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Chapter 3: Classification of Elements and Periodicity in Properties

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In 1869, a Russian chemist named Dmitri Mendeleev was trying to bring order to a growing crowd of about sixty known elements. He wrote the properties of each element on cards and arranged them, again and again, in the increasing order of their atomic weights. A striking pattern emerged: after fixed intervals, elements with similar properties reappeared, almost like the beats of a song.

But Mendeleev did something bold. Where an element with the “right” properties did not exist, he refused to force a wrong element into the slot. Instead, he left the gap empty and declared that an undiscovered element must belong there. He even predicted the weight, density, and the formulas of the oxides and chlorides of these missing elements, naming them “eka-aluminium” and “eka-silicon.”

Years later, when gallium and germanium were finally discovered, their measured properties matched Mendeleev’s predictions almost exactly. A table built from careful observation had successfully foretold the existence of elements no one had yet seen. That is the power of the Periodic Table you are about to study — not just a chart to memorise, but a map that reveals why elements behave the way they do, and lets us predict properties before we ever meet an element in the laboratory.

Fig. 3.0

Fig. 3.0 – Mendeleev and the prediction of eka-aluminium and eka-silicon.

NoteCuriosity Corner

Q1. Mendeleev ordered elements by atomic weight, yet he deliberately placed iodine after tellurium even though iodine is the lighter of the two. Could there be a more fundamental property than weight that fixes such “misfits” automatically? Q2. Why should elements sitting in the same vertical column of the table — sometimes with very different weights — end up showing such similar chemical behaviour? Q3. As you move from left to right across a period you keep adding electrons, yet the atoms actually get smaller. How can gaining more electrons make an atom shrink? Q4. The noble gases sit quietly at the far right of the table and refuse to take part in most reactions. What is so special about their electron arrangement that makes them this unwilling to gain or lose electrons?

This chapter will give you the reasoning tools to answer each of these questions for yourself.

3.1 Why Do We Need to Classify Elements?

Everything around you — the air you breathe, the water you drink, the metal casing of your phone — is built from a small set of basic substances called elements. An element cannot be broken down into anything simpler by ordinary chemical means. This is why we treat elements as the basic units of all types of matter.

The number of known elements has grown steadily over time. In 1800, only 31 elements were known. By 1865, that number had more than doubled to 63, and at present 114 elements are known. Some of the most recently discovered ones do not occur in nature at all; they are man-made, created in the laboratory, and efforts to synthesise still newer elements are continuing. The steady rise in the number of known elements is easy to picture as a simple timeline, shown in Fig. 3.1.

Fig. 3.1

Fig. 3.1 – Growth in the number of known elements over time.

The heaviest members of the modern Periodic Table — the man-made “superheavy” elements — are not dug out of the earth; they are built one atom at a time. In specialised laboratories, scientists accelerate lighter atomic nuclei to high speeds and smash them together, hoping they will fuse into a single heavier nucleus. These artificial elements are so unstable that only a handful of atoms are ever produced, and each may survive for only a tiny fraction of a second before it decays. This is exactly why detecting, studying, and even naming such elements is slow, painstaking, and highly competitive work.

Why We Group Elements

With such a large number of elements, it becomes very difficult to study the chemistry of each one individually, together with their innumerable compounds. Imagine trying to learn the behaviour of every single element in isolation — the task would be almost endless, and we would never see the connections between them.

To ease this problem, scientists searched for a systematic way to organise their knowledge by classifying the elements. A good classification does two powerful things. First, it rationalizes the chemical facts we already know, by placing elements that behave alike into the same groups. Second — and this is the truly remarkable part — it can even help us predict the properties of elements that have not yet been discovered or studied, pointing chemists toward fresh directions for further research.

Q. How can a mere classification predict an element that no one has discovered yet?

A. When elements are arranged by their shared behaviour, a gap in the repeating pattern signals a “missing” element. Its likely properties can then be estimated from the known elements that surround it in the same group and period.

In the sections that follow, we will trace the historical development of the Periodic Table as it stands today, arrive at the Modern Periodic Law, see how the periodic classification follows as a logical consequence of the electronic configuration of atoms, and finally examine the periodic trends in the physical and chemical properties of the elements.

3.2 Genesis of Periodic Classification

The Periodic Law and the Periodic Table did not appear overnight. They grew slowly, as the consequence of many scientists patiently systematising the knowledge gained through their observations and experiments. Each researcher noticed a small piece of the pattern, and the next built upon it. Before we look at the final table, it is worth walking through these early attempts, because each one teaches us something about how a good classification is judged.

A recurring idea in this story is that certain properties of elements seem to repeat at regular intervals when the elements are lined up by their atomic weight. Here, atomic weight simply means the relative mass of an atom of the element compared with a standard. As we shall see, spotting where the repetition happens is exactly what allowed chemists to organise the elements.

Döbereiner’s Triads

The German chemist Johann Döbereiner, in the early 1800s, was the first to seriously consider the idea of trends among the properties of elements. By 1829, he had noticed a similarity among the physical and chemical properties of several groups of three elements, which he called Triads.

