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Chapter 5: The Periodic Table

In the autumn of 1875, a French chemist called Paul-Émile Lecoq de Boisbaudran was studying the light given out by a mineral from the Pyrenees. He let the glow pass through a spectroscope, an instrument that spreads light into a band of coloured lines, and saw two violet lines that matched no element then known. Working patiently through a great deal of ore, he drew out a few grams of a soft, silvery metal so low-melting that it turned liquid in the warmth of his hand. He called it gallium.

Then a letter arrived from St Petersburg. Dmitri Mendeleev wrote that six years earlier he had left an empty square in his table of elements, in the column below aluminium, and had worked out from its neighbours what the missing element must be like — how heavy its atoms, how dense the metal, how low its melting point. The density Boisbaudran had measured, he said, was too low. Boisbaudran purified his sample again, weighed it again, and found the value Mendeleev had calculated.

A man in Russia who had never seen gallium, never held it, never tested it, had described it correctly from an empty box in a table. That is what this chapter is about: how elements were sorted, and why the sorting turned out to carry so much information.

Diagram

Fig. 5.0-1 — Two New Violet Lines In The Spectroscope A laboratory bench of the 1870s: a burner heating a sample holder, its light entering a spectroscope tube fitted with a prism, and the spread-out spectrum shown as a horizontal band of coloured lines on the right; label the burner, the sample of zinc ore, the prism, the eyepiece, the familiar bright lines of zinc, and two new violet lines marked ‘unidentified’. Beside the bench, show a small dish holding a bead of silvery metal melting in a warm palm, labelled gallium.

NoteCuriosity Corner

Q1. How could an empty square in a table be trusted enough to describe a metal that nobody had ever seen? Q2. Early chemists tried sorting elements into groups of three, and later into rows of seven. Why did neither arrangement survive? Q3. Argon is heavier than potassium, yet it is placed before potassium in today’s table. Why is this not treated as a mistake? Q4. Lithium, sodium and potassium are three different metals. What could they possibly share that makes all three behave so alike in water? Q5. Going from sodium to chlorine, each atom gains more protons and more electrons — so why does the atom get smaller rather than bigger?

5.1 Reasons For Classification Of Elements

Sorting things out is something you do without being told. Spoons go in one drawer, knives in another; textbooks stand on one shelf, storybooks on the next. Putting like with like is the first step in making sense of a large collection.

Chemists faced exactly this problem, on a much larger scale. Towards the end of the eighteenth century and the beginning of the nineteenth, new elements were being discovered one after another. Each had its own properties, and each formed its own set of compounds. Studying them one by one was becoming impossible, and comparing them was harder still.

5.1.1 The Need For A Scheme Of Classification

An element is a pure substance made of only one kind of atom, and cannot be broken down into anything simpler by a chemical change. Once dozens of such substances were known, chemists began looking for a way of arranging them that would make their study manageable. Three gains were expected from such an arrangement.

  1. It allows the elements to be studied in an organised way, family by family, instead of one at a time.
  2. It allows the properties of an element to be linked with the fundamental properties of matter in all its states.
  3. It allows the relationship of one element to another to be stated clearly, rather than guessed at.

The third point is worth pausing on. A good arrangement does not merely store information about elements that are already known — it lets you say something about an element you have not yet met, simply from where it sits. You saw this at the start of the chapter, and the rest of the chapter explains how it became possible.

Roughly 30 elements were known in 1789. By the time Mendeleev began his work, the count was 63. Today the table holds 118, the heaviest of them made a few atoms at a time in accelerators.

5.1.2 The Earliest Attempts And Why They Failed

The first attempts used properties that a chemist could measure directly in the laboratory: density, malleability — how far a solid can be beaten into sheets — and ductility, how far it can be drawn into wire. Elements were also sorted in a broader way, into metals and non-metals. Metals are lustrous solids that conduct heat and electricity well; non-metals mostly do not.

Antoine Lavoisier’s Table Of Simple Substances In 1789, in his textbook of chemistry, the French chemist Antoine Lavoisier printed a list of substances that could not be broken down further. He grouped them as gases, metals, non-metals and earths. It was the first serious attempt at classification — though two of his ‘simple substances’, light and heat, are not elements at all.

None of these early schemes proved satisfactory, and for three plain reasons.

  1. Far too many elements fell into a single group. A group holding a large fraction of all known elements tells you very little about any one of them.
  2. Some of the properties used were not fixed. Density and malleability shift with temperature, with pressure, and with how a sample was prepared, so two chemists could sort the same element differently.
  3. Some elements refused to sit on either side of the metal–non-metal line, showing a little of both characters.

Notice what all three failures have in common. The property being used to sort was either too coarse to separate elements, or not constant enough to be relied on. What chemists needed was a property of an element that never changes, whatever the sample and whatever the conditions.

Take twenty objects from around you and sort them by colour, then re-sort them by material. Which sorting puts the same things together each time you repeat it, and which one depends on your judgement? That difference is exactly the problem the early chemists ran into.

The Search For A Fixed Property

Attention turned to atomic mass, the mass of an atom of an element compared with a standard atom. In 1815 the English chemist William Prout suggested that the atomic masses of the elements are whole-number multiples of the atomic mass of hydrogen, hinting that atomic mass was a fixed and fundamental quantity rather than an accident of the sample.

Prout’s exact claim did not survive later measurement, but the shift in thinking did. From this point onwards, chemists tried to classify elements by atomic mass — and the schemes in the next three sections were all built on it.

5.2 Döbereiner’s Triads

Once atomic mass was accepted as a fixed property of an element, chemists began laying the elements out in order of it, looking for a pattern. The first person to find one was a German chemist, J. W. Döbereiner.

5.2.1 The Law Of Triads

Döbereiner noticed that certain elements fall naturally into sets of three, in which all three members behave alike chemically. Such a set of three elements is called a triad. Lithium, sodium and potassium make one such set.

What made the observation important was the second thing he noticed. Within a triad, the atomic masses are not scattered at random: the middle element sits almost exactly halfway between the other two. Döbereiner was the first chemist to show a clear link between the properties of an element and its atomic mass.

NotePrinciple / Law

Döbereiner’s Law Of Triads When three elements with similar chemical properties are arranged in order of increasing atomic mass, the atomic mass of the middle element is approximately equal to the arithmetic mean (the average) of the atomic masses of the other two.

The arithmetic mean of two numbers is their sum divided by two. If the atomic masses of the first and third members of a triad are \(A\) and \(C\), then the atomic mass of the middle member \(B\) is given by

\[B \approx \frac{A + C}{2}\]

Relative Atomic Mass An atomic mass such as 40 for calcium is a comparison, not a weight in grams. It says the atom is 40 times as heavy as one-twelfth of a carbon-12 atom, so it carries no unit.

5.2.2 Examples Of Döbereiner’s Triads

Two examples show what a triad looks like — one made of metals, one made of non-metals.

Calcium, Strontium And Barium

These three elements resemble one another closely.

  1. All three are metals.
  2. Each forms an oxide that is alkaline in nature.
  3. Each has a valency of 2 — that is, one atom of the element combines with, or replaces, two atoms of hydrogen.

Because their chemical behaviour is so alike, they were placed together in one family. Their atomic masses are 40 for calcium, 88 for strontium and 137 for barium. Taking the first and the last,

\[\frac{40 + 137}{2} = \frac{177}{2} = 88.5\]

which is very nearly the atomic mass of strontium, 88. The relationship is easier to see on a scale of atomic mass, as Fig. 5.2-1 shows.

Diagram

Fig. 5.2-1 — A Triad Seen On A Scale Of Atomic Mass A horizontal number line of atomic mass running from 0 to 150, with three labelled markers: calcium at 40, strontium at 88 and barium at 137; a dashed vertical line at 88.5 labelled ‘mean of 40 and 137’, drawn almost on top of the strontium marker; a curved brace joining calcium and barium with the label ‘first and third members of the triad’.

Chlorine, Bromine And Iodine

The second triad is made of non-metals, and the same rule holds.

  1. All three are non-metals.
  2. All three react with water to form acids.
  3. All three have a valency of 1.

Their atomic masses are 35.5 for chlorine, 80 for bromine and 127 for iodine.

\[\frac{35.5 + 127}{2} = \frac{162.5}{2} = 81.25\]

This is close to 80, the atomic mass of bromine. Some further triads that Döbereiner and others recognised are collected in Table 5.1.

Triad First element Middle element Third element
Alkali metals \(^{7}\mathrm{Li}\) \(^{23}\mathrm{Na}\) \(^{39}\mathrm{K}\)
Nitrogen family \(^{14}\mathrm{N}\) \(^{31}\mathrm{P}\) \(^{75}\mathrm{As}\)
Oxygen family \(^{32}\mathrm{S}\) \(^{79}\mathrm{Se}\) \(^{128}\mathrm{Te}\)
Iron family \(^{56}\mathrm{Fe}\) \(^{59}\mathrm{Co}\) \(^{59}\mathrm{Ni}\)

Table 5.1 — Some other triads, with the atomic mass written above each symbol

Example 5.1 — Predicting The Atomic Mass Of The Middle Element

1. Sulphur (atomic mass 32) and tellurium (atomic mass 128) are the first and third members of a triad. Predict the atomic mass of the middle element, selenium, and compare it with the accepted value of 79.

  1. By the law of triads, the middle element’s atomic mass is the arithmetic mean of the other two. Adding the two given masses gives \(32 + 128 = 160\). Dividing by 2 gives \[\frac{32 + 128}{2} = \frac{160}{2} = 80\]

  2. The predicted atomic mass is 80, against an accepted value of 79 for selenium. The prediction is close, but not exact — and that small gap is a fair warning of what happens to the law when it is tested more widely.

5.2.3 Reasons For Discarding The Law Of Triads

A rule in science is judged by how far it stretches, not by how well it fits the cases that suggested it. The law of triads was abandoned for two reasons.

  1. Döbereiner could not arrange all the elements known at the time into triads. Only a handful of sets obeyed the rule, and the great majority of elements were left out of the scheme altogether.
  2. The rule failed even inside a family whose members clearly belong together. Take the first three halogens: fluorine (19), chlorine (35.5) and bromine (80). The mean of the first and the third is

\[\frac{1}{2}(19 + 80) = 49.5\]

which is nowhere near 35.5, the atomic mass of chlorine.

So Döbereiner’s scheme was not successful as a classification. It was, however, far from worthless: it established that elements with similar properties have atomic masses that are related to one another, and that idea survived every scheme that followed.

Lithium, sodium and potassium have atomic masses 7, 23 and 39. Show whether they satisfy the law of triads, and state clearly which two masses you use in the calculation and why.

