Classification of Elements and Periodicity in Properties: Class 11 Chemistry · Chapter 3

Walk into a busy medical store and ask for a cough syrup. The pharmacist does not search every shelf: syrups sit together, tablets of one kind sit together, and a new brand goes straight into its place. The periodic table does the same job for chemistry. It is a single chart that tells you where each element belongs, which elements behave alike, and roughly how an element you have never handled will behave.

This chapter follows the table from its first rough drafts to the long form printed in your book, and shows why its shape follows from electronic configuration. It then uses the table to explain how size, ionization enthalpy, electron gain enthalpy, electronegativity, valence and reactivity change from element to element. Almost every later chapter of chemistry leans on these trends.

3.1 Why do we need to classify elements?

Around the year 1800 chemists knew roughly thirty elements. By the 1860s the count had passed sixty, and today 118 elements are confirmed, the heaviest of them made in accelerators a few atoms at a time. Each element forms many compounds. Learning all of them one by one would be hopeless.

A good classification fixes this in two ways. It organises what is already known: once you know that sodium and potassium are alike, a fact about one tells you a lot about the other. It also predicts what is not yet known: an empty place in a pattern points to an element waiting to be found, and even hints at its properties. Both payoffs drove the search for a natural arrangement of the elements in the nineteenth century.

3.2 Genesis of Periodic Classification

No single person invented the periodic table. Several chemists, each working with the atomic weights (relative atomic masses) of their time, noticed pieces of the pattern. Follow their attempts in Fig. 3.1 and in the timeline below.

Triads (Döbereiner, 1829). The German chemist Johann Döbereiner found sets of three similar elements in which the middle member had an atomic weight close to the average of the other two, and properties between theirs. Sulphur, selenium and tellurium form such a triad. The idea worked for only a handful of sets, so most chemists treated it as a coincidence.

A spiral on a cylinder (de Chancourtois, 1862). The French geologist A. E. B. de Chancourtois wrote the elements in order of atomic weight along a spiral line wound round a cylinder. Similar elements then fell one above another. His paper reached few chemists and the idea was overlooked.

Octaves (Newlands, 1865). The English chemist John Newlands listed the elements by atomic weight and saw that each element resembled the one seven places after it, the eighth in the count. He compared this with music: after Sa Re Ga Ma Pa Dha Ni comes the upper Sa, which sounds like the first. The rule held only as far as calcium; beyond it, metals such as iron landed among non-metals. Fellow chemists mocked the idea, yet the Royal Society awarded Newlands its Davy Medal in 1887.

Curves and tables (Lothar Meyer and Mendeleev, 1868–1870). In Germany, Lothar Meyer took a graphical route. He drew a curve of atomic volume, the volume occupied by a mole of atoms, against atomic weight. It climbs to a sharp peak at each alkali metal and dips in between, like a row of waves, and the waves get wider for heavier elements, something Newlands' fixed octaves could not capture. Meyer's own table dates from 1868, but the Russian chemist Dmitri Mendeleev (1834–1907) got his into print first, in 1869, so Mendeleev usually receives the credit for the periodic law.

Mendeleev is said to have written each element, its atomic weight and its properties on a card, then rearranged the cards until families lined up. His table ran by atomic weight along the rows, with chemically similar elements stacked in columns, or groups. What made it stronger than Meyer's was the evidence behind it: not one physical property but many chemical ones, above all the formulas of each element's oxide and hydride. Two bold decisions made it stand out (Fig. 3.2).

1. Properties over weights. Tellurium (127.6) is heavier than iodine (126.9), so strict weight order would put iodine with oxygen and sulphur and tellurium with the halogens. Mendeleev swapped them, trusting chemistry over the numbers, and suspected the weights were wrong.
2. Gaps for the unknown. Where no known element fitted a column, he left a blank and described the missing element. He called the one below aluminium eka-aluminium and the one below silicon eka-silicon (eka is Sanskrit for 'one'). Gallium (1875) and germanium (1886) were later found with properties very close to his forecasts.