In each triad, he observed something remarkable about the middle element. Its atomic weight fell about half way between the atomic weights of the other two, and its properties, too, lay in between those of the other two members. The triads he identified are listed in Table 3.1.

Table 3.1 Döbereiner’s Triads

Element Atomic weight Element Atomic weight Element Atomic weight
Li 7 Ca 40 Cl 35.5
Na 23 Sr 88 Br 80
K 39 Ba 137 I 127

Notice, for example, that sodium’s atomic weight of 23 is very close to the average of lithium (7) and potassium (39), which is \(\frac{7+39}{2}=23\). This numerical regularity, referred to as the Law of Triads, was the first hint that the properties of elements might be linked to their atomic weights in a systematic way.

However, the relationship seemed to work only for a few sets of elements. Because it could not be extended to most of the known elements, it was dismissed at the time as a mere coincidence.

(question only — no solution) Using Döbereiner’s law of triads, and the atomic weights of chlorine (35.5) and iodine (127), predict the approximate atomic weight of bromine, the middle member of this halogen triad.

de Chancourtois and Newlands

The next reported attempt to classify the elements was made by the French geologist A. E. B. de Chancourtois in 1862. He arranged the then-known elements in order of increasing atomic weight and wrapped them around a cylinder, making a cylindrical table designed to display the periodic recurrence of properties. This clever idea, unfortunately, did not attract much attention.

A few years later, in 1865, the English chemist John Alexander Newlands put forward the Law of Octaves. He arranged the elements in increasing order of their atomic weights and noticed that every eighth element had properties similar to the first, as shown in Table 3.2.

The comparison he drew was musical. In the octaves of music, the eighth note resembles the first, and Newlands saw the same kind of repeat-after-eight pattern in the elements.

Table 3.2 Newlands’ Octaves

Element Li Be B C N O F
At. wt. 7 9 11 12 14 16 19
Element Na Mg Al Si P S Cl
At. wt. 23 24 27 29 31 32 35.5
Element K Ca
At. wt. 39 40

Newlands’ Law of Octaves seemed to hold true only for the elements up to calcium; beyond that, the pattern broke down. Although his idea was not widely accepted in his own time, his contribution was recognised much later, when he was awarded the Davy Medal in 1887 by the Royal Society, London.

When Newlands first presented his octaves to the Chemical Society, the reception was far from warm. His idea of a musical repetition among elements was met with ridicule, and the Society even declined to publish his paper. Yet the underlying insight — that properties recur periodically — was essentially correct. History eventually vindicated him: more than two decades later the same Royal Society honoured him with its prestigious Davy Medal for exactly this work.

Meyer and Mendeleev

The Periodic Law, as we understand it today, owes its development chiefly to two chemists working independently: the Russian chemist Dmitri Mendeleev (1834–1907) and the German chemist Lothar Meyer (1830–1895).

In 1869, both proposed that when elements are arranged in the increasing order of their atomic weights, similarities in their physical and chemical properties appear at regular intervals.

Lothar Meyer approached this graphically. He plotted physical properties such as atomic volume, melting point, and boiling point against atomic weight, and obtained a curve that rose and fell in a repeating, periodically recurring pattern. Unlike Newlands, who assumed a fixed repeat of eight, Meyer observed that the length of the repeating pattern actually changed as one moved through the elements — an important and more realistic insight.

By 1868, Meyer had developed a table of elements that closely resembles the Modern Periodic Table. However, his work was not published until after that of Mendeleev, who is therefore generally credited with the development of the Modern Periodic Table.

Mendeleev’s Periodic Law

While Döbereiner had initiated the study of periodic relationships, it was Mendeleev who was responsible for publishing the Periodic Law for the first time.

NotePrinciple / Law

The properties of the elements are a periodic function of their atomic weights. (This is Mendeleev’s original Periodic Law, based on atomic weights.)

Mendeleev arranged the elements in horizontal rows and vertical columns of a table, in order of their increasing atomic weights, in such a way that elements with similar properties fell into the same vertical column, or group.

His system was more elaborate than Meyer’s. Mendeleev fully recognised the significance of periodicity and used a broader range of physical and chemical properties to classify the elements. In particular, he relied heavily on the similarities in the empirical formulas and properties of the compounds that the elements formed. In other words, he did not judge an element by the element alone, but also by the kinds of oxides, hydrides, and chlorides it produced.

Gaps and Predictions

Mendeleev soon realised that some elements simply did not fit his scheme if the order of atomic weight was followed too strictly. His response was daring. He decided that the atomic weight measurements of the day might be slightly incorrect, and he chose to place elements with similar properties together even when this meant breaking the strict weight order.

For example, iodine has a lower atomic weight than tellurium, yet Mendeleev placed iodine after tellurium so that it would sit with the other similar elements — fluorine, chlorine, and bromine — in the same family, as reflected in his table (Fig. 3.2). He trusted chemical behaviour over a raw number.