5.3 Newland’s Law Of Octaves

Döbereiner had shown that similar elements come in related sets, but he had only reached small islands of elements with a wide sea between them. The next step was taken by John Newland, an English chemist who was also a keen musician.

5.3.1 Arranging The Elements In Octaves

Newland did something simple and bold. Instead of hunting for sets of three, he wrote out all the elements known to him in one long line, in ascending order of atomic mass, and then looked along the line for repetitions.

He found that starting from any element and counting onwards, the eighth element resembled the first in its properties. The comparison he drew was musical. In a scale of eight notes — Do, Re, Me, Fa, So, La, Ti and back to Do, or Sa, Re, Ga, Ma, Pa, Dha, Nee and back to Sa — the eighth note is the same note as the first, sounded higher. Such a set of eight is called an octave, and Newland gave the elements the same name.

NotePrinciple / Law

Newland’s Law Of Octaves When the elements are arranged in the increasing order of their atomic masses, the properties of every eighth element, counting from any element, are a repetition of the properties of that starting element.

Diagram

Fig. 5.3-1 — The Musical Analogy Behind The Law Of Octaves Two parallel strips. The upper strip shows one octave on a keyboard, the eight keys labelled Sa, Re, Ga, Ma, Pa, Dha, Nee, Sa, with the first and last keys shaded alike and joined by an arrow labelled ‘same note, sounded higher’. The lower strip shows eight element boxes in order of atomic mass — Li, Be, B, C, N, O, F, Na — with the Li and Na boxes shaded alike and joined by an arrow labelled ‘similar properties’.

To display the repetition, Newland set the elements out in horizontal rows of seven, so that members of the same family fell one below the other in a vertical column. His arrangement is shown in Table 5.2.

Western music Do Re Me Fa So La Ti
Indian music Sa Re Ga Ma Pa Dha Nee
1 2 3 4 5 6 7
H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
Co and Ni Cu Zn Y In As Se
Br Rb Sr Ce and La Zr — —

Table 5.2 — Newland’s arrangement of the elements in octaves

Seven columns, and not eight, is the point to notice. The eighth element was not given a fresh column; it began the next row, directly under the element it resembled. The noble gases would have upset this count badly, but they had not yet been discovered, so Newland never had to deal with them.

Triads And Octaves A triad is a set of three similar elements, and the test applied to it is arithmetic — the middle atomic mass is the mean of the other two. An octave is a repetition after every eighth element, and the test applied to it is a resemblance in properties.

5.3.2 Merits Of Newland’s Classification

Newland’s arrangement failed as a complete scheme, but three things about it were genuine advances.

  1. It worked quite well for the lighter elements. Lithium, sodium and potassium, for instance, came together in one column, exactly as their chemistry demands.
  2. It related the properties of the elements to their atomic masses, over the whole range of light elements rather than in isolated sets of three.
  3. For the first time, it showed that there is a distinct periodicity in the properties of elements — that is, that similar properties reappear at regular intervals along the list.

A Medal, Twenty Years Late When Newland first presented his law to a meeting of chemists in 1866, it was not taken seriously; one member asked whether he had thought of arranging the elements alphabetically instead. In 1887, long after the periodic table had proved him broadly right, the Royal Society awarded him the Davy Medal for his work on the classification of elements.

5.3.3 Reasons For Discarding The Law Of Octaves

The law held only for a short stretch of the list, and broke down badly beyond it.

  • It did not work for the heavier elements, that is, for those lying beyond calcium. As more and more elements were discovered, there was simply no way to fit them into the pattern of octaves.
  • Newland forced two elements, cobalt and nickel, into a single slot in order to keep the count of seven working. That slot fell in the same column as fluorine, chlorine and bromine — elements whose properties are nothing like those of cobalt and nickel.
  • Iron, which does resemble cobalt and nickel closely, was placed far away from them, in a different column altogether.

The last two failures are the more serious. A classification is meant to put like with like; Newland’s table was, in places, doing the opposite, and doing it because the arrangement demanded it. That is a sign that the rule, not the chemistry, was in charge.

Q2. Early chemists tried sorting elements into groups of three, and later into rows of seven. Why did neither arrangement survive?

A. Both schemes were built on real patterns, but each fixed the size of the pattern in advance — always three, always seven. Döbereiner’s rule could gather only a few elements into triads and failed even among the halogens. Newland’s rule held for the light elements and then collapsed, because he had no room for the elements still to be discovered and none at all for the noble gases. What was needed was an arrangement that let the length of the repeat change as the elements got heavier.

Using Table 5.2, name the element that Newland’s arrangement places in the same column as chlorine but which does not belong there chemically. State the property of chlorine you would test to show that the placing is wrong.

5.4 Mendeleev’s Periodic Table

In 1869 a Russian chemist, Dmitri Ivanovich Mendeleev, took up the same task once more. He wrote out all the 63 elements known at that time in increasing order of atomic mass, and set elements with similar properties one below another in the same vertical column.

Where the repetition failed — where the next element in order of atomic mass did not match the element above it — Mendeleev did not force it into place. He left the square blank and moved the element down to where it belonged. That single decision is what separates his table from Newland’s.

Working this way, he found that elements with similar properties turned up at regular intervals along the list. He called this the periodicity of the properties of elements, and built a law on it.

5.4.1 Mendeleev’s Periodic Law

NotePrinciple / Law

Mendeleev’s Periodic Law The physical and chemical properties of elements are a periodic function of their atomic weights.

The phrase ‘a periodic function of’ is worth unpacking. It means that if you list the elements in order of atomic mass and then read off any property — say, the formula of the oxide — that property does not change randomly along the list. It rises and falls, and returns to something like its earlier value at fixed intervals.

The arrangement that results is what we now call a periodic table.

NoteDefinition

Periodic Table A chart of the elements prepared in such a way that elements with similar properties occur in the same vertical column, called a group. It is called periodic because elements with similar properties occur at fixed intervals, and a table because the elements are set out in rows and columns.

5.4.2 Essential Features Of Mendeleev’s Table

The vertical columns of Mendeleev’s table are called groups and the horizontal rows are called periods. Four features define the arrangement.

  • There are eight groups in all, Group I to Group VIII. Each of the groups from I to VII is divided into two sub-groups, A and B. Group VIII has no sub-group. The inert gases were not known at the time, so no group was set aside for them.
  • All the elements of a sub-group — or of Group VIII, which is undivided — have similar properties and show the same valency.
  • That valency is fixed by the group number. For Groups I to IV, the valency equals the group number. For Groups V to VII, it equals the group number subtracted from eight.
  • Along a period, the elements change gradually from metallic to non-metallic character; that is, metallic character decreases as one moves from left to right across a horizontal row.
Reading The Formulae At The Head Of A Group

Mendeleev published a revised version of his table, complete with his predictions, in a German journal in 1871. At the head of each group he wrote the general formula of the oxide and of the hydride of that group, using the letter \(\mathrm{R}\) to stand for any element of the group.

So the hydride of any element of Group IV is written \(\mathrm{RH_4}\); for carbon this becomes \(\mathrm{CH_4}\). The oxide of the same group is written \(\mathrm{RO_2}\), which for carbon is \(\mathrm{CO_2}\). The full table, with these general formulae along the top, is given in Table 5.3.

Group I II III IV V VI VII VIII
Oxide \(\mathrm{R_2O}\) \(\mathrm{RO}\) \(\mathrm{R_2O_3}\) \(\mathrm{RO_2}\) \(\mathrm{R_2O_5}\) \(\mathrm{RO_3}\) \(\mathrm{R_2O_7}\) \(\mathrm{RO_4}\)
Hydride \(\mathrm{RH}\) \(\mathrm{RH_2}\) \(\mathrm{RH_3}\) \(\mathrm{RH_4}\) \(\mathrm{RH_3}\) \(\mathrm{RH_2}\) \(\mathrm{RH}\) —
Period 1 H 1.008
Period 2 Li 6.939 Be 9.012 B 10.81 C 12.011 N 14.007 O 15.999 F 18.998
Period 3 Na 22.99 Mg 24.31 Al 26.98 Si 28.09 P 30.974 S 32.06 Cl 35.453
Period 4, first series K 39.102 Ca 40.08 Sc 44.96 Ti 47.90 V 50.94 Cr 52.00 Mn 54.94 Fe 55.85, Co 58.93, Ni 58.71
Period 4, second series Cu 63.54 Zn 65.37 Ga 69.72 Ge 72.59 As 74.92 Se 78.96 Br 79.909
Period 5, first series Rb 85.47 Sr 87.62 Y 88.91 Zr 91.22 Nb 92.91 Mo 95.94 Tc 99 Ru 101.07, Rh 102.91, Pd 106.4
Period 5, second series Ag 107.87 Cd 112.40 In 114.82 Sn 118.69 Sb 121.75 Te 127.60 I 126.90
Period 6, first series Cs 132.90 Ba 137.34 La 138.91 Hf 178.49 Ta 180.95 W 183.85 Os 190.2, Ir 192.2, Pt 195.2
Period 6, second series Au 196.97 Hg 200.59 Tl 204.37 Pb 207.19 Bi 208.98

Table 5.3 — Mendeleev’s periodic table, with the atomic mass written after each symbol. In a long period, the elements of the first series and those of the second series belong to different sub-groups of the same group

Example 5.2 — Predicting Formulae From A Group Number

1. Phosphorus is placed in Group V of Mendeleev’s table. Work out its valency, and hence the formula of its hydride and of its oxide.

  1. Phosphorus lies in Group V, so the rule for Groups V to VII applies: valency \(= 8 - 5 = 3\). A valency of 3 means one phosphorus atom combines with three hydrogen atoms, giving the hydride \(\mathrm{RH_3}\), that is \(\mathrm{PH_3}\). For the oxide, the group heading gives \(\mathrm{R_2O_5}\); putting phosphorus in place of \(\mathrm{R}\) gives \(\mathrm{P_2O_5}\).

  2. Phosphorus has valency 3, its hydride is \(\mathrm{PH_3}\) and its oxide is \(\mathrm{P_2O_5}\) — both of which are the formulae found by experiment.

5.4.3 Merits Of Mendeleev’s Table

Order Brought To Sixty-Three Elements

Mendeleev’s first achievement was simply one of scale. The study of 63 separate elements was reduced to the study of eight groups, because once you knew how a group behaved, you knew a good deal about every element in it.

Gaps Left For Elements Not Yet Discovered

To keep elements of similar properties in the same column, Mendeleev had to leave some squares of his table empty, as shown in Fig. 5.4-1. He did not treat these gaps as untidiness to be corrected. He argued that each gap stood for an element that had not yet been discovered, and would be found later.