3.3 Modern Periodic Law and the Present Form of the Periodic Table

Mendeleev worked without knowing what an atom contains. Electrons were discovered only in 1897 and the nucleus in 1911, so he could not explain why his law worked, nor why a few pairs such as tellurium and iodine refused to obey weight order. The answer came from X-rays.

In 1913 the English physicist Henry Moseley fired fast electrons at targets made of different elements. Each element gave off X-rays of its own characteristic frequencies $\nu$. When he plotted $\sqrt{\nu }$ for the strongest line against the element's position in the table, the points fell on a straight line; against atomic mass they did not (Fig. 3.3). The quantity that rose by exactly one from each element to the next was the atomic number $Z$. It is the charge on the nucleus, equal to the number of protons and to the number of electrons in the neutral atom.

With atomic number as the order, the old misfits behave. Argon ($Z=18$) comes before potassium ($Z=19$) though it is heavier; cobalt (27) comes before nickel (28); tellurium (52) before iodine (53). Mendeleev's chemical instinct had been right, and the weights were never wrong: they simply were not the right thing to sort by.

Ninety-four elements are found in nature, if we count the traces of neptunium and plutonium that form in uranium ores; every heavier element has been made artificially, in reactors or accelerators. Today we see the periodic law as a consequence of something deeper: electronic configurations repeat as shells fill, and configuration fixes chemistry. Section 3.5 makes this connection precise.

The table in your book is the long form, the most widely used of the many layouts chemists have tried, from spirals to pyramids. Read across it and you move along a period (आवर्त); read down and you stay within a group (वर्ग), a family of elements whose outer electronic configurations match. Groups are numbered 1 to 18 from left to right, as IUPAC recommends; older books use labels such as IA, IIB, VIII and 0, so expect to meet both.

• There are seven periods. The period number is the principal quantum number $n$ of the outermost shell.
• The periods hold 2, 8, 8, 18, 18, 32 and 32 elements. Period 7 has been complete since elements 113 to 118 were confirmed.
• Fourteen elements of period 6 (the lanthanoids) and fourteen of period 7 (the actinoids) are printed as two separate rows at the foot of the table, so the main table is not too wide.

3.4 Nomenclature of Elements with Atomic Numbers > 100

Who gets to name a new element? Traditionally its discoverers, with IUPAC's approval. That works until two teams claim the same discovery. From the late 1960s, laboratories at Berkeley in the USA and at Dubna in the Soviet Union both reported element 105. Berkeley called it hahnium, Dubna nielsbohrium, and the argument dragged on until IUPAC settled on dubnium in 1997. Disputes like this arise easily, because superheavy elements are made only a few atoms at a time, in a handful of laboratories with very costly equipment, and many of the atoms decay within seconds.

IUPAC's fix is a temporary name that nobody can argue about, because it is simply the atomic number spelt out. Write each digit of $Z$ as its root from Table 3.2 (the roots come from Latin and Greek numbers), run the roots together and finish with -ium. The first letters of the roots make a three-letter symbol.

• Two spelling rules: when enn comes before nil, one n is dropped (ennil); when the name would end in bi-ium or tri-ium, the i is dropped (bium, trium).
• Element 116, for example, was ununhexium (Uuh) until it was named livermorium (Lv).

The temporary name lasts only until a joint IUPAC–IUPAP panel has checked the discovery claim. The discoverers then propose a permanent name with a one- or two-letter symbol, often honouring a place, a laboratory or a scientist, and IUPAC approves it after a period of public comment. Every element up to 118 has now been through this process; the last four names were confirmed in 2016 (Table 3.3).

3.5 Electronic Configurations of Elements and the Periodic Table

Here is the payoff of the previous chapter. Take the aufbau order, 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, and give each new element one more electron than the last. Start a new row every time a new shell is opened, and line up each element under the one whose outer subshells match. The long form of the table appears without any further rules.