Fig. 3.2

Fig. 3.2 – Mendeleev’s original Periodic Table.

Keeping his primary aim of grouping similar elements together, Mendeleev did something even bolder. He proposed that some elements were still undiscovered, and he deliberately left several gaps in his table for them. At the time, both gallium and germanium were unknown. He left a gap under aluminium and another under silicon, and named these missing elements Eka-Aluminium and Eka-Silicon.

The prefix “eka” comes from Sanskrit and means “one.” So “eka-aluminium” was Mendeleev’s way of saying “the element one place beyond aluminium,” and “eka-silicon” meant “one place beyond silicon.” Using an ordinary counting word, he gave temporary names to elements that no one had yet seen.

Mendeleev did not merely predict that these elements existed; he described some of their general physical properties in advance. When gallium and germanium were later discovered, the properties measured for them matched Mendeleev’s predictions with striking accuracy, as shown in Table 3.3.

Table 3.3 Mendeleev’s Predictions for the Elements Eka-aluminium (Gallium) and Eka-silicon (Germanium)

Property Eka-aluminium (predicted) Gallium (found) Eka-silicon (predicted) Germanium (found)
Atomic weight 68 70 72 72.6
Density / (g/cm³) 5.9 5.94 5.5 5.36
Melting point / K Low 302.93 High 1231
Formula of oxide \(\mathrm{E_2O_3}\) \(\mathrm{Ga_2O_3}\) \(\mathrm{EO_2}\) \(\mathrm{GeO_2}\)
Formula of chloride \(\mathrm{ECl_3}\) \(\mathrm{GaCl_3}\) \(\mathrm{ECl_4}\) \(\mathrm{GeCl_4}\)

The boldness of these quantitative predictions, and their eventual success, is what made Mendeleev and his Periodic Table famous. A version of Mendeleev’s Periodic Table, as published in 1905, is shown in Fig. 3.2.

Q. Why was Mendeleev’s willingness to leave gaps such an important step forward?

A. Leaving a gap turned the table from a mere summary of known elements into a tool that made testable predictions. If the predicted element was later found with the expected properties, it proved the classification was capturing something real about nature.

Taken together, these attempts show a clear line of progress, which is easiest to appreciate side by side.

Döbereiner’s Triads vs Newlands’ Octaves vs Mendeleev’s Periodic Table | Feature | Döbereiner’s Triads (1829) | Newlands’ Octaves (1865) | Mendeleev’s Periodic Table (1869) | |—|—|—|—| | Basis of arrangement | Small groups of three elements, by atomic weight | All elements in increasing atomic weight | Rows and columns, by increasing atomic weight | | Key idea | Middle element’s weight ≈ average of the other two; its properties also lie in between | Every eighth element resembles the first (like musical octaves) | Similar properties recur at regular intervals; similar elements share a group | | Main strength | First numerical hint of a pattern among properties | Recognised a genuinely repeating (periodic) pattern | Predicted undiscovered elements and their properties; grouped by chemistry, not just weight | | Main limitation | Worked for only a few sets of elements | Held only up to calcium | Based on atomic weight, which sometimes forced anomalies (e.g., iodine and tellurium) |

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

We must keep one important fact in mind: when Mendeleev built his Periodic Table, chemists knew nothing about the internal structure of the atom. They had no idea about protons, electrons, or the arrangement of electrons. Mendeleev’s ordering by atomic weight was, in a sense, the best that could be done at the time.

The beginning of the twentieth century, however, witnessed profound developments in the theories about sub-atomic particles. Once the atom’s inner structure came to light, the true basis of the Periodic Table could finally be uncovered.

Moseley and the Atomic Number

In 1913, the English physicist Henry Moseley studied the characteristic X-rays given off by different elements. Every element, when bombarded suitably, emits X-rays of a frequency that is characteristic of that element — a kind of fingerprint.

Moseley looked for a regular relationship between this X-ray frequency and the identity of the element. He found that a plot of \(\sqrt{\nu}\) (where \(\nu\) is the frequency of the X-rays emitted) against the atomic number \(Z\) gave a straight line. Crucially, plotting \(\sqrt{\nu}\) against atomic mass did not give such a clean straight line.

This was a decisive result. A property that varies in a simple, straight-line fashion with atomic number, but not with atomic mass, is telling us that the atomic number — not the atomic mass — is the more fundamental property of an element. The atomic number, you may recall, is equal to the nuclear charge, that is, the number of protons in the nucleus, and in a neutral atom it also equals the number of electrons.

Because of Moseley’s work, Mendeleev’s Periodic Law was accordingly modified. Atomic weight was replaced by atomic number as the ordering principle, giving us what we now call the Modern Periodic Law.

NotePrinciple / Law

The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Henry Moseley was only in his mid-twenties when he established the central role of the atomic number. His careful X-ray measurements even allowed him to correctly predict that certain elements were still missing from the table. Tragically, his promising career was cut short: he was killed in action during the First World War, in 1915, at Gallipoli. Had he lived, many believe he would have received a Nobel Prize for this fundamental contribution.