Diagram

Fig. 5.4-1 — Gaps Left In Mendeleev’s Table A fragment of the table showing Groups II, III and IV as three columns and three periods as rows. Fill in Zn, Al and Si in their places; mark the square below aluminium and the square below silicon with a large question mark each, labelled ‘eka-aluminium’ and ‘eka-silicon’; add an arrow from each question mark to a caption box reading ‘square kept empty for an element not yet discovered’.

Why ‘Eka’? Mendeleev named his missing elements using Sanskrit number words. Eka means one, so eka-aluminium meant the element one place below aluminium in the group, and eka-silicon the element one place below silicon.

Properties Predicted Before Discovery

Leaving a gap was bold; describing what would fill it was bolder. In 1871 Mendeleev used the properties of the elements lying next to each vacant square — above, below and on either side — to predict the properties of the elements that would occupy them.

Eka-aluminium was discovered in 1875 and named gallium, the metal that opened this chapter. Eka-silicon proved to be germanium. Table 5.4 sets his predictions beside the values later measured.

Property Eka-aluminium Gallium Eka-silicon Germanium
Atomic mass 68 69.7 72 72.6
Density 5.9 g/cc 5.91 g/cc 5.5 g/cc 5.36 g/cc
Melting point Low 302 K High 1231 K
Valency 3 3 4 4
Formula of oxide \(\mathrm{M_2O_3}\) \(\mathrm{Ga_2O_3}\) \(\mathrm{MO_2}\) \(\mathrm{GeO_2}\)

Table 5.4 — Mendeleev’s predicted properties compared with the measured properties of the elements that filled the gaps

Q1. How could an empty square in a table be trusted enough to describe a metal that nobody had ever seen?

A. Because the square was not empty of information. Mendeleev knew the properties of the elements above, below and beside it, and he knew that properties change smoothly along a group and along a period. Reading the four neighbours of the gap told him roughly how heavy an atom of the missing element must be, how dense the metal, what valency it would show and what the formula of its oxide had to be. The gap described the element before anyone had it in a bottle.

Wrong Atomic Masses Put Right

The table could also be turned back on the measurements that built it. Where an element’s chemistry clearly placed it in one group but its accepted atomic mass placed it elsewhere, Mendeleev trusted the chemistry. The atomic masses of gold and platinum were corrected in this way, by placing these elements strictly according to the similarities in their properties.

5.4.4 Defects In Mendeleev’s Table

Anomalous Pairs

Mendeleev’s own rule was ‘increasing order of atomic mass’, yet in a few places he had to break it to keep chemically similar elements together. Three pairs did not obey his principle.

  1. Argon, with atomic mass 39.9, is placed before potassium, whose atomic mass is 39.1.
  2. Cobalt, with atomic mass 58.9, is placed before nickel, whose atomic mass is 58.6.
  3. Tellurium, with atomic mass 127.6, is placed before iodine, whose atomic mass is 126.9.

In each pair the heavier element comes first. Within Mendeleev’s law there was no way to justify this — the placings were right chemically and wrong by the rule.

Position Of Isotopes

Isotopes are atoms of the same element that have similar chemical properties but different atomic masses. Mendeleev’s law arranges elements by atomic mass, so two isotopes, having different atomic masses, ought to occupy two separate places in the table. They were never given separate places, and no reason within the law could be offered for that.

Chemically Dissimilar Elements Grouped Together

Copper and silver bear no resemblance to the alkali metals lithium and sodium, yet they were placed with them in the first group. Putting unlike elements in one group defeats the purpose of the arrangement.

Chemically Similar Elements Separated

The opposite fault also appears. Gold and platinum are chemically similar metals, but they were placed in separate groups.

Electron Arrangement Not Explained

The table showed that properties repeat, but it could not say why. Nothing in it explained the arrangement of electrons in an atom, and so nothing in it explained where the repetition came from.

Position Of Hydrogen

Hydrogen was never given a fixed position. It was placed in Group IA with the alkali metals, and also in Group VIIA with the halogens, because it behaves in both ways: it forms a positive ion, as in hydrogen chloride, \(\mathrm{HCl}\), and a negative ion, as in sodium hydride, \(\mathrm{NaH}\).

Silicon lies in Group IV of Mendeleev’s table. Using only the group headings, write the formula of its hydride and of its oxide, and state the valency of silicon.

5.5 Atomic Number As The Basis Of The Modern Periodic Law

Every defect listed at the end of the last section has the same root. Mendeleev was sorting the elements by a property of the whole atom’s mass, when the property that actually decides an element’s chemistry lies in its nucleus. Finding that property took another forty years.

5.5.1 Moseley And The Discovery Of Atomic Number

The size of the positive charge in the nucleus of an atom was measured by Henry Moseley, an English physicist.

In 1913 Moseley set up a discharge tube in which the anode could be changed from one metal to another, and let cathode rays strike the metal anode. When cathode rays hit an anode, rays of very short wavelength are given out, and Moseley found that their wavelength depends on which metal the anode is made of. His apparatus is shown in Fig. 5.5-1.

Diagram

Fig. 5.5-1 — Moseley’s Experiment With Different Anodes An evacuated discharge tube connected to a high-voltage supply: label the cathode at the left, the beam of cathode rays crossing the tube, and a replaceable metal anode at the right marked ‘anode metal changed one element at a time’. Show short-wavelength rays leaving the anode and passing to a crystal that spreads them into a spectrum recorded on a photographic plate; label the recorded lines and mark an arrow along the plate reading ‘wavelength decreases in a regular way as the anode metal moves along the periodic table’.

The result was strikingly regular. As he changed the anode metal element by element, in the order in which those elements stand in the periodic table, the wavelength of the rays decreased in a steady, regular manner — not in jumps, and not in the order of atomic mass.

Moseley concluded that the quantity changing regularly from one element to the next is the number of positive charges in the nucleus, and that this number is the most fundamental property of an atom.

NoteDefinition

Atomic Number The number of unit positive charges present in the nucleus of an atom of a particular element.

The positive charge of the nucleus comes from its protons, each of which carries one unit of positive charge. So the atomic number of an element is simply the number of protons in one atom of it, and it is written as a subscript before the symbol, as in \(^{35}_{17}\mathrm{Cl}\).

Atomic Number And Mass Number The atomic number counts protons only, and fixes which element the atom is. The mass number counts protons and neutrons together, and fixes how heavy that particular atom is. Change the atomic number and you have a different element; change only the mass number and you have the same element.

5.5.2 The Modern Periodic Law

Moseley had found what chemists had been looking for since the beginning of this chapter: a property of an element that never changes, and that increases by exactly one from each element to the next. Atomic number is a better fundamental property of elements than atomic mass, and the periodic law was restated on it.

NotePrinciple / Law

The Modern Periodic Law The physical and chemical properties of elements are a periodic function of their atomic numbers.

If the elements are arranged in order of increasing atomic number, those with similar properties are repeated after regular intervals. The periodicity in the table therefore runs on atomic numbers — that is, on the number of protons — and not on atomic masses.

Q. Two of Mendeleev’s rules stayed and one changed. Which changed?

A. The idea of periodicity stayed, and grouping like with like stayed. What changed is the property the elements are ordered by — atomic mass gave way to atomic number.

5.5.3 How The Anomalies Were Explained

Position Of Isotopes

All the isotopes of an element contain the same number of protons; only their numbers of neutrons differ. Since the number of protons is the atomic number, all the isotopes of an element have the same atomic number.

The modern law places elements by atomic number, so all the isotopes of an element belong at one and the same position in the table. What was an unanswered objection under Mendeleev’s law becomes an obvious consequence under the modern one.

Example 5.3 — Placing Two Isotopes

1. Carbon occurs as \(^{12}_{6}\mathrm{C}\) and as \(^{14}_{6}\mathrm{C}\). Decide how many places in the periodic table these two atoms need, giving your reasoning.

  1. Read the lower index of each symbol: both atoms have atomic number 6, so each nucleus holds 6 protons. The upper index differs, 12 against 14, so the two atoms differ in their number of neutrons — 6 in the first, 8 in the second. The modern periodic law arranges elements by atomic number, and the atomic number here is the same for both.

  2. Both atoms occupy the same single position in the periodic table, the position of carbon. Their difference in mass does not move either of them.

Position Of Anomalous Pairs

Take argon and potassium. The atomic number of argon is 18 and its mass number is 40, while the atomic number of potassium is 19 and its mass number is 39. Order these two by mass and potassium comes first; order them by atomic number and argon comes first.

Under the modern periodic law the elements are arranged by increasing atomic number. So argon, with the lower atomic number, comes first, and potassium, with the higher atomic number, comes later — which is exactly where the chemistry of the two elements demands they be put.

The positions of cobalt and nickel are justified in the same way. Cobalt has atomic number 27 and nickel 28, so cobalt comes first, even though its atomic mass is the greater of the two.

Q3. Argon is heavier than potassium, yet it is placed before potassium in today’s table. Why is this not treated as a mistake?

A. Because the table is no longer ordered by mass. Argon has 18 protons and potassium has 19, so argon’s atomic number is the smaller and argon must come first. Its atoms happen to be heavier because an ordinary argon atom carries more neutrons, and neutrons add mass without adding positive charge. Mass number is a fact about a particular atom; atomic number is what makes the element the element it is.

An atom of an element has 20 protons and 22 neutrons. State its atomic number and its mass number, and say whether an atom of the same element with 20 neutrons would be placed at a different position in the periodic table.

5.6 The Modern Periodic Table

The table built on atomic number is the one in use today. When the elements are arranged in order of increasing atomic number, those having an equal number of valence electrons turn up at regular intervals.

That last sentence is the whole engine of the table, so it is worth saying slowly. The valence electrons of an atom are the electrons in its outermost shell. Because the number of valence electrons repeats at regular intervals along the list, chemical properties repeat along with it.

The full modern table is given in Fig. 5.6-1. Keep it open beside you for the rest of the chapter.

Diagram

Fig. 5.6-1 — The Long Form Of The Periodic Table Of Elements The complete modern periodic table: 18 vertical columns numbered 1 to 18 across the top, with the older notation IA, IIA, IIIB to VIIB, VIII, IB, IIB, IIIA to VIIA and Zero printed beneath the new numbers; 7 horizontal rows numbered 1 to 7 down the left. Each box carries the atomic number, the symbol and the name of the element, from hydrogen (1) to oganesson (118). Mark the s-block (groups 1–2), the d-block (groups 3–12), the p-block (groups 13–18) and, below the main body, the f-block as two rows labelled Lanthanide series (57–71) and Actinide series (89–103), joined by arrows to their place in group 3 of periods 6 and 7. Include a colour key for active metals, heavy metals, metalloids, non-metals, halogens and inert gases, and a small key box showing how atomic number, symbol, name and physical state are printed in each cell.