(a) Electronic configurations in periods

A new period begins every time electrons start a new shell, so the period number equals $n$ of the valence shell. The length of a period equals the number of electrons that can go into the subshells filled during it, that is, twice the number of orbitals being filled. Build the table row by row in Fig. 3.5.

• Period 1 ($n=1$): only the 1s orbital, so two elements, hydrogen ($1s^1$) and helium ($1s^2$).
• Periods 2 and 3: one $ns$ and three $np$ orbitals, so eight elements each, lithium to neon and sodium to argon. The 3d orbitals lie above 4s in energy, so they are not used in period 3.
• Period 4: potassium and calcium fill 4s; then the ten elements from scandium ($Z=21$, $3d^{1}4s^2$) to zinc ($Z=30$, $3d^{10}4s^2$) fill 3d, the first transition series; gallium to krypton fill 4p. Nine orbitals, eighteen elements.
• Period 5: rubidium to xenon in the same pattern, with the 4d series starting at yttrium ($Z=39$). Again eighteen.
• Period 6: 6s, 4f, 5d and 6p, sixteen orbitals, 32 elements. The 4f subshell fills from cerium ($Z=58$) to lutetium ($Z=71$): the lanthanoids.
• Period 7: 7s, 5f, 6d and 7p, again 32 elements, ending with oganesson ($Z=118$) in group 18. The 5f subshell fills after actinium ($Z=89$) from thorium to lawrencium: the actinoids, most of them made artificially.

Why print the lanthanoids and actinoids separately? Squeezing 28 more elements into periods 6 and 7 would make the table very wide. More importantly, the fourteen members of each series are so alike that they would all deserve the single column of group 3. Printing them as two rows keeps every column a family of similar elements.

(b) Groupwise electronic configurations

Read down a column instead of along a row and the outer configuration stays the same while $n$ grows. Elements of one group therefore have the same number of valence electrons arranged the same way, and that is why they behave alike. The alkaline earth metals of group 2 all end in $n{s}^2$ (Table 3.4).

Notice that the atomic numbers down a group do not rise by a fixed amount: they jump by 8, 8, 18, 18 and 32, the lengths of the periods in between. Configuration, and so chemistry, depends on atomic number, which is why the modern law uses it and not atomic mass.

3.6 Electronic Configurations and Types of Elements: s-, p-, d-, f-Blocks

Because the subshell that receives the last electron decides an element's column, the table splits naturally into four blocks, named after that subshell. Their shapes match the subshell capacities: two columns for s, six for p, ten for d and fourteen for f.

Two elements break the simple rule. Helium is $1s^2$, so strictly it is an s-block element, but its shell is full and it is as unreactive as the other noble gases, so it stands at the top of group 18. Hydrogen has a single 1s electron. Like the group 1 metals it can lose it to form $\mathrm{H}{\phantom{X}}^{+}$; like the halogens it can gain one to reach the helium configuration, forming $\mathrm{H}{\phantom{X}}^{-}$. Because it fits neither family neatly, many tables place it on its own at the top.

3.6.1 The s-Block Elements

The s-block is just two columns wide: group 1, the alkali metals ($n{s}^1$), and group 2, the alkaline earth metals ($n{s}^2$). Sodium has to be kept under kerosene, and none of these metals is found uncombined in rocks or seawater. Their ionization enthalpies are low, so they give up their one or two outer electrons readily and become $\mathrm{M}{\phantom{X}}^{+}$ or $\mathrm{M}{\phantom{X}}^{2+}$ ions. Down either column the atoms get bigger, the outer electrons go even more easily and the metals grow more reactive; lithium fizzes gently in water, caesium explodes. Most s-block compounds, common salt among them, are ionic. Lithium and beryllium, the tiny first members, are the exceptions: their compounds are noticeably covalent.