Q1. Mendeleev ordered elements by atomic weight, yet he deliberately placed iodine after tellurium even though iodine is the lighter of the two. Could there be a more fundamental property than weight that fixes such “misfits” automatically?

A. Yes — that property is the atomic number. Moseley showed the atomic number, not the atomic weight, is the true basis of periodicity. When the elements are arranged by atomic number, tellurium (\(Z = 52\)) naturally comes before iodine (\(Z = 53\)), even though iodine is lighter. The very anomaly that forced Mendeleev to break his own weight-order rule simply disappears once atomic number is used as the ordering principle.

Why Electronic Configuration is the Real Basis

Once we accept that atomic number governs periodicity, its deeper meaning becomes clear. Since the atomic number fixes the number of electrons in a neutral atom, it also fixes how those electrons are arranged — the electronic configuration.

It is now recognised that the Periodic Law is essentially the consequence of the periodic variation in electronic configurations. As we move through the elements, the pattern of filling electrons into shells and orbitals repeats itself, and it is this repeating pattern of configurations that indeed determines the physical and chemical properties of elements and their compounds.

The Modern Periodic Law revealed important analogies among the 94 naturally occurring elements. Even elements once thought purely artificial, such as neptunium and plutonium, are, like actinium and protoactinium, found in trace amounts in pitchblende — an ore of uranium. This law renewed interest in Inorganic Chemistry and has carried into the present day, extended further by the creation of artificially produced short-lived elements.

The Long Form of the Periodic Table

Numerous forms of the Periodic Table have been devised from time to time. Some forms emphasise chemical reactions and valence, while others stress the electronic configuration of elements. The modern version most widely used is the so-called “long form” of the Periodic Table, shown in Fig. 3.3. It is the most convenient and is the one we shall use throughout.

Fig. 3.3

Fig. 3.3 – The long form of the modern Periodic Table.

In this table, the horizontal rows — which Mendeleev had called series — are now called periods, and the vertical columns are called groups. Elements having similar outer electronic configurations in their atoms are placed one below another in the same vertical column, which is why groups are also called families.

Following the recommendation of the International Union of Pure and Applied Chemistry (IUPAC), the groups are numbered from 1 to 18. This simple numbering replaces the older, more confusing notation of groups IA to VIIA, VIII, IB to VIIB, and 0.

(question only — no solution) Elements were earlier grouped as IA to VIIA, VIII, IB to VIIB and 0. State the modern IUPAC group numbers that now correspond to (i) the old Group IA and (ii) the old Group 0 (the noble gases).

Periods and Groups

There are altogether seven periods. The period number corresponds to the highest principal quantum number (\(n\)) of the elements in that period. Here, the principal quantum number \(n\) is the number that labels the main energy level, or shell, of an electron.

The first period contains only 2 elements. The subsequent periods consist of 8, 8, 18, 18, and 32 elements, respectively. The seventh period is still incomplete; like the sixth period, it would have a theoretical maximum of 32 elements based on the number of available quantum numbers.

In this form of the Periodic Table, 14 elements each of the sixth and seventh periods — the lanthanoids and the actinoids, respectively — are placed in separate panels at the bottom. This is done purely to keep the main table compact and readable; otherwise the table would become inconveniently wide.

The period number of an element equals the highest principal quantum number \(n\) of its atoms. To which period does an element whose valence shell has \(n = 5\) belong, and how many elements does that period contain?

Since the period number equals the highest principal quantum number, an element with a valence shell of \(n = 5\) belongs to the fifth period. From the counts given above (2, 8, 8, 18, 18, 32 for periods one to six), the fifth period contains 18 elements.

The placing of the actinoids as a separate row below the lanthanoids is largely due to Glenn T. Seaborg. Starting with the discovery of plutonium around 1940, and then of all the transuranium elements from atomic number 94 to 102, Seaborg’s work forced a reconfiguration of the Periodic Table. He was awarded the Nobel Prize in Chemistry in 1951 for this contribution, and in his honour element 106 has been named Seaborgium (Sg).

3.4 Nomenclature of Elements with Atomic Numbers > 100

Why Naming Became a Problem

Traditionally, the naming of a newly discovered element was the privilege of its discoverer (or discoverers), and the suggested name was then ratified by IUPAC. In recent years, however, this simple tradition has led to some controversy.

The reason lies in the nature of the newest elements. Those with very high atomic numbers are so unstable — that is, so quickly radioactive — that only minute quantities, sometimes only a few atoms, can ever be obtained. Their synthesis and characterisation (confirming exactly which element has been made and measuring its properties) therefore require highly sophisticated and costly equipment and laboratories.

Such work is carried out with a competitive spirit in only a handful of laboratories around the world. Under this pressure, scientists have sometimes been tempted to claim a discovery even before collecting reliable data on the new element.

A famous case is element 104. Both American and Soviet scientists claimed the credit for discovering it. The Americans named it Rutherfordium, whereas the Soviets named it Kurchatovium — leaving the world with two names for one element.