5.6.1 The Cause Of Periodicity

Take the first group of the table as an example. Lithium, sodium, potassium, rubidium, caesium and francium have atomic numbers 3, 11, 19, 37, 55 and 87. They share a long list of properties.

  1. They are soft metals.
  2. They have a valency of one and form unipositive ions, that is, ions carrying a single positive charge.
  3. They are good reducing agents.
  4. They are very reactive, and are therefore found mainly in the combined state.
  5. They impart colour to a flame when they burn.
  6. They form hydrides with hydrogen.
  7. They form basic oxides with oxygen.
  8. They react with water to form metal hydroxides and hydrogen.

Now look at how the electrons of these atoms are arranged. The electronic configuration of an atom lists how many electrons occupy each shell, counting outwards from the nucleus: lithium is 2, 1; sodium is 2, 8, 1; potassium is 2, 8, 8, 1. Every one of these atoms has exactly one electron in its outermost shell, as Fig. 5.6-2 shows.

Diagram

Fig. 5.6-2 — One Valence Electron, Three Different Atoms Three shell diagrams side by side, each with a small labelled nucleus and circular shells carrying dots for electrons: lithium (2, 1), sodium (2, 8, 1) and potassium (2, 8, 8, 1). Circle the single outermost electron in each atom and connect the three circles with a horizontal bracket labelled ‘one valence electron in every case’; label the increasing number of shells below the three diagrams.

There is the answer. Elements with similar electronic configurations have similar properties, so the cause of periodicity is the recurrence of the same number of electrons in the outermost orbit — that is, the recurrence of a similar electronic configuration.

NoteDefinition

Periodic Properties And Periodicity Properties that reappear at regular intervals, or in which there is a gradual variation — an increase or a decrease at regular intervals — are called periodic properties. The phenomenon itself is known as the periodicity of elements.

Within a single group, the number of electrons in the outermost orbit stays the same as you go down, so the electronic configuration is similar and the valency stays the same as well. The elements of a group therefore have similar properties, even though the number of shells increases steadily down the group.

Q4. Lithium, sodium and potassium are three different metals. What could they possibly share that makes all three behave so alike in water?

A. One electron in the outermost shell. Lithium is 2, 1; sodium is 2, 8, 1; potassium is 2, 8, 8, 1. The three atoms differ in size and in the number of shells they carry, but each has a single valence electron that it loses easily, and it is that electron which does the chemistry. Losing it gives a unipositive ion in every case, so all three attack water in the same way, forming a hydroxide and releasing hydrogen.

5.6.2 Groups In The Modern Periodic Table

The vertical columns of the table are its groups, and the horizontal rows are its periods. Take the groups first.

International Union of Pure and Applied Chemistry (IUPAC) recommends that the groups be numbered from 1 to 18, replacing the older notation of Groups IA to VIIA, VIII, IB to VIIB and 0. Both notations are still met with, and Table 5.5 sets them side by side.

Old notation IA IIA IIIB IVB VB VIB VIIB VIII IB IIB IIIA IVA VA VIA VIIA 0
New notation 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

Table 5.5 — The older group notation compared with the IUPAC numbering. Group VIII of the old notation spreads over three of the new columns, 8, 9 and 10

Four Things To Know About The Groups
  1. The modern periodic table has eighteen vertical columns, made up of the groups I to VIII and the zero group. Groups I to VII are each divided into sub-groups A and B, which is how eight groups come to fill eighteen columns.
  2. The elements of the A sub-groups — IA, IIA, IIIA, IVA, VA, VIA and VIIA, that is, groups 1, 2, 13, 14, 15, 16 and 17 — are known as the main group elements or representative elements or the normal elements. Their outermost shells are incomplete.
  3. The elements of groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 are known as transition elements. They have their two outermost shells incomplete.
  4. The elements of the zero group, group 18, are called inert gases or noble gases. Owing to their stable electronic configurations they hardly react with other elements. They carry eight electrons in the outermost orbit, except helium, which has two.

5.6.3 Periods In The Modern Periodic Table

There are seven horizontal rows in the modern periodic table, each called a period. What decides an element’s period is the number of shells present in its atom: an atom with three occupied shells lies in period 3, and so on.

Along a period, electrons are added to the outermost shell one at a time until an octet — eight electrons — is reached. All the elements over which this filling happens belong to the same period, and the last element of each period therefore has eight electrons in its outermost shell. The one exception is the first period, which ends at helium with two.

  1. The first period contains only two elements, of atomic numbers 1 and 2. It is the shortest period.
  2. The second and third periods contain eight elements each — atomic numbers 3 to 10 in the second, and 11 to 18 in the third. These are the short periods.
  3. The fourth and fifth periods contain eighteen elements each — atomic numbers 19 to 36 in the fourth, and 37 to 54 in the fifth. These are the long periods.
  4. The sixth period contains 32 elements, of atomic numbers 55 to 86. It is the longest period.
  5. The seventh period also contains 32 elements, of atomic numbers 87 to 118, and is likewise a longest period.

Table 5.6 gathers these features together.

Period Type of period Number of elements Atomic numbers Number of shells First element Last element
1 Shortest period 2 1 – 2 1 H (1) He (2)
2 Short period 8 3 – 10 2 Li (3) Ne (10)
3 Short period 8 11 – 18 3 Na (11) Ar (18)
4 Long period 18 19 – 36 4 K (19) Kr (36)
5 Long period 18 37 – 54 5 Rb (37) Xe (54)
6 Longest period 32 55 – 86 6 Cs (55) Rn (86)
7 Longest period 32 87 – 118 7 Fr (87) Og (118)

Table 5.6 — Salient features of the periods of the modern periodic table

What A Group Tells You And What A Period Tells You The group number tells you how many valence electrons the atom has, and so how it will react. The period number tells you how many shells it has, and so how large the atom is. One fixes chemistry, the other fixes size.

Example 5.4 — Finding The Period From The Atomic Number

1. An element has atomic number 15. Write its electronic configuration, state the period it belongs to, and give its number of valence electrons.

  1. Atomic number 15 means 15 electrons in the neutral atom. Filling the shells outwards, the first takes 2 and the second takes 8, which accounts for 10; the remaining 5 go into the third shell. The configuration is 2, 8, 5. Three shells are occupied, and the period number equals the number of shells.

  2. The element has the configuration 2, 8, 5, lies in period 3, and has 5 valence electrons. Reading along period 3 in Table 5.6 and Fig. 5.6-1, the element is phosphorus, in group 15.

The Lanthanides And The Actinides

Two sets of elements need a word of their own. In group IIIB — group 3 — of the sixth period there is a set of elements with atomic numbers 57 to 71, beginning with lanthanum (La-57). These are known as the lanthanides, or the rare earths.

In group IIIB of the seventh period there is a matching set, with atomic numbers 89 to 103, beginning with actinium (Ac-89). These are the actinides, and all of them are radioactive.

The lanthanides have similar properties to one another because they all belong to the same group, and so do the actinides. They are printed in two separate rows at the foot of the table only because there are so many of them: setting all thirty into the main body would stretch the table out of shape.

An atom of an element has the electronic configuration 2, 8, 8. State the number of its valence electrons, the period to which it belongs, and whether you would expect it to react readily with other elements.

5.6.4 Bridge Elements And Typical Elements

Two smaller patterns run across the table rather than along it, and both have names.

The first is a diagonal one. An element of the second period often resembles, in its properties, the element of the next group in the third period. This is called a diagonal relationship, and such elements are called bridge elements. There are three pairs: lithium and magnesium; beryllium and aluminium; boron and silicon.

Group → 1 2 13 14
Period 2 Li Be B C
Period 3 Na Mg Al Si

Table 5.7 — The bridge elements, each pair lying diagonally: Li with Mg, Be with Al, and B with Si

The second is a matter of being representative. The third-period elements — sodium, magnesium, aluminium, silicon, phosphorus, sulphur and chlorine — summarise the properties of their respective groups particularly well, and are for that reason called typical elements. If you want to know how group 16 behaves in general, sulphur is the member to look at.

5.7 Types Of Elements

Groups and periods tell you where an element sits. There is a second, coarser way of sorting the same elements, which cuts the table into four broad types:

  1. Representative elements
  2. Transition elements
  3. Inner transition elements
  4. Inert gases or noble gases

Each type occupies its own region of the table, as Fig. 5.7-1 shows. A further classification into s, p, d and f blocks is based on electronic configuration, and you will study it in a higher class.

Diagram

Fig. 5.7-1 — The Four Types Of Elements On The Table A plain outline of the periodic table with no element symbols, only the group numbers 1 to 18 across the top and the period numbers 1 to 7 down the side. Shade and label four regions distinctly: groups 1 and 2 together with groups 13 to 17 as ‘representative elements’; groups 3 to 12 as ‘transition elements’; group 18 as ‘inert gases’; and the two detached rows below the main body as ‘inner transition elements’.

5.7.1 Representative Elements

These are the s-block and p-block elements: the elements whose outermost orbit is incomplete. They include every element of the following seven groups.

  • Group 1: the alkali metals, which form strong alkalis with water. Hydrogen, which heads the column, is not one of them.
  • Group 2: the alkaline earth metals, which form alkalis that are weaker than those of the group 1 elements.
  • Group 13: the boron family, boron being the first member of the group.
  • Group 14: the carbon family, carbon being the first member.
  • Group 15: the nitrogen family, nitrogen being the first member.
  • Group 16: the oxygen family, oxygen being the first member.
  • Group 17: the halogens, which are salt formers.

The two ends of this set are the two most reactive families in the table. The alkali metals are the most reactive metals that occur; they react vigorously with water to produce hydrogen and an alkali solution, and it is this that gives them their name. The halogens are the most reactive non-metals, and form salts on reacting with metals.

Main Characteristics Of Representative Elements
  • They include both metals and non-metals, with a regular gradation from metallic to non-metallic character as one moves from left to right across a period.
  • They form electrovalent as well as covalent compounds with non-metals.
  • Metallic nature increases on moving down any of these seven groups.
  • The metals, which are good conductors of heat and electricity, are found in groups 1 and 2. The non-metals of groups 16 and 17 are poor conductors of both.
  • Some of the heavier members, such as tin and lead, exhibit variable valencies — that is, the same element can show more than one combining capacity.

5.7.2 Transition Elements

The transition elements are the d-block elements, and they occupy groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 — the wide block in the middle of the table.

Main Characteristics Of Transition Elements
  • All of them are metals, with high melting points and boiling points.
  • They are good conductors of heat and electricity.
  • Some of them are attracted towards a magnet.
  • Most of them are used as catalysts. A catalyst is a substance that speeds up a chemical reaction without being used up in it.
  • Most of them exhibit variable valencies.
  • Most of them form coloured ions and coloured compounds.