3.6.2 The p-Block Elements

Six columns, groups 13 to 18, form the p-block, where the outer configuration runs from $n{s}^{2}n{p}^1$ to $n{s}^{2}n{p}^6$. Read it from the right. Group 18 holds the noble gases: every valence orbital is full, so they neither give up nor accept electrons easily and hardly react at all. Their neighbours are the opposite. The chalcogens of group 16 and the halogens of group 17 sit just short of a full shell, by two electrons and by one. Their electron gain enthalpies are strongly negative, so they grab electrons readily; they are some of the most important non-metals. Moving left within the block, or down any of its columns, metallic character increases. The s- and p-blocks together make up the main group, and their members are known as representative elements.

3.6.3 The d-Block Elements (Transition Elements)

Iron, copper, silver, gold and platinum all come from the ten-column d-block in the middle of the table, groups 3 to 12. The old name transition elements (संक्रमण तत्व) comes from their position, between the very reactive s-block metals on the left and the less active metals of groups 13 and 14 on the right. What sets them apart electronically is where the last electron goes: not into the outermost shell but into the $(n-1)d$ subshell beneath it, giving the pattern $(n-1)d^{1-10}n{s}^{0-2}\mathrm{.}$ Palladium, $4d^{10}5s^0$, has no outer s electron at all. Every d-block element is a metal. Many show several oxidation states, give coloured ions (think of blue copper sulphate solution), are paramagnetic and work as catalysts. Zinc, cadmium and mercury, whose $(n-1)d^{10}n{s}^2$ arrangement is complete, show few of these features. Class 12 chemistry gives the d-block a chapter of its own.

3.6.4 The f-Block Elements (Inner-Transition Elements)

The two rows printed at the foot of the table make up the f-block. Here the incoming electron goes two shells in from the outside, into 4f or 5f, so the outer pattern is $(n-2)f^{1-14}(n-1)d^{0-1}n{s}^2$ and the elements are called inner-transition elements. The first row, cerium ($Z=58$) to lutetium ($Z=71$), is the lanthanoids; the second, thorium ($Z=90$) to lawrencium ($Z=103$), is the actinoids. Because the electrons being added are buried so deep, neighbours within each series are remarkably alike, and all of them are metals. Every actinoid is radioactive. The early actinoids show many oxidation states, which makes their chemistry richer than that of the lanthanoids. Many later ones have been made only in nanogram amounts or less, so little of their chemistry is known. Elements beyond uranium ($Z=92$) are called transuranium elements.

3.6.5 Metals, Non-metals and Metalloids

Long before electrons were known, chemists sorted elements into metals and non-metals, and the division still matters. More than three quarters of all elements are metals, filling the left side and the centre of the table; the non-metals are crowded into its upper right corner.

You can tell the two kinds apart by handling them. A metal is lustrous, conducts heat and electricity well, and can be beaten into sheets (malleable, like the silver varq on sweets) or drawn into wire (ductile, like copper). Nearly all metals are solids with high melting points; mercury is a liquid, and gallium (about 30 °C) and caesium (about 28 °C) melt on a hot day. Non-metals are solids or gases at room temperature, bromine being the one liquid. They usually melt and boil at low temperatures, conduct poorly, and break rather than bend when solid; boron and carbon, with extremely high melting points, are the exceptions.

The change from one kind to the other is gradual. Across a period, metallic character fades and non-metallic character grows; down a group, elements become more metallic. Along a stepped diagonal line through the p-block lie elements with a mixture of properties: boron, silicon, germanium, arsenic, antimony and tellurium (Fig. 3.7). They are called metalloids (उपधातु) or semi-metals. Silicon, for instance, looks shiny like a metal but is brittle and conducts electricity only modestly, which is exactly what makes it useful in computer chips and solar cells.

3.7 Periodic Trends in Properties of Elements

Look along period 3 and reactivity follows a U-shape: sodium reacts furiously, silicon hardly at all, chlorine furiously again. Look down a group and the direction depends on the family: the alkali metals grow more reactive from lithium to caesium, the halogens less reactive from fluorine to iodine. Patterns like these call for an explanation in terms of the electrons and energy levels of the atoms. This section builds it, starting with five measurable or estimated properties of isolated atoms and then turning to chemical behaviour.