The tug-of-war over element 104 was only one battle in what came to be nicknamed the “Transfermium Wars” — a decades-long naming rivalry, running through the Cold War, between the American laboratory at Berkeley and the Soviet laboratory at Dubna over the elements just beyond fermium. Each group claimed priority and proposed its own names. The disputes became so tangled that IUPAC finally had to step in with an impartial system, which is exactly the systematic nomenclature described below.

The IUPAC Systematic Names

To avoid such problems, IUPAC recommended that until a new element’s discovery is proved and its name is officially recognised, a systematic name should be derived directly from the atomic number of the element. This is done using fixed numerical roots for 0 and for the numbers 1 to 9, listed in Table 3.4.

Table 3.4 Notation for IUPAC Nomenclature of Elements

Digit Name Abbreviation
0 nil n
1 un u
2 bi b
3 tri t
4 quad q
5 pent p
6 hex h
7 sept s
8 oct o
9 enn e

The rule is simple: take the digits that make up the atomic number, replace each digit by its root, join the roots together in the same order as the digits, and add “ium” at the end. The three-letter symbol is formed by joining the abbreviations of the same roots. The full sequence of steps is summarised in FC 3.1.

Flowchart

FC 3.1 – IUPAC systematic naming: atomic number → split into individual digits → replace each digit with its numerical root (Table 3.4) → join the roots in order → add “ium” to get the temporary name; in parallel, join the digit abbreviations to get the three-letter symbol → later, IUPAC vote assigns the permanent name and symbol.

The systematic IUPAC names for the elements with \(Z\) above 100 are shown in Table 3.5, alongside the permanent official names that IUPAC has since assigned.

Table 3.5 Nomenclature of Elements with Atomic Number Above 100

Atomic Number Name according to IUPAC nomenclature Symbol IUPAC Official Name IUPAC Symbol
101 Unnilunium Unu Mendelevium Md
102 Unnilbium Unb Nobelium No
103 Unniltrium Unt Lawrencium Lr
104 Unnilquadium Unq Rutherfordium Rf
105 Unnilpentium Unp Dubnium Db
106 Unnilhexium Unh Seaborgium Sg
107 Unnilseptium Uns Bohrium Bh
108 Unniloctium Uno Hassium Hs
109 Unnilennium Une Meitnerium Mt
110 Ununnillium Uun Darmstadtium Ds
111 Unununnium Uuu Rontgenium Rg
112 Ununbium Uub Copernicium Cn
113 Ununtrium Uut Nihonium Nh
114 Ununquadium Uuq Flerovium Fl
115 Ununpentium Uup Moscovium Mc
116 Ununhexium Uuh Livermorium Lv
117 Ununseptium Uus Tennessine Ts
118 Ununoctium Uuo Oganesson Og

What would be the IUPAC name and symbol for the element with atomic number 120?

From Table 3.4, the roots for the digits 1, 2, and 0 are un, bi, and nil, respectively. Joining these roots in order and adding “ium” gives the name, while joining the abbreviations gives the symbol. Hence, the symbol and the name are respectively Ubn and unbinilium.

Write the systematic IUPAC name and symbol for the element with atomic number 119, the first element beyond those currently discovered.

The digits of 119 are 1, 1, and 9. From Table 3.4, the roots are un (u), un (u), and enn (e). Joining the roots and adding “ium” gives ununennium, and joining the abbreviations gives the symbol Uue.

(question only — no solution) Write the systematic IUPAC name and three-letter symbol for the element with atomic number 126.

From Temporary to Permanent Names

In this way, a newly made element first receives a temporary name, with a symbol consisting of three letters. Later, once its discovery is firmly established, a permanent name and symbol are decided by a vote of IUPAC representatives from each country.

The permanent name might reflect the country (or the state within a country) in which the element was discovered, or it may pay tribute to a notable scientist. You can see both kinds of tribute in Table 3.5 — for instance, Nihonium (after Japan) and Mendelevium (after Mendeleev).

As of now, elements with atomic numbers up to 118 have been discovered, and the official names of all these elements have been announced by IUPAC.

3.5 Electronic Configurations of Elements and the Periodic Table

In the preceding chapter you learnt that an electron in an atom is described by a set of four quantum numbers, and that the principal quantum number (\(n\)) defines the main energy level, known as a shell. You also studied how electrons are filled into different subshells, also called orbitals (\(s\), \(p\), \(d\), \(f\)), within an atom.

The distribution of electrons into the orbitals of an atom is called its electronic configuration. As a quick reminder, the four subshell types can hold different numbers of electrons: an \(s\) subshell holds up to 2, a \(p\) subshell up to 6, a \(d\) subshell up to 10, and an \(f\) subshell up to 14 electrons. These numbers will matter throughout this section.

An element’s location in the Periodic Table reflects the quantum numbers of the last orbital that was filled. In this section we will see the direct and beautiful connection between the electronic configurations of the elements and the long form of the Periodic Table.