The Metals Inside A Catalytic Converter The exhaust of a petrol vehicle carries unburnt fuel and poisonous carbon monoxide. Fitted into the exhaust pipe is a honeycomb of ceramic coated with a thin film of platinum, palladium and rhodium — three transition metals. On their surface the harmful gases are converted to carbon dioxide, water vapour and nitrogen before leaving the pipe. The metals are not consumed, which is why so little of them is needed and why the same converter lasts the life of the vehicle.

Representative And Transition Elements A representative element has only its outermost shell incomplete. A transition element has its two outermost shells incomplete. That single difference is why the transition metals show variable valencies and coloured compounds, and the representative elements usually do not.

5.7.3 Inner Transition Elements

The elements of the sixth and seventh periods of group 3 — that is, the lanthanides and the actinides — are together known as the inner transition elements, or the f-block elements.

Main Characteristics Of Inner Transition Elements
  • They are heavy metals with high melting points and boiling points.
  • They show variable valencies.
  • They form coloured ions.
  • The actinides are all radioactive. Some of them are not found in nature at all, and are made in specially equipped laboratories.

5.7.4 Inert Gases Or Noble Gases

The elements of the zero group, which is the eighteenth vertical column, are known as inert gases or noble gases.

They have 8 electrons in their outermost orbits, except helium, which has only 2. Because these arrangements are already stable, the atoms have nothing to gain by combining, and so these elements do not react with other elements. This is the reason they are called inert. Their properties are taken up in detail in section 5.11.4.

Example 5.5 — Sorting Three Elements By Type

1. Elements with atomic numbers 12, 26 and 54 are given. Using Table 5.6 and Fig. 5.6-1, state which type each belongs to.

  1. Atomic number 12 falls in the range 11–18, so the element lies in period 3, in group 2; group 2 is one of the seven representative groups. Atomic number 26 falls in the range 19–36, so the element lies in period 4, in group 8; groups 3 to 12 hold the transition elements. Atomic number 54 falls in the range 37–54, and 54 is the last element of period 5, so it lies in group 18.

  2. Element 12 is a representative element, element 26 is a transition element, and element 54 is an inert gas.

An element is a hard metal, forms a green compound, and is used as a catalyst in industry. State which of the four types of elements it most probably belongs to, and give two further properties you would expect it to show.

5.8 Merits Of The Modern Periodic Table

Set beside the table Mendeleev built, the modern table is an improvement in four ways.

  1. It is based on atomic number, which is an even better fundamental property of an element than atomic mass. Atomic number cannot vary from sample to sample, because changing it would change the element itself.
  2. The position of an element in the table is related to its electronic configuration. Its group tells you the number of valence electrons and its period the number of shells, so the position and the atom’s structure carry the same information.
  3. It shows regular changes in the properties of elements on moving across a period or down a group. These regular changes are set out in section 5.10.
  4. It is easier to remember, to understand and to reproduce, because every one of its features follows from the way electrons are arranged rather than from a list of exceptions.

The second point deserves the most weight. Mendeleev could show that properties repeat; the modern table also shows why they repeat, and it does so from the structure of the atom. Table 5.8 sets the two tables side by side.

Table 5.8 — Mendeleev’s Periodic Table Compared With The Modern Periodic Table | Basis of comparison | Mendeleev’s periodic table | Modern periodic table | |—|—|—| | Property used for arrangement | Increasing atomic mass | Increasing atomic number | | Number of groups | Eight, I to VIII, with sub-groups A and B | Eighteen, numbered 1 to 18 | | Number of periods | Seven, with long periods split into two series | Seven | | Position of isotopes | Not explained; different atomic masses demand separate places | All isotopes share one place, since they share one atomic number | | Anomalous pairs | Placed against the rule, with no justification | Justified, since the order of atomic number differs from that of atomic mass | | Link with atomic structure | None; electron arrangement not explained | Position follows directly from the electronic configuration | | Noble gases | No place kept for them; they were unknown | Placed together in group 18 |

State one merit of the modern periodic table that Mendeleev’s table could not claim, and explain in a sentence what makes it possible.

5.9 Defects Of The Modern Periodic Table

No classification is perfect, and two difficulties survive even in the modern table.

  1. The position of hydrogen is still not satisfactory. Its properties relate to group 1 on one side and to group 17 on the other, so wherever it is printed, part of its chemistry is being ignored. This is taken up in section 5.10.3.
  2. The table fails to accommodate the inner transition elements — the lanthanides and the actinides — within its main body. Thirty elements have to be lifted out and printed separately below, which means the table is not quite the single continuous arrangement it claims to be.

Notice that neither defect is of the kind that troubled Mendeleev. Nothing here places an element in the wrong company or leaves a regularity unexplained; the difficulties are about how the arrangement is drawn on paper, and about one element that genuinely belongs to two families at once.

Q. Why do the lanthanides and actinides have to be shown below the table at all?

A. They all belong to group 3, and a group can hold only one element per period. Fitting fifteen elements into a single square is impossible, and spreading them sideways would stretch the sixth and seventh periods far past the others.

5.11 Study Of Specific Groups

Four groups of the table are worth meeting properly, because between them they cover the extremes of chemical behaviour: the two groups on the far left that give electrons away most readily, the group on the far right but one that takes them most greedily, and the group at the very edge that does neither.

5.11.1 Group 1 — The Alkali Metals

Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr) have one electron in their outermost orbit and therefore show a valency of one. They are placed in group IA, the first column on the left of the periodic table.

They are called alkali metals because they react with water to form their hydroxides, and these hydroxides are strong alkalis — that is, bases that dissolve in water.

alkali metal + water → alkali metal hydroxide + hydrogen

\[2\,\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{MOH} + \mathrm{H_2}\]

where \(\mathrm{M}\) stands for any metal of the IA group.

Characteristics Of The Alkali Metals
  1. Electronic configuration. Every one of them ends in a single outermost electron.
Element K L M N O P Q
\(_{3}\mathrm{Li}\) 2 1
\(_{11}\mathrm{Na}\) 2 8 1
\(_{19}\mathrm{K}\) 2 8 8 1
\(_{37}\mathrm{Rb}\) 2 8 18 8 1
\(_{55}\mathrm{Cs}\) 2 8 18 18 8 1
\(_{87}\mathrm{Fr}\) 2 8 18 32 18 8 1

Table 5.17 — Electronic configurations of the alkali metals, shell by shell

  1. Occurrence. All of them are very reactive and do not occur in the free state.

  2. Nature. They are soft — sodium can be cut with a knife — and have low melting and boiling points.

  3. Bonding. All alkali metal salts, except some salts of lithium, are ionic in nature.

  4. Action of air. They react rapidly with the oxygen and the water vapour of the air. Sodium shows the sequence clearly. First the metal is oxidised:

sodium + oxygen → sodium oxide

\[4\,\mathrm{Na} + \mathrm{O_2} \rightarrow 2\,\mathrm{Na_2O}\]

The oxide then takes up moisture to give the hydroxide:

\[\mathrm{Na_2O} + \mathrm{H_2O} \rightarrow 2\,\mathrm{NaOH}\]

and the hydroxide absorbs carbon dioxide from the air to give the carbonate:

\[2\,\mathrm{NaOH} + \mathrm{CO_2} \rightarrow \mathrm{Na_2CO_3} + \mathrm{H_2O}\]

The reactivity of these metals increases down the group.

Why Sodium Lives Under Kerosene Because they attack both oxygen and water vapour, the group 1 metals cannot be left in an open jar; a lump of sodium exposed to air is losing itself to the three reactions above even as you watch. They are therefore stored under an inert solvent such as kerosene, which keeps air and moisture off the surface. This is also why a piece is always cut on a dry tile with dry instruments, never handled with wet fingers, and why the offcuts go back under the kerosene rather than into a sink.

  1. Action of water. They react violently with water and produce hydrogen.

\[2\,\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{MOH} + \mathrm{H_2}\]

  1. Action of acids. They react violently with dilute hydrochloric acid and dilute sulphuric acid, producing hydrogen.

alkali metal + dilute hydrochloric acid → alkali metal chloride + hydrogen

\[2\,\mathrm{M} + 2\,\mathrm{HCl} \rightarrow 2\,\mathrm{MCl} + \mathrm{H_2}\]

  1. Formation of ions. They form unipositive ions easily, and are therefore strong reducing agents — a reducing agent is a substance that gives up electrons to something else.

\[\mathrm{M} - e^{-} \rightarrow \mathrm{M^{+}}\]

  1. Appearance. When freshly cut they show a typical silvery white metallic lustre, but the shine soon dulls, or tarnishes, because of their reaction with the air.

  2. Flame colour. All alkali metal salts impart colour to a flame: lithium a crimson red, sodium a golden yellow and potassium a lilac.

Diagram

Fig. 5.11-1 — Flame Colours Of Some Metal Salts Five identical burner flames drawn side by side, each with a nichrome wire loop dipped in a salt held in it. Colour and label the flames: lithium salt — crimson red; sodium salt — golden yellow; potassium salt — lilac; calcium salt — brick red; barium salt — apple green. Below the row, label the loop, the salt and the hottest part of the flame.

  1. Extraction. Alkali metals are obtained by the electrolysis of their molten salts.

Example 5.8 — Using A General Equation

1. Write the balanced equation for the reaction of potassium with water, starting from the general equation for a group 1 metal, and name the alkali formed.

  1. The general equation is \(2\,\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{MOH} + \mathrm{H_2}\). Potassium belongs to group 1, so \(\mathrm{M}\) is replaced throughout by \(\mathrm{K}\), keeping every coefficient as it stands. In words: potassium + water → potassium hydroxide + hydrogen.

  2. \[2\,\mathrm{K} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{KOH} + \mathrm{H_2}\] The alkali formed is potassium hydroxide. Counting atoms on both sides — 2 K, 4 H and 2 O each way — confirms that the equation is balanced.

5.11.2 Group 2 — The Alkaline Earth Metals

Beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra) are known as the alkaline earth metals. The name has two halves to it: their oxides occur in the earth’s crust, and their hydroxides are weak alkalis.

All the elements of this group have two electrons in their outermost orbit and show a valency of two, so they are placed in the IIA group of the periodic table.

Characteristics Of The Alkaline Earth Metals
  1. Electronic configuration. Every one of them ends in two outermost electrons.
Element K L M N O P Q
\(_{4}\mathrm{Be}\) 2 2
\(_{12}\mathrm{Mg}\) 2 8 2
\(_{20}\mathrm{Ca}\) 2 8 8 2
\(_{38}\mathrm{Sr}\) 2 8 18 8 2
\(_{56}\mathrm{Ba}\) 2 8 18 18 8 2
\(_{88}\mathrm{Ra}\) 2 8 18 32 18 8 2

Table 5.18 — Electronic configurations of the alkaline earth metals, shell by shell

  1. Occurrence. They are reactive metals, though less reactive than the alkali metals, and so they too do not occur in the free state in nature.