3.7.1 Trends in Physical Properties

Plot almost any physical property against atomic number, whether the temperature at which an element melts or the energy needed to boil or atomise it, and a repeating pattern appears. Five atomic properties explain most of the chemistry, so we study those. They are the sizes of atoms and of ions, the energy needed to remove an electron, the energy change when one is added, and the pull an atom exerts on shared electrons.

(a) Atomic radius

Where does an atom end? Its electron cloud has no surface: the chance of finding an electron simply fades with distance from the nucleus (first step of Fig. 3.8). Atoms are also about ${10}^{-10}\mathrm{m}$ across, far too small to hold against any scale. So chemists never measure a lone atom. They measure something that is sharp, the distance between two nuclei held together in a molecule or a crystal, which X-ray diffraction and spectroscopy give very precisely, and take half of it as the radius.

• Covalent radius, for non-metals: half the length of a single bond between two identical atoms. The I–I bond in iodine is 267 pm long, so the covalent radius of iodine is about 133 pm.
• Metallic radius, for metals: half the distance between neighbouring nuclei in the metal crystal. In aluminium the neighbours are 286 pm apart, giving 143 pm.
• van der Waals radius: half the distance between atoms that touch without bonding, as in solid argon. It is much larger, so noble gas radii must not be compared with covalent radii.

The numbers show a clear pattern. Across a period the radius shrinks: from roughly 190 pm for sodium to roughly 100 pm for chlorine. Down a group it grows: an alkali metal atom goes from about 150 pm for lithium to about 265 pm for caesium. To see why, we need the idea of effective nuclear charge (प्रभावी नाभिकीय आवेश).

An outer electron is attracted by the nucleus but repelled by every other electron. Inner electrons sit between it and the nucleus and shield (or screen) part of the nuclear charge, so the outer electron feels a net charge $Z_{\text{eff}}$ smaller than $Z$. Electrons in the same shell shield one another only weakly. Across period 2, from lithium to fluorine, $Z$ rises by one at each step but the core stays the same two 1s electrons, so the outer electrons feel an ever stronger pull and are drawn closer. Going from lithium to sodium, by contrast, a whole new shell is added and ten core electrons now shield the nucleus. The outer electron feels about the same net charge but sits in $n=3$, much farther out (Fig. 3.9).

(b) Ionic radius

Turn a potassium atom (227 pm) into $\mathrm{K}{\phantom{X}}^{+}$ and its radius drops to 138 pm; turn a chlorine atom (99 pm) into $\mathrm{C}\mathrm{l}{\phantom{X}}^{-}$ and it grows to 181 pm (Fig. 3.10). The two changes mirror each other. The cation has lost an electron but kept all its protons, and here that one 4s electron was the entire outer shell, so what remains is a smaller, tightly held core. The anion has gained an electron but no protons, so repulsion among the electrons grows, each feels a little less of the nucleus, and the cloud swells. Ionic radii themselves come from the spacing of cations and anions in ionic crystals, and they follow the usual period and group trends.

Atoms and ions with the same number of electrons are called isoelectronic species (समइलेक्ट्रॉनी). $\mathrm{S}{\phantom{X}}^{2-}$, $\mathrm{C}\mathrm{l}{\phantom{X}}^{-}$, $\mathrm{K}{\phantom{X}}^{+}$ and $\mathrm{C}\mathrm{a}{\phantom{X}}^{2+}$ all have 18 electrons, like argon. With the electron count fixed, size depends on nuclear charge alone: the more protons, the tighter the cloud. So the size order is $\mathrm{S}{\phantom{X}}^{2-}>\mathrm{C}\mathrm{l}{\phantom{X}}^{-}>\mathrm{K}{\phantom{X}}^{+}>\mathrm{C}\mathrm{a}{\phantom{X}}^{2+}:$ the anion with the larger negative charge is the largest, the cation with the larger positive charge the smallest.