Electronic Configurations Across Periods

The period tells us the value of \(n\) for the outermost, or valence, shell. In other words, each successive period in the Periodic Table is associated with the filling of the next higher principal energy level (\(n = 1\), then \(n = 2\), and so on).

A very useful rule follows from this: the number of elements in any period is exactly twice the number of atomic orbitals available in the energy level being filled. (The factor of two appears because each orbital holds two electrons, and each electron added means one more element.)

The First Three Periods

The first period (\(n = 1\)) begins with the filling of the lowest available level, \(1s\). Since there is only one \(1s\) orbital, it holds two electrons, so this period has just two elements — hydrogen (\(1s^1\)) and helium (\(1s^2\)) — after which the first shell (the K shell) is complete.

The second period (\(n = 2\)) starts with lithium, whose third electron enters the \(2s\) orbital. Beryllium, with four electrons, has the configuration \(1s^2\,2s^2\). From the next element, boron, the \(2p\) orbitals begin to fill, and the L shell is completed at neon (\(2s^2\,2p^6\)). Thus the second period contains 8 elements.

The third period (\(n = 3\)) begins at sodium, where the added electron enters a \(3s\) orbital. The successive filling of the \(3s\) and \(3p\) orbitals then gives a third period of 8 elements, from sodium to argon.

Why is the first shell labelled “K”? The letters K, L, M for the shells come from early X-ray spectroscopy. The physicist Charles Barkla, studying the X-rays that different elements emit, labelled the two main types of emission “K” and “L.” He deliberately started in the middle of the alphabet, leaving room in case more penetrating types were found closer to the nucleus. The naming stuck, and so the innermost shell has been “K” ever since.

Where the Transition Series Appear

The fourth period (\(n = 4\)) starts at potassium, and the added electrons first fill the \(4s\) orbital. Now something new happens. Before the \(4p\) orbital fills, the filling of the \(3d\) orbitals becomes energetically more favourable, and we meet the so-called \(3d\) transition series of elements.

This filling order — where \(4s\) fills before \(3d\) — is not a mistake; it happens because, at this stage, the \(4s\) orbital is actually slightly lower in energy than the \(3d\). The \(3d\) transition series starts from scandium (\(Z = 21\)), which has the configuration \(3d^1\,4s^2\), and the \(3d\) orbitals are completely filled at zinc (\(Z = 30\)), with configuration \(3d^{10}\,4s^2\). The fourth period then ends at krypton with the filling of the \(4p\) orbitals. Altogether there are 18 elements in this fourth period.

The fifth period (\(n = 5\)), beginning with rubidium, is similar to the fourth period and contains the \(4d\) transition series, which starts at yttrium (\(Z = 39\)). This period ends at xenon with the filling of the \(5p\) orbitals.

The Inner Transition Series

The sixth period (\(n = 6\)) contains 32 elements. Here, successive electrons enter the \(6s\), \(4f\), \(5d\), and \(6p\) orbitals, in that broad order. The filling of the \(4f\) orbitals begins with cerium (\(Z = 58\)) and ends at lutetium (\(Z = 71\)), giving the \(4f\)-inner transition series, which is called the lanthanoid series.

The seventh period (\(n = 7\)) is similar to the sixth, with the successive filling of the \(7s\), \(5f\), \(6d\), and \(7p\) orbitals, and it includes most of the man-made radioactive elements. This period will end at the element with atomic number 118, which would belong to the noble gas family. The filling of the \(5f\) orbitals, beginning after actinium (\(Z = 89\)), gives the \(5f\)-inner transition series, known as the actinoid series.

The overall order in which orbitals are filled as we build up the elements is worth having at a glance, and is shown in FC 3.2.

Flowchart

FC 3.2 – Order of orbital filling (aufbau order): 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p; mark 3d/4d/5d/6d as the transition-series entries and 4f/5f as the inner-transition (lanthanoid/actinoid) entries.

The \(4f\)- and \(5f\)-inner transition series are placed separately at the bottom of the Periodic Table. This is done to keep the table’s structure neat and, more importantly, to preserve the principle of classification by keeping elements with similar properties in a single column.

How would you justify the presence of 18 elements in the 5th period of the Periodic Table?

When \(n = 5\), the allowed values of the azimuthal quantum number are \(l = 0, 1, 2, 3\). The order in which the energy of the available orbitals \(4d\), \(5s\), and \(5p\) increases is \(5s < 4d < 5p\). The total number of orbitals available is 9 (one \(5s\), five \(4d\), and three \(5p\)). The maximum number of electrons that can be accommodated is therefore 18, and so there are 18 elements in the 5th period.

Using the same kind of reasoning, justify why the 6th period contains 32 elements.

As we move through the 6th period, electrons fill the \(6s\), \(4f\), \(5d\), and \(6p\) orbitals. Counting the orbitals: one \(6s\), seven \(4f\), five \(5d\), and three \(6p\), giving a total of \(1 + 7 + 5 + 3 = 16\) orbitals. Since each orbital holds 2 electrons, the maximum number of electrons is \(16 \times 2 = 32\). Hence the 6th period contains 32 elements.