  2. Nature. They are greyish white in colour and harder than the alkali metals.

  3. Bonding. All the alkaline earth metal salts, except the salts of beryllium, are ionic compounds.

  4. Action of air. They are less reactive towards air than the alkali metals. Their reactivity increases on moving down the group.

  5. Action of water. They react with water to produce hydrogen.

alkaline earth metal + water → alkaline earth metal hydroxide + hydrogen

\[\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow \mathrm{M(OH)_2} + \mathrm{H_2}\]

  1. Action of acids. They react with dilute hydrochloric acid and dilute sulphuric acid to produce hydrogen.

\[\mathrm{M} + 2\,\mathrm{HCl} \rightarrow \mathrm{MCl_2} + \mathrm{H_2}\]

\[\mathrm{M} + \mathrm{H_2SO_4} \rightarrow \mathrm{MSO_4} + \mathrm{H_2}\]

  1. Formation of ions. They form dipositive ions, and are also strong reducing agents.

\[\mathrm{M} - 2e^{-} \rightarrow \mathrm{M^{2+}}\]

  1. Melting and boiling points. These are comparatively low, but higher than those of the alkali metals — the dipositive ions hold the metal together more firmly than unipositive ones do.

  2. Flame colour. Except beryllium and magnesium, all the alkaline earth metals impart colour to a flame: calcium a brick red, strontium a crimson, barium an apple green and radium a crimson.

  3. Extraction. They too are obtained by the electrolysis of their molten salts.

Table 5.19 — Alkali Metals Compared With Alkaline Earth Metals | Basis of comparison | Alkali metals (group 1) | Alkaline earth metals (group 2) | |—|—|—| | Valence electrons | One | Two | | Valency | 1 | 2 | | Ion formed | Unipositive, \(\mathrm{M^{+}}\) | Dipositive, \(\mathrm{M^{2+}}\) | | Hardness | Soft; sodium can be cut with a knife | Harder than the alkali metals | | Melting and boiling points | Low | Low, but higher than those of the alkali metals | | Reactivity | Very reactive; react violently with water | Reactive, but less so than the alkali metals | | Nature of hydroxides | Strong alkalis | Weak alkalis | | Bonding in their salts | Ionic, except some salts of lithium | Ionic, except the salts of beryllium |

Alkali Metals And Alkaline Earth Metals Alkali metals are the group 1 elements, with one valence electron, whose hydroxides are strong alkalis. Alkaline earth metals are the group 2 elements, with two valence electrons, whose oxides are found in the earth’s crust and whose hydroxides are weak alkalis.

Hold a clean, dry steel spoon carrying a few grains of common salt at the edge of a kitchen gas flame, with an adult present, and watch the colour that appears. Predict the colour first from the flame-colour list above, then check whether your prediction was right.

5.11.3 Group 17 — The Halogens

The elements of group VIIA, or group 17, are known as the halogens. The name comes from Greek: halo, meaning sea or salt, and gen, meaning producing. A halogen is a salt-former.

Characteristics Of The Halogens
  1. Electronic configuration. These elements have 7 electrons in their outermost orbit.
Element K L M N O P Physical state and colour
\(_{9}\mathrm{F}\) 2 7 Pale yellow gas
\(_{17}\mathrm{Cl}\) 2 8 7 Greenish yellow gas
\(_{35}\mathrm{Br}\) 2 8 18 7 Reddish brown liquid
\(_{53}\mathrm{I}\) 2 8 18 18 7 Violet solid
\(_{85}\mathrm{At}\) 2 8 18 32 18 7 Highly radioactive and rare

Table 5.20 — Electronic configurations, physical states and colours of the halogens

  1. Reactivity. They are the most reactive non-metals. Their reactivity decreases down the group: fluorine is the most reactive halogen and iodine the least reactive.

  2. Melting point and boiling point. There is a steady increase in melting points and boiling points as we go down the group. This is why fluorine is a gas, chlorine is a gas, bromine is a liquid and iodine is a solid at ordinary temperature. Astatine is highly radioactive and a rare element.

  3. Colour. The intensity of the colour of the element also increases from pale to dark down the group, as Table 5.20 shows and Fig. 5.11-2 illustrates.

Diagram

Fig. 5.11-2 — The Halogens Darken Down The Group Four stoppered glass jars in a row, drawn left to right and labelled with the element name and physical state: fluorine as a pale yellow gas, chlorine as a greenish yellow gas, bromine as a reddish brown liquid with a dark vapour above it, and iodine as violet-black solid crystals with a faint violet vapour. Draw a rightward arrow beneath the row labelled ‘colour deepens and melting point rises down the group’.

  1. Nature. They are all poisonous and have a similar strong, pungent and unpleasant odour. For this reason they are handled only in a fume cupboard, and never smelt directly.

  2. Atomicity. They all form diatomic molecules: \(\mathrm{F_2}\), \(\mathrm{Cl_2}\), \(\mathrm{Br_2}\) and \(\mathrm{I_2}\).

  3. Valency. They all have a valency of 1, and so form compounds with similar formulae — for example hydrogen chloride (\(\mathrm{HCl}\)), hydrogen bromide (\(\mathrm{HBr}\)) and hydrogen iodide (\(\mathrm{HI}\)).

  4. Halides. They produce a whole series of compounds with other elements — chlorides, bromides and iodides — which together are known as halides.

  5. Reaction with metals. The halogens can react directly with metals to form metal halides, that is, salts. This is what the name of the group records.

  6. Ions formed. They all form negative ions carrying a single charge: the fluoride ion \(\mathrm{F^{-}}\), the chloride ion \(\mathrm{Cl^{-}}\), the bromide ion \(\mathrm{Br^{-}}\) and the iodide ion \(\mathrm{I^{-}}\).

5.11.4 Group 18 — The Noble Gases

The elements of the zero group, the eighteenth vertical column, are the inert gases or noble gases: helium, neon, argon, krypton, xenon and radon.

A Whole Group That Nobody Missed When Mendeleev first constructed his table there was no indication that an entire group of elements remained to be discovered, precisely because these elements react with nothing and so leave no chemical trace of themselves. The first clue came from a careful analysis of the gases of the air, which led to argon. Helium had already been spotted differently — in the spectrum of sunlight studied during an eclipse, before anyone had found it on earth. William Ramsay isolated these gases, and was awarded the Nobel Prize for the work.

The name has changed once. These gases were called inert gases, meaning that they do not react at all. Then, in the 1960s, some compounds of xenon and krypton were made, and the name was changed to noble gases — noble in the sense of aloof rather than incapable.

Characteristics Of The Noble Gases
  1. Electronic configuration. Helium has 2 electrons in its only shell, and all the other members have eight electrons in their outermost orbit. Their electronic arrangement is very stable, and that is why they are unreactive.
Element K L M N O P
\(_{2}\mathrm{He}\) 2
\(_{10}\mathrm{Ne}\) 2 8
\(_{18}\mathrm{Ar}\) 2 8 8
\(_{36}\mathrm{Kr}\) 2 8 18 8
\(_{54}\mathrm{Xe}\) 2 8 18 18 8
\(_{86}\mathrm{Rn}\) 2 8 18 32 18 8

Table 5.21 — Electronic configurations of the noble gases, shell by shell

2. All are colourless, tasteless and odourless gases. They are neither inflammable nor supporters of combustion.

3. They are monoatomic — their atoms go about singly — because their electronic configuration is already stable and there is nothing to be gained by pairing up.

4. They are liquefied only with great difficulty.

5. They are slightly soluble in water, and on moving down the group from helium to radon their solubility increases.

6. They emit coloured light when an electrical discharge is passed through them.

7. Their melting points and boiling points are extremely low. Helium has the lowest melting point of any element, and cannot be solidified by cooling alone — pressure is needed as well. All these properties point to atoms that are particularly stable and hold on to nothing.

  1. Uses. Every use of a noble gas depends on its unreactive nature.
  • Helium is used in airships and balloons, because it is both light and unreactive.
  • Argon is used to fill light bulbs, because it will not react with the glowing filament even at high temperatures.
  • The best known use is in ‘neon’ lights. The brightly coloured advertising sign works when an electric discharge is passed through a tube containing a little of a noble gas, and different noble gases give different colours.
Why Their Neighbours Are So Reactive

The elements of group 18 sit between the two most reactive groups of elements — group 1 on one side and group 17 on the other. It is precisely this closeness to a stable electron arrangement that makes the alkali metals and the halogens so reactive.

A group 17 atom is one electron short of the noble gas arrangement, so it gains or shares an electron and reaches it. A group 1 atom has one electron too many, so it loses that electron and reaches it. Both are within a single step of stability, and both take that step at the first opportunity.

An element is a gas that glows red-orange in a discharge tube, does not burn, and forms no compounds under ordinary conditions. State the group it belongs to, give its number of valence electrons, and explain why it does not react with sodium.

5.12 Uses Of The Periodic Table

The table began as a filing system for what chemists already knew. It turned out to be far more than that.

  1. It has been useful in predicting the existence of new elements, as the gaps left by Mendeleev showed.
  2. It has been useful in the past in correcting the position of elements in relation to their properties.
  3. It has made the study of the elements and their compounds systematic, and much easier to remember, because a whole group can be learnt at once.
  4. The position of an element in the table reveals its atomic number, its electronic configuration, its number of valence electrons and hence its properties.
Predicting A Compound Before Making It

The fifth use is the one a student will lean on most. From the positions of two elements you can work out the nature of the chemical bond between them, the formula of the compound they form, and something of that compound’s properties.

The elements of groups 1, 2 and 3 are metals, and they lose electrons to form positive ions. The elements of groups 15, 16 and 17 are non-metals, and they gain electrons to form negative ions. When such a positive ion and such a negative ion combine, the compound formed is an electrovalent compound — one held together by the attraction between oppositely charged ions.

If instead the elements of groups 14, 15, 16 and 17 combine with those of groups 15, 16 and 17 — that is, non-metals with non-metals — they form covalent compounds, in which the atoms share electrons rather than transferring them.

Finally, the position of an element also reveals its valency, and whether it is a metal or a non-metal: metals occupy the extreme left positions of the periodic table, while non-metals are at the extreme right.

The whole chain of reasoning, from a position on the table to the formula of a compound, is set out in FC 5.1.