(c) Ionization enthalpy

How tightly does an atom hold its electrons? The measure is the ionization enthalpy (आयनन एन्थैल्पी), ${\mathrm{\Delta }}_{i}H$. It is the energy needed to remove the most loosely bound electron from an isolated gaseous atom in its ground state. 'Isolated' and 'gaseous' mean no neighbours pull on the atom; 'ground state' means every atom starts from the same lowest-energy arrangement. Only then can values for different elements be compared fairly. The first ionization enthalpy is the enthalpy change of

and it is quoted in $\mathrm{k}\mathrm{J}\mathrm{m}\mathrm{o}{\mathrm{l}}^{-1}\mathrm{.}$ Removing a second electron, from the ion, gives the second ionization enthalpy:

Pulling a negative electron away from a positive nucleus always costs energy, so every ionization enthalpy is positive. Each successive one is bigger than the last: the second electron has to escape from an ion that already carries a +1 charge, the third from a +2 ion, and so on. The rise is gradual while valence electrons are removed and then jumps enormously when the next electron has to come out of the noble-gas core (Fig. 3.11). When chemists say 'ionization enthalpy' without a qualifier, they mean the first.

Because a smaller atom holds its electrons closer, ionization enthalpy tracks atomic radius in reverse, and the same two trends appear (Fig. 3.12).

• Across a period, ${\mathrm{\Delta }}_{i}H$ generally increases. New electrons enter the same shell, shield one another poorly, and feel a rising $Z_{\text{eff}}$, so they are held more tightly.
• Down a group, ${\mathrm{\Delta }}_{i}H$ decreases. The outer electron sits in a higher shell, farther out, behind more inner electrons, and together these effects beat the growing nuclear charge.

Over a longer stretch of the table the result is a saw-tooth. The noble gases, with their closed shells, sit at the peaks; the alkali metals, each with a single $ns$ electron outside a noble-gas core, sit in the troughs, which is why they are so reactive. You can draw this pattern yourself in the activity below.

The dips are worth understanding. Magnesium to aluminium: within one shell, an s electron penetrates closer to the nucleus than a p electron, so a 3p electron is shielded more by the inner electrons. Aluminium's 3p electron is therefore easier to remove than magnesium's 3s electron, even though aluminium has one more proton. Phosphorus to sulphur: phosphorus has three 3p electrons in three separate orbitals (Hund's rule), while in sulphur one 3p orbital holds a pair. The repulsion within that pair makes one of them easier to remove. The same two dips appear in period 2, between beryllium and boron and between nitrogen and oxygen.

(d) Electron gain enthalpy

Ionization enthalpy measures how hard it is to take an electron away. The opposite question, how readily an atom accepts one, is answered by the electron gain enthalpy (इलेक्ट्रॉन लब्धि एन्थैल्पी), ${\mathrm{\Delta }}_{eg}H$, the enthalpy change when an isolated gaseous atom gains an electron to become an anion:

Think of it as the welcome an atom gives an extra electron. A halogen atom, one electron short of a noble-gas shell, welcomes it warmly: a lot of energy is released and ${\mathrm{\Delta }}_{eg}H$ is large and negative. A noble gas atom has no room left in its valence shell, so the newcomer would have to sit alone in the next shell, far out and weakly held; energy has to be put in and ${\mathrm{\Delta }}_{eg}H$ is positive. Most elements fall in between, with modest negative values. Adding a second electron to an ion that is already negative always needs energy, because of the repulsion.

Size and effective nuclear charge set the broad pattern, although it is bumpier than for ionization enthalpy. Small atoms with a high $Z_{\text{eff}}$, towards the right of a period, hold an added electron tightly; larger atoms lower down a group hold it less tightly. So the most negative values sit in the top right corner, in groups 16 and 17.