(question only — no solution) Using the same orbital-counting reasoning, justify the number of elements present in the 4th period of the Periodic Table.

Group-wise Electronic Configurations

Elements in the same vertical column, or group, have similar valence shell electronic configurations. They possess the same number of electrons in their outer orbitals, and this is precisely why they show similar chemical properties.

For example, the Group 1 elements (the alkali metals) all have an \(ns^1\) valence shell electronic configuration, as shown below.

Atomic number Symbol Electronic configuration
3 Li \(1s^2\,2s^1\) (or) \([\mathrm{He}]\,2s^1\)
11 Na \(1s^2\,2s^2\,2p^6\,3s^1\) (or) \([\mathrm{Ne}]\,3s^1\)
19 K \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,4s^1\) (or) \([\mathrm{Ar}]\,4s^1\)
37 Rb \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,3d^{10}\,4s^2\,4p^6\,5s^1\) (or) \([\mathrm{Kr}]\,5s^1\)
55 Cs \(1s^2\,2s^2\,2p^6\,3s^2\,3p^6\,3d^{10}\,4s^2\,4p^6\,4d^{10}\,5s^2\,5p^6\,6s^1\) (or) \([\mathrm{Xe}]\,6s^1\)
87 Fr \([\mathrm{Rn}]\,7s^1\)

Notice that although lithium and caesium differ enormously in atomic mass, they share the same \(ns^1\) outer configuration — and that shared outer arrangement is what makes them all soft, reactive alkali metals.

Q2. Why should elements sitting in the same vertical column of the table — sometimes with very different weights — end up showing such similar chemical behaviour?

A. Because chemistry is decided mainly by the outermost (valence) electrons, and every element in a group has the same valence shell electronic configuration. The alkali metals, for example, all end in \(ns^1\): one loosely held outer electron. Whether the atom is small (lithium) or large (caesium), that identical outer arrangement makes them behave alike. The similarity comes from the electron pattern, not from the atomic weight.

Thus we can clearly see that the properties of an element have a periodic dependence upon its atomic number, and not upon its relative atomic mass.

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

The aufbau (build-up) principle, together with the electronic configuration of atoms, provides a solid theoretical foundation for the periodic classification. Recall that the aufbau principle simply says electrons fill the lowest-energy orbitals first as we build up an atom.

The elements in a vertical column of the Periodic Table make up a group, or family, and show similar chemical behaviour. This similarity arises because these elements have the same number, and the same distribution, of electrons in their outermost orbitals.

We can classify all the elements into four blocks — the s-block, p-block, d-block, and f-block — depending on the type of atomic orbital that is being filled with electrons. This block structure is illustrated in Fig. 3.4.

Fig. 3.4

Fig. 3.4 – The s-, p-, d-, and f-blocks of the Periodic Table.

There are two exceptions worth noting at once. Strictly, helium belongs to the s-block (its configuration is \(1s^2\)), yet it is positioned in the p-block along with the other Group 18 elements. This is justified because its completely filled valence shell (\(1s^2\)) makes it behave exactly like the other noble gases.

The other exception is hydrogen. With only one s-electron, it can be placed in Group 1 with the alkali metals. But it can also gain one electron to reach a noble-gas arrangement, which makes it resemble a Group 17 (halogen) element. Because it is such a special case, hydrogen is placed separately at the top of the Periodic Table, as shown in Fig. 3.3 and Fig. 3.4.

Q. If hydrogen has just one s-electron like the alkali metals, why isn’t it simply placed in Group 1?

A. Because hydrogen also resembles the halogens — it can gain one electron to reach a noble-gas (helium) arrangement, just as a halogen does. Since it fits partly in Group 1 and partly in Group 17, it is treated as a special case and shown on its own at the top of the table.

A quick side-by-side view of the four blocks makes their differences easy to hold in mind.

s-block vs p-block vs d-block vs f-block | Feature | s-block | p-block | d-block | f-block | |—|—|—|—|—| | Groups | 1 and 2 | 13 to 18 | 3 to 12 | lanthanoids and actinoids (bottom rows) | | Orbital being filled | outer \(ns\) | outer \(np\) | inner \((n-1)d\) | inner \((n-2)f\) | | General valence configuration | \(ns^{1-2}\) | \(ns^2np^{1-6}\) | \((n-1)d^{1-10}ns^{0-2}\) | \((n-2)f^{1-14}(n-1)d^{0-1}ns^2\) | | Overall nature | reactive metals | metals, non-metals, metalloids and noble gases | all metals (transition) | all metals (inner transition) | | Typical traits | form +1/+2 ions; mostly ionic compounds | full range of character; noble gases inert; halogens/chalcogens electron-hungry | coloured ions, variable oxidation states, catalysts, paramagnetism | many oxidation states (especially actinoids); actinoids radioactive |

We will now briefly discuss the salient features of these four types of elements. Some of the terms used here — such as ionization enthalpy and electron gain enthalpy — will be defined properly in Section 3.7; for now, ionization enthalpy can be read simply as a measure of how easily an atom loses an electron.

