Assertion (A): Concentrated sulphuric acid is not used to dry hydrogen. Reason (R): Concentrated sulphuric acid reacts with hydrogen. [1]
Assertion (A): Aluminium does not readily give hydrogen when treated with a dilute acid. Reason (R): Aluminium lies below hydrogen in the activity series. [1]
Assertion (A): The reaction between red-hot iron and steam is a reversible reaction. Reason (R): The hydrogen formed can reduce the iron oxide back to iron. [1]
Match The Following
- Match each item in Column A with the correct item in Column B. [5]
| Column A | Column B |
|---|---|
| a. Lead nitrate solution | i. removes carbon monoxide |
| b. Ammoniacal cuprous chloride solution | ii. catalyst for the hydrogenation of oils |
| c. Anhydrous calcium chloride | iii. removes hydrogen sulphide |
| d. Finely divided iron | iv. removes water vapour |
| e. Nickel | v. catalyst in the Haber process |
Fill In The Blanks
- Complete each statement. [4]
- The atomicity of hydrogen is ____________.
- Magnesium reacts with steam to form ____________ and hydrogen.
- During the electrolysis of acidulated water, the gas collected at the anode is ____________.
- Metals lying ____________ hydrogen in the activity series cannot displace it from dilute acids.
True Or False
- State whether each statement is true or false. [4]
- Water is a basic oxide.
- Hydrogen is combustible but does not support combustion.
- Lead reacts with hot concentrated sodium hydroxide solution to liberate hydrogen.
- Pure hydrogen burns in air with a lilac flame.
Name The Following
- Name each of the following. [4]
- The gas obtained along with hydrogen when steam is passed over red-hot coke.
- The alloy of sodium with mercury from which hydrogen can be prepared safely.
- The solid fat obtained by the hydrogenation of liquid vegetable oils.
- The metal that takes up about 900 times its own volume of hydrogen at room temperature.
Give One Word Or Phrase
- Give one word or phrase for each of the following. [3]
- The taking up of a gas by a solid.
- A reaction in which oxidation and reduction take place at the same time.
- An oxide that reacts with acids as well as with alkalis.
Odd One Out
- Pick the odd one out in each set and give a reason for your choice. [3]
- Potassium, sodium, calcium, copper
- Zinc, lead, aluminium, iron
- Hydrogen sulphide, carbon monoxide, chlorine, ammonia
6.18 Check Your Concepts
State two respects in which hydrogen resembles the halogens. [2]
Give a reason why hydrogen, although lighter than air, is not collected by the downward displacement of air. [2]
Define atomicity, and state the atomicity of hydrogen. [2]
Calcium is added to cold water and a drop of red litmus solution is then added to the mixture. Explain the colour change that follows. [2]
Differentiate between an oxidising agent and a reducing agent, stating clearly the basis of the difference. [3]
Give a reason for each of the following. [3]
- Sodium is stored under kerosene oil.
- The lower end of the thistle funnel must dip below the surface of the acid.
- No flame is allowed to burn near the apparatus during the preparation of hydrogen.
Explain why zinc is preferred to both iron and lead for the laboratory preparation of hydrogen. [3]
Name three properties of metals that change as you move down the activity series, and state in each case the direction of the change. [3]
Differentiate between the action of hot water and the action of steam on magnesium, stating the basis of the difference and naming the product formed in each case. [4]
What would happen in each case, and why? [4]
- Dilute nitric acid is used in place of dilute sulphuric acid in the laboratory preparation of hydrogen.
- Pure zinc is used in place of granulated zinc.
Explain, with reference to its behaviour with acids and with alkalis, why aluminium is said to have a unique nature. [4]
Describe how the impurities present in hydrogen prepared in the laboratory are removed, naming the reagent used at each stage and the impurity it removes. [4]
Describe the manufacture of hydrogen by the Bosch process, giving the conditions used at each stage and explaining how the carbon dioxide and carbon monoxide are separated from the hydrogen. [5]
With reference to the action of steam on red-hot iron, explain what is meant by a reversible reaction and by chemical equilibrium, and state how this reaction can be made to proceed continuously in one direction. [5]
6.19 Equations, Conversions & Identification
Write balanced chemical equations for the following. [4]
- Potassium reacts with cold water.
- Magnesium reacts with steam.
- Zinc reacts with hot concentrated potassium hydroxide solution.
- Aluminium reacts with dilute sulphuric acid.
Complete and balance the following equations. [4]
- \(\\mathrm{Fe} + \\mathrm{H\_2O} \\rightarrow\)
- \(\\mathrm{Ca} + \\mathrm{H\_2O} \\rightarrow\)
- \(\\mathrm{Pb} + \\mathrm{NaOH} \\rightarrow\)
- \(\\mathrm{Fe\_2O\_3} + \\mathrm{H\_2} \\rightarrow\)
Granulated zinc is added to dilute hydrochloric acid. Name the gas evolved and describe the test by which you would identify it. [2]
State what you would observe in each of the following. [3]
- A small piece of potassium is dropped into a trough of cold water.
- Dry hydrogen is passed over heated copper(II) oxide.
- A piece of calcium is dropped into cold water.
How would you obtain the following? [4]
- Hydrogen from water gas.
- Hydrogen from zinc, by two different routes.
- Copper from copper(II) oxide, using hydrogen.
A metal lies below hydrogen in the activity series. It gives no hydrogen with dilute acids, but its black oxide, when heated in a stream of hydrogen, gives a reddish-brown metal and a colourless liquid. Identify the metal and the oxide, and name the colourless liquid. [2]
A gas is lighter than air, almost insoluble in water, burns in air with a pale blue flame, and explodes with a pop sound when a premixed sample is lit. Identify the gas and state one method by which it is collected. [2]
With reference to Fig. 6.9-1, name the two pieces of apparatus fitted into the holes of the cork, and state the purpose of each. State also why the cork must be airtight. [3]
With reference to Fig. 6.10-1, name the gas collected in each of the two tubes, state which electrode each tube stands over, and give a reason for the difference in the volumes collected. [3]
For the reaction \(\\mathrm{Zn} + \\mathrm{Cu^{2+}} \\rightarrow \\mathrm{Zn^{2+}} + \\mathrm{Cu}\), write the two half reactions, state which substance has been oxidised, and name the oxidising agent. [3]