Yet the very top of those groups breaks the rule. Chlorine's value (−349 kJ mol⁻¹) is more negative than fluorine's (−328 kJ mol⁻¹), and sulphur's (−200 kJ mol⁻¹) more negative than oxygen's (−141 kJ mol⁻¹). Fluorine and oxygen are so small that their $n=2$ shell is already crowded, and an extra electron there is strongly repelled by the others. Chlorine and sulphur add the electron to the larger $n=3$ shell, where it has more room (Fig. 3.13).

(e) Electronegativity

In a hydrogen chloride molecule the bonding pair spends more of its time near chlorine than near hydrogen. The tendency of an atom in a molecule to draw a shared pair towards itself is its electronegativity (विद्युत्-ऋणात्मकता). Because it describes a bonded atom rather than an isolated one, no single experiment measures it; it is expressed on constructed scales instead. Linus Pauling built the first and most popular in 1932 from bond energies, placing fluorine, the strongest electron-puller, near 4 (3.98 on today's revised scale). The Mulliken–Jaffé and Allred–Rochow scales use other data but give similar orderings.

Electronegativity increases across a period (lithium 0.98, fluorine 3.98) and decreases down a group (fluorine 3.98, iodine 2.66), following ionization enthalpy, because a small atom with a high $Z_{\text{eff}}$ holds bonding electrons close. Treat the numbers as guides rather than constants: an element's electronegativity shifts somewhat with its oxidation state and its bonding partners. Even so, the difference in electronegativity between two atoms is a good predictor of the kind of bond they form, as the next chapter shows.

High electronegativity is the mark of a non-metal, which gains electrons readily; low electronegativity is the mark of a metal, which loses them. So a map of electronegativity doubles as a map of character: as electronegativity climbs across a period, metallic character falls, and as it drops down a group, metallic character rises. Explore all three atomic properties together in Fig. 3.14, then use Table 3.6 as a one-glance summary.

3.7.2 Periodic Trends in Chemical Properties

Some chemical trends, such as the inert pair effect and the lanthanoid contraction, make most sense group by group and come in later chapters. Two follow so directly from configuration that we take them now: how valence repeats along the periods, and why the second-period elements are the odd ones out.

(a) Periodicity of valence or oxidation states

Count an element's outer electrons and you can usually predict how many bonds it forms. For a representative element the valence is normally either that count or eight minus it. Sodium, with one outer electron, has valence 1; chlorine, with seven, shows valence 1 (eight minus seven) towards hydrogen and up to 7 towards oxygen. Across period 3 the valence towards hydrogen rises from 1 to 4 and falls back to 1, while the highest valence towards oxygen climbs all the way to 7 (Table 3.7).

Today chemists usually speak of oxidation state (ऑक्सीकरण अवस्था) rather than valence. To find it, give each shared pair to the more electronegative atom and count the charge left on each atom. In $\mathrm{H}\mathrm{C}\mathrm{l}$ chlorine is more electronegative than hydrogen, so chlorine is $-1$ and hydrogen $+1$. In $\mathrm{C}\mathrm{l}{\phantom{X}}_2\mathrm{O}$ oxygen is more electronegative than chlorine, so oxygen is $-2$ and each chlorine is $+1$. The same element can therefore show different oxidation states, depending on its partner.

Valence is not always a single number. Phosphorus forms both $\mathrm{P}\mathrm{C}\mathrm{l}{\phantom{X}}_3$ and $\mathrm{P}\mathrm{C}\mathrm{l}{\phantom{X}}_5$, and among the transition metals and actinoids several oxidation states for one element are the rule rather than the exception.

(b) Anomalous properties of second period elements

Lithium is an alkali metal, yet in several ways it behaves more like magnesium than like sodium; beryllium, likewise, resembles aluminium more than magnesium. Both form compounds with marked covalent character, whereas the heavier members of groups 1 and 2 form mostly ionic ones. In the same way, each second-period element from lithium to fluorine stands apart from the rest of its group. The resemblance of an element to its lower-right neighbour is called the diagonal relationship (विकर्ण संबंध); see Fig. 3.15.