Flowchart

FC 5.1 — From A Position In The Table To The Formula Of A Compound Position of the element (group and period) → number of valence electrons, read from the group number → valency, and whether electrons are lost or gained → metal on the left, non-metal on the right → metal with non-metal gives an electrovalent compound; non-metal with non-metal gives a covalent compound → formula of the compound, obtained from the two valencies

Example 5.9 — Predicting A Compound From Two Positions

1. Element A lies in group 2 and element B in group 17, both in period 3. Predict the type of bond between them and the formula of the compound they form.

  1. A is in group 2, so it has 2 valence electrons and a valency of 2; lying on the left of the table it is a metal, and it loses both electrons to give \(\mathrm{A^{2+}}\). B is in group 17, so it has 7 valence electrons and a valency of \(18 - 17 = 1\); lying on the right it is a non-metal, and it gains one electron to give \(\mathrm{B^{-}}\). A metal combining with a non-metal by transfer of electrons gives an electrovalent compound. One \(\mathrm{A^{2+}}\) ion needs two \(\mathrm{B^{-}}\) ions for the charges to cancel.

  2. The compound is electrovalent, and its formula is \(\mathrm{AB_2}\). Reading period 3 in Fig. 5.6-1, A is magnesium and B is chlorine, so the compound is magnesium chloride, \(\mathrm{MgCl_2}\).

The Table As A Shopping List Industrial chemists searching for a new catalyst, magnet or battery material rarely start at random. They start from a material that already works and try its neighbour, one square along or one square down, because that element is the one most likely to behave in a similar way.

Element X lies in group 14 and element Y in group 16. State whether the compound they form is electrovalent or covalent, give a reason, and write the formula of the compound.

5.13 Summary

  1. Elements are classified so that they can be studied in an organised way, so that their properties can be linked with the properties of matter, and so that the relationship of one element to another can be stated.

  2. The earliest attempts used density, malleability, ductility and the division into metals and non-metals. They failed because groups were too large, the properties used varied with conditions, and some elements showed both metallic and non-metallic character.

  3. Attention then turned to atomic mass, which was thought to be a fixed and fundamental property of an element.

  4. Döbereiner’s law of triads: when three chemically similar elements are arranged in order of increasing atomic mass, the atomic mass of the middle element is approximately the arithmetic mean of the other two, \(B \approx \dfrac{A + C}{2}\).

  5. Calcium (40), strontium (88) and barium (137) form a triad, since \(\dfrac{40 + 137}{2} = 88.5\); so do chlorine (35.5), bromine (80) and iodine (127), since \(\dfrac{35.5 + 127}{2} = 81.25\).

  6. The law of triads was discarded because most known elements could not be arranged in triads, and because it failed even within a family — for the halogens, \(\dfrac{1}{2}(19 + 80) = 49.5\), not 35.5.

  7. Newland’s law of octaves: when the elements are arranged in increasing order of atomic mass, the properties of every eighth element, counting from any element, repeat the properties of the starting element.

  8. Its merits were that it worked well for the lighter elements, that it linked properties to atomic masses, and that it showed for the first time a distinct periodicity in the properties of elements.

  9. It was discarded because it failed beyond calcium, because newly discovered elements could not be fitted in, because cobalt and nickel were forced into one slot in the halogen column, and because iron was separated from cobalt and nickel.

  10. Mendeleev arranged the 63 elements then known in increasing order of atomic mass, placing similar elements in the same vertical column and leaving gaps where the periodicity was disrupted.

  11. Mendeleev’s periodic law: the physical and chemical properties of elements are a periodic function of their atomic weights.

  12. His table had eight groups, I to VIII, with groups I to VII divided into sub-groups A and B; the valency of an element equals its group number for groups I to IV, and eight minus the group number for groups V to VII. The general oxide and hydride formulae head each group, as \(\mathrm{RO_2}\) and \(\mathrm{RH_4}\) for group IV.

  13. Merits of Mendeleev’s table: it reduced 63 elements to eight groups, it left gaps for undiscovered elements, it predicted the properties of eka-aluminium (gallium) and eka-silicon (germanium), and it corrected the atomic masses of gold and platinum.

  14. Defects of Mendeleev’s table: anomalous pairs (argon before potassium, cobalt before nickel, tellurium before iodine), no place for isotopes, chemically dissimilar elements grouped together (copper and silver with the alkali metals), chemically similar elements separated (gold and platinum), no explanation of electron arrangement, and no fixed position for hydrogen.

  15. In 1913 Moseley showed that the wavelength of the rays produced at a metal anode changes regularly with the element, and concluded that the positive charge of the nucleus — the atomic number — is the most fundamental property of an atom.

  16. The atomic number is the number of unit positive charges, and hence the number of protons, in the nucleus of an atom.

  17. The modern periodic law: the physical and chemical properties of elements are a periodic function of their atomic numbers.

  18. The modern law explains the anomalies: all the isotopes of an element share one atomic number and so one position, and argon (18) rightly precedes potassium (19), as cobalt (27) precedes nickel (28).

  19. The cause of periodicity is the recurrence of the same number of electrons in the outermost orbit, that is, of a similar electronic configuration. Properties that reappear at regular intervals, or vary gradually at regular intervals, are periodic properties.

  20. The modern table has eighteen vertical columns, numbered 1 to 18 by IUPAC, and seven horizontal periods. The number of shells in an atom fixes its period; the number of valence electrons fixes its group.

  21. Period 1 holds 2 elements, periods 2 and 3 hold 8 each, periods 4 and 5 hold 18 each, and periods 6 and 7 hold 32 each. The lanthanides (57–71) and the actinides (89–103) belong to group 3 of periods 6 and 7 and are printed below the table.

  22. The elements are of four types: representative elements (outermost shell incomplete), transition elements of groups 3 to 12 (two outermost shells incomplete), inner transition elements (the lanthanides and actinides) and the inert gases of group 18.

  23. Second-period elements resemble the third-period element of the next group; these diagonally related pairs — Li and Mg, Be and Al, B and Si — are the bridge elements. The third-period elements Na, Mg, Al, Si, P, S and Cl are the typical elements.

  24. Merits of the modern table: it is based on atomic number, position is related to electronic configuration, properties change regularly across a period and down a group, and the whole arrangement is easier to remember and reproduce. Its defects are the position of hydrogen and its failure to accommodate the inner transition elements in the main body.

  25. Down a group: the number of shells increases, the number of valence electrons and the valency stay the same, atomic size increases, metallic character increases and the oxides become more basic. Reactivity increases down group 1, \(\mathrm{Li} < \mathrm{Na} < \mathrm{K} < \mathrm{Rb} < \mathrm{Cs}\), and decreases down group 17, \(\mathrm{F} > \mathrm{Cl} > \mathrm{Br} > \mathrm{I}\).

  26. Across a period: the number of shells stays the same, valence electrons increase one at a time, atomic size decreases because the nuclear charge rises, metallic character decreases, and the oxides change from basic to acidic. Valency with respect to hydrogen rises from 1 to 4 and falls back to 1; valency with respect to oxygen rises from 1 to 7.

  27. Hydrogen is placed above group 1 because it has one valence electron, but its atom is very small and many of its properties differ from those of the alkali metals; it is therefore left out when the alkali metals are discussed.

  28. The alkali metals of group 1 have one valence electron, are soft, very reactive and stored under kerosene, form unipositive ions, are strong reducing agents, colour a flame, and are obtained by the electrolysis of their molten salts.

\[2\,\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow 2\,\mathrm{MOH} + \mathrm{H_2}\]

  1. The alkaline earth metals of group 2 have two valence electrons, are harder and less reactive than the alkali metals, form dipositive ions, and give weak alkalis.

\[\mathrm{M} + 2\,\mathrm{H_2O} \rightarrow \mathrm{M(OH)_2} + \mathrm{H_2}\]

  1. The halogens of group 17 have seven valence electrons, are the most reactive non-metals with reactivity decreasing down the group, are diatomic and poisonous, deepen in colour and rise in melting point down the group, and form singly charged negative ions and metal halides.

  2. The noble gases of group 18 have eight valence electrons (helium two), are colourless, odourless, monoatomic and unreactive, have extremely low melting and boiling points, and are used in balloons, light bulbs and discharge tubes. It is their stable arrangement that the alkali metals and halogens are reaching for when they react.

  3. The periodic table predicts new elements, corrects positions, and reveals an element’s atomic number, electronic configuration, valency, metallic or non-metallic nature, the type of bond it will form and the formula of its compounds.

Flowchart

FC 5.2 — The Road To The Modern Periodic Table Elements classified by density, malleability and metal or non-metal character → attempts fail, as the properties vary → atomic mass taken as a fixed property → Döbereiner’s triads, covering only a few elements → Newland’s octaves, failing beyond calcium → Mendeleev’s table, ordered by atomic mass, with gaps left and properties predicted → anomalous pairs and isotopes remain unexplained → Moseley finds the atomic number in 1913 → modern periodic law, ordered by atomic number → modern periodic table of 18 groups and 7 periods, explained by electronic configuration

5.14 Key Terms

  1. Actinides — the fifteen elements of atomic numbers 89 to 103, belonging to group 3 of the seventh period, all of them radioactive.
  2. Alkali metals — the elements of group 1, each with one valence electron, whose hydroxides are strong alkalis.
  3. Alkaline earth metals — the elements of group 2, each with two valence electrons, whose oxides occur in the earth’s crust and whose hydroxides are weak alkalis.
  4. Atomic mass — the mass of an atom of an element compared with a standard atom, carrying no unit.
  5. Atomic number — the number of unit positive charges, and hence of protons, in the nucleus of an atom of an element.
  6. Bridge elements — the diagonally related pairs Li and Mg, Be and Al, and B and Si, in which a second-period element resembles the third-period element of the next group.
  7. Catalyst — a substance that speeds up a chemical reaction without being used up in it.
  8. Covalent compound — a compound in which the atoms are held together by sharing electrons.
  9. Electronic configuration — the arrangement of the electrons of an atom in its shells, counted outwards from the nucleus.
  10. Electrovalent compound — a compound held together by the attraction between oppositely charged ions formed by the transfer of electrons.
  11. Element — a pure substance made of only one kind of atom, which cannot be broken down into anything simpler by a chemical change.
  12. Groups — the vertical columns of the periodic table, whose elements have the same number of valence electrons.
  13. Halogens — the elements of group 17, the most reactive non-metals, whose name means salt-former.
  14. Inert gases — the elements of group 18, whose stable electronic arrangement makes them virtually unreactive; also called noble gases.
  15. Inner transition elements — the lanthanides and the actinides together, the f-block elements printed below the main body of the table.
  16. Isotopes — atoms of the same element having similar chemical properties but different atomic masses.
  17. Lanthanides — the fifteen elements of atomic numbers 57 to 71, belonging to group 3 of the sixth period; also called the rare earths.
  18. Mass number — the total number of protons and neutrons in the nucleus of an atom.
  19. Metals — lustrous elements that conduct heat and electricity well and tend to lose electrons.
  20. Non-metals — elements that lack metallic lustre, conduct heat and electricity poorly, and tend to gain or share electrons.
  21. Octave — a set of eight elements in Newland’s arrangement, in which the eighth resembles the first.
  22. Periodic properties — properties that reappear at regular intervals, or that increase or decrease steadily at regular intervals, along the table.
  23. Periodicity — the regular repetition of similar properties among the elements, caused by the recurrence of similar electronic configurations.
  24. Periods — the horizontal rows of the periodic table, whose elements all have the same number of shells.
  25. Representative elements — the elements of groups 1, 2 and 13 to 17, which have only their outermost shell incomplete.
  26. Transition elements — the elements of groups 3 to 12, which have their two outermost shells incomplete.
  27. Triad — a set of three chemically similar elements in which the middle atomic mass is about the mean of the other two.
  28. Typical elements — the third-period elements Na, Mg, Al, Si, P, S and Cl, which summarise the properties of their groups.
  29. Valence electrons — the electrons present in the outermost shell of an atom.
  30. Valency — the combining capacity of an element, equal to the number of hydrogen atoms one of its atoms can combine with or replace.