Why are the second-period elements so different? Their atoms are unusually small, so their ions carry a lot of charge for their size and they attract electrons strongly. On top of this, a second-period atom has only four valence orbitals, 2s and three 2p, so it cannot form more than four covalent bonds. From period 3 onward the valence shell also offers d orbitals (3s, 3p and 3d make nine), so the heavier members can take on more than four electron pairs. Nitrogen stops at $\mathrm{N}\mathrm{C}\mathrm{l}{\phantom{X}}_3$ while phosphorus goes on to $\mathrm{P}\mathrm{C}\mathrm{l}{\phantom{X}}_5$, and oxygen stops at $\mathrm{O}\mathrm{F}{\phantom{X}}_2$ while sulphur forms $\mathrm{S}\mathrm{F}{\phantom{X}}_6$.

Small second-period atoms can also come close enough for their p orbitals to overlap sideways, forming strong $p\pi$–$p\pi$ multiple bonds. Nitrogen is the triple-bonded $\mathrm{N}{\phantom{X}}_2$ and oxygen the double-bonded $\mathrm{O}{\phantom{X}}_2$, while phosphorus prefers $\mathrm{P}{\phantom{X}}_4$ molecules and sulphur $\mathrm{S}{\phantom{X}}_8$ rings, both held together by single bonds. Carbon dioxide is a small molecule with two C=O double bonds; silicon dioxide is a giant network of Si–O single bonds.

3.7.3 Periodic Trends and Chemical Reactivity

Put the trends together and reactivity follows. Going left to right along a period, atoms shrink, ionization enthalpy climbs and electron gain enthalpy grows more negative. The far-left element therefore parts with an electron more easily than any other, and the far-right element (leaving out the noble gas) takes one on most eagerly. Both ends are highly reactive, in opposite ways: one forms cations and behaves as a metal (and, as later chapters show, as a reducing agent); the other forms anions and behaves as a non-metal (an oxidising agent). In the middle of the period neither process is easy, so reactivity is lowest there.

Oxides make this visible. Sodium's oxide, $\mathrm{N}\mathrm{a}{\phantom{X}}_2\mathrm{O}$, dissolves in water to give a strong alkali, so it is basic; the highest oxide of chlorine, $\mathrm{C}\mathrm{l}{\phantom{X}}_2\mathrm{O}{\phantom{X}}_7$, gives a strong acid, so it is acidic. Between them the character shifts step by step (Fig. 3.16). Aluminium oxide is amphoteric (उभयधर्मी), dissolving in acids and in alkalis alike, and a few oxides of non-metals, such as carbon monoxide and dinitrogen oxide, are neutral and react with neither.

Transition metals change more gently. Their added electrons go into an inner d subshell, so atomic radii shrink only slightly across the 3d series, and across the 4f series the change is smaller still. Their ionization enthalpies lie above those of the s-block metals and below those of most p-block elements, so they are less electropositive than the metals of groups 1 and 2. Iron rusts slowly, while a freshly cut piece of sodium tarnishes within seconds.

Down a main group, bigger atoms lose electrons more readily and attract extra ones less strongly (apart from the period-2 exceptions in part (d)). Metallic character therefore grows down a group, which makes caesium the most reactive alkali metal, while non-metallic character fades, which makes fluorine the most reactive halogen. Transition metals buck this trend: gold is far less reactive than copper two places above it. Its atoms are only a little larger than copper's (144 against 128 pm) while its nuclear charge is far higher (+79 against +29). So gold holds its electrons more tightly: its first ionization enthalpy is 890 kJ mol⁻¹ against copper's 745 kJ mol⁻¹. Reactivity also decides how elements occur in nature. Sodium, potassium and fluorine are always found combined in compounds, while unreactive gold and platinum can be dug up as the free metals.