5.15 Points To Ponder

  1. Mendeleev corrected the accepted atomic masses of gold and platinum because his table disagreed with them — that is, he trusted a pattern over a measurement. When is it reasonable in science to doubt a measurement because a pattern disagrees with it, and when would doing so be dangerous?

  2. Newland’s arrangement broke partly because an entire family of elements, the noble gases, had not yet been discovered, and nothing in his scheme showed that anything was missing. Is there any way, from the modern table alone, to tell whether a family of elements is still missing today?

  3. The lanthanides and actinides belong inside group 3, but have to be printed below the table because the page is flat. If you were free to draw the periodic table in any shape at all, what shape would keep every element in its proper group without distorting the rest?

  4. Hydrogen has one valence electron like the alkali metals, and is one electron short of a stable arrangement like the halogens. Should a classification ever allow one member two places, or does that defeat the purpose of classifying?

  5. Reactivity increases down group 1 but decreases down group 17, and both trends have the same cause — the growing distance between the outermost shell and the nucleus. What does a single cause producing opposite effects tell you about learning trends by memory rather than by reason?

  6. The heaviest elements of period 7 exist only as a handful of atoms made in a laboratory, surviving for a fraction of a second. In what sense are they elements in the way that sodium and chlorine are, and what evidence would you want before giving one of them a permanent square in the table?

5.16 Objective Questions

  1. Choose the correct option.
    1. The property on which the modern periodic table is based is: (a) atomic mass (b) atomic number (c) valency (d) number of neutrons [1]
    2. An element with the electronic configuration 2, 8, 2 belongs to: (a) period 2, group 2 (b) period 3, group 2 (c) period 2, group 12 (d) period 3, group 12 [1]
    3. The number of elements in the fifth period of the modern periodic table is: (a) 8 (b) 18 (c) 32 (d) 2 [1]
    4. The valency of an element belonging to group 15 is: (a) 15 (b) 5 (c) 3 (d) 2 [1]
    5. The halogen that exists as a liquid at ordinary temperature is: (a) fluorine (b) chlorine (c) bromine (d) iodine [1]
    6. Down group 1, the atomic size: (a) decreases (b) increases (c) first increases then decreases (d) remains unchanged [1]
    7. The elements that have their two outermost shells incomplete are called: (a) representative elements (b) transition elements (c) inert gases (d) typical elements [1]
    8. The oxide of the third-period element that is the most acidic in character is: (a) \(\mathrm{Na_2O}\) (b) \(\mathrm{MgO}\) (c) \(\mathrm{SiO_2}\) (d) \(\mathrm{Cl_2O_7}\) [1]
  2. Each of the following gives an assertion (A) and a reason (R). Choose from: (a) both A and R are true and R is the correct explanation of A; (b) both A and R are true but R is not the correct explanation of A; (c) A is true but R is false; (d) A is false but R is true.
    1. Assertion (A): All the isotopes of an element occupy the same position in the modern periodic table. Reason (R): All the isotopes of an element have the same number of protons. [1]
    2. Assertion (A): Caesium is more reactive than lithium. Reason (R): The valence electron of caesium lies further from the nucleus and is lost more easily. [1]
    3. Assertion (A): The atomic radius of chlorine is greater than that of sodium. Reason (R): Chlorine has more protons in its nucleus than sodium has. [1]
    4. Assertion (A): Argon does not form compounds under ordinary conditions. Reason (R): Argon is a monoatomic gas. [1]
  3. Match each item in Column A with the correct item in Column B. [5]
Column A Column B
a. Group 3 of periods 6 and 7 i. Salt formers
b. Group 17 ii. Predicted the properties of eka-silicon
c. Moseley iii. Inner transition elements
d. Mendeleev iv. Diagonal relationship
e. Beryllium and aluminium v. Determined the nuclear positive charge
  1. Fill in the blanks.
    1. The number of shells in an atom is equal to its ……… number. [1]
    2. Moving from left to right across a period, the metallic character of the elements ……… . [1]
    3. The elements of group 2 form ……… ions carrying a charge of \(2+\). [1]
    4. Elements of the sixth period from atomic number 57 to 71 are called the ……… . [1]
    5. Alkali metals are stored under ……… to keep air and moisture away from them. [1]
  2. State whether each statement is true or false.
    1. The last element of every period has eight electrons in its outermost shell. [1]
    2. Newland’s law of octaves held good for elements heavier than calcium. [1]
    3. The valency of the elements of group 18 is zero. [1]
    4. Transition elements generally form colourless compounds. [1]
  3. Name the following.
    1. The scientist who proposed that the atomic mass of the middle element of a set of three is the mean of the other two. [1]
    2. The alkaline earth metal that imparts an apple green colour to a flame. [1]
    3. The noble gas used to fill electric light bulbs. [1]
    4. The metalloid of period 3 belonging to group 14. [1]
    5. The element placed at the top of group 1 that is not an alkali metal. [1]
  4. Give one word or phrase for each of the following.
    1. The number of unit positive charges present in the nucleus of an atom. [1]
    2. The horizontal rows of the modern periodic table. [1]
    3. Atoms of the same element having different atomic masses. [1]
    4. An oxide that reacts both with acids and with alkalis. [1]
  5. Pick the odd one out in each set and give a reason for your choice.
    1. Li, Na, K, Ca [1]
    2. F, Cl, Br, Ne [1]
    3. Fe, Cu, Zn, Mg [1]

5.17 Check Your Concepts

  1. Define the term valency, and state the valency of an element belonging to group 16. [2]

  2. State the modern periodic law. How does it differ from the law stated by Mendeleev? [2]

  3. Give a reason: the atoms of all the elements of a group have the same valency. [2]

  4. Name the two sets of inner transition elements and state where they are printed in the periodic table. [2]

  5. Why are the elements of group 18 described as noble rather than inert today? [2]

  6. An element has the electronic configuration 2, 8, 7. State its period, its group and its valency, and say whether it is a metal or a non-metal. [3]

  7. Give a reason for each of the following.

    1. Alkali metals do not occur in the free state in nature.
    2. The atom of sodium is larger than the atom of chlorine.
    3. Helium is used in balloons. [3]

  8. Differentiate between representative elements and transition elements, stating the basis of the difference. [3]

  9. Explain, with one example each, why the oxides of the elements become more basic down a group but more acidic across a period. [3]

  10. Explain ‘Newland’s law of octaves’ with one example, and give two reasons why it was discarded. [4]

  11. What would happen to the arrangement of the modern periodic table if the elements were ordered by atomic mass instead of atomic number? Support your answer with two examples. [4]

  12. Differentiate between the alkali metals and the alkaline earth metals with respect to (a) valence electrons (b) hardness (c) strength of their hydroxides (d) reactivity with water. [4]

  13. Describe the main characteristics of the transition elements, and state why they are placed between groups 2 and 13 rather than at either end of the table. [5]

  14. Mendeleev’s table carried four clear merits and several defects. State any three merits, and explain how the discovery of the atomic number removed two of the defects. [5]

5.18 Equations, Conversions & Identification

  1. Write the word equation and then the balanced chemical equation for the reaction of sodium with water. Name the gas evolved and state the test by which you would identify it. [3]

  2. Complete and balance the following equations.

    1. \(\mathrm{Ca} + \mathrm{HCl} \rightarrow\)
    2. \(\mathrm{Na_2O} + \mathrm{H_2O} \rightarrow\)
    3. \(\mathrm{Mg} + \mathrm{H_2SO_4} \rightarrow\) [3]

  3. Write the balanced equation for the burning of sodium in oxygen, and state what you would observe when the freshly cut metal is left exposed to air for a few minutes. [3]

  4. State what you would see in each case.

    1. A salt of potassium is held in a flame.
    2. A jar of bromine vapour is compared with a jar of chlorine.
    3. A small piece of calcium is dropped into water. [3]

  5. How would you obtain the following? Give the reasoning or the equation in each case.

    1. Sodium hydroxide starting from sodium metal, in two steps.
    2. Magnesium chloride starting from magnesium metal, in one step. [4]

  6. Identify the element in each case, giving your reasoning.

    1. An element of period 3 whose oxide has the formula \(\mathrm{X_2O_3}\).
    2. A gas of group 18 with only one shell in its atom.
    3. A solid non-metal of group 17 that gives a violet vapour. [3]

  7. Look at Fig. 5.5-1. Name the labelled parts of the apparatus, state what is changed from one reading to the next, and state what quantity is measured from the photographic plate. [3]

5.19 Practice With Numericals

  1. Three elements P, Q and R form a Döbereiner’s triad. The atomic mass of P is 32 and that of R is 128. Calculate the expected atomic mass of Q, showing your working. [2]

  2. Lithium, sodium and potassium have atomic masses 7, 23 and 39. Show by calculation whether they satisfy the law of triads. [2]

  3. An atom of an element has 17 protons and 18 neutrons.

    1. State its atomic number and its mass number.
    2. Write its electronic configuration.
    3. State its period, its group and its valency. [4]

  4. An element T has atomic number 20.

    1. Write the electronic configuration of T.
    2. State the number of its valence electrons and its valency.
    3. Write the formula of the compound formed between T and chlorine, and state whether it is electrovalent or covalent. [4]

  5. Using the mean rule for a triad, calculate the atomic mass expected for the middle element of the set formed by fluorine (19), chlorine (35.5) and bromine (80), and state whether these three elements form a triad. [3]