3.6.1 The s-Block Elements

The elements of Group 1 (the alkali metals) and Group 2 (the alkaline earth metals), which have \(ns^1\) and \(ns^2\) outermost electronic configurations respectively, belong to the s-Block Elements.

These are all reactive metals with low ionization enthalpies, meaning they give up their outer electrons easily. They lose the outermost electron(s) readily to form a 1+ ion (in the case of the alkali metals) or a 2+ ion (in the case of the alkaline earth metals).

The metallic character and the reactivity both increase as we go down the group. Because of this high reactivity, s-block elements are never found pure (in the free state) in nature; they always occur combined with other elements. The compounds of the s-block elements are predominantly ionic, with the important exception of those of lithium and beryllium, which show more covalent character.

The name “alkali” comes from an Arabic word for the ashes of certain plants. When these ashes (rich in sodium and potassium compounds) are dissolved in water, they give a basic, or alkaline, solution — hence the Group 1 metals are called the alkali metals, the “makers of alkali.”

Because alkali metals react so vigorously — potassium and caesium can react violently with even a trace of water, and sodium tarnishes instantly in air — they cannot be stored in the open. In the laboratory they are kept submerged under kerosene or mineral oil, which seals them off from air and moisture. This is a direct, practical consequence of their low ionization enthalpy and high reactivity.

(question only — no solution) Caesium belongs to Group 1 of the Periodic Table. State the charge of the ion it most readily forms, and predict whether its compounds are expected to be mainly ionic or mainly covalent, giving a reason.

3.6.2 The p-Block Elements

The p-Block Elements are those belonging to Groups 13 to 18. Together with the s-Block Elements, they are called the Representative Elements or Main Group Elements.

Their outermost electronic configuration varies from \(ns^2np^1\) to \(ns^2np^6\) across each period. At the end of every period sits a noble gas, with a closed valence shell of \(ns^2np^6\).

All the orbitals in the valence shell of the noble gases are completely filled with electrons, and it is very difficult to disturb this stable arrangement by adding or removing electrons. This is why the noble gases show very low chemical reactivity.

Q4. The noble gases sit quietly at the far right of the table and refuse to take part in most reactions. What is so special about their electron arrangement that makes them this unwilling to gain or lose electrons?

A. Their valence shell is completely filled (\(ns^2np^6\), or \(1s^2\) for helium). A full valence shell is an especially stable, low-energy arrangement. Removing an electron from it costs a great deal of energy, and adding one would force the electron into a much higher energy level. With no easy way to gain or lose electrons, the noble gases have little tendency to react — which is exactly why they are so unreactive.

Just before the noble gas family lie two chemically important groups of non-metals: the halogens (Group 17) and the chalcogens (Group 16). These two groups have highly negative electron gain enthalpies and readily add one or two electrons, respectively, to attain the stable noble-gas configuration.

Several p-block family names describe what the elements do. The “noble gases” are called noble because, like the noble metals, they hold themselves aloof and rarely react. The “halogens” take their name from Greek roots meaning “salt-formers,” since they combine with metals to make salts (such as common salt). The “chalcogens” are the “ore-formers,” a nod to the fact that many metal ores are oxides and sulphides of Group 16 elements.

Across a period, the non-metallic character increases as we move from left to right, while down a group the metallic character increases.

3.6.3 The d-Block Elements (Transition Elements)

These are the elements of Groups 3 to 12, sitting in the centre of the Periodic Table. They are characterised by the filling of inner \(d\) orbitals by electrons, and are therefore referred to as the d-Block Elements.

Their general outer electronic configuration is \((n-1)d^{1-10}ns^{0-2}\), with the exception of palladium (Pd), whose configuration is \(4d^{10}5s^0\). They are all metals. They mostly form coloured ions, exhibit variable valence (that is, several oxidation states), show paramagnetism (attraction to a magnetic field, arising from unpaired electrons), and are frequently used as catalysts.

However, zinc, cadmium, and mercury, which have the electronic configuration \((n-1)d^{10}ns^2\), do not show most of the properties of transition elements, because their \(d\) subshell is completely filled.

In a way, the transition metals form a bridge between the chemically active metals of the s-block and the less active elements of Groups 13 and 14 — and it is from this “in-between” position that they take their familiar name, “Transition Elements.”

The special properties of the d-block show up all around us. Their coloured compounds give many gemstones and pigments their hues; their variable oxidation states make them excellent catalysts in industry; and their strength and resistance to corrosion make them the basis of countless alloys, from stainless steel to the metals in coins.

(question only — no solution) The general outer electronic configuration of the d-block elements is \((n-1)d^{1-10}ns^{0-2}\). Explain why zinc, cadmium, and mercury are usually not regarded as typical transition elements.

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

The two rows of elements placed at the bottom of the Periodic Table are the Lanthanoids, from cerium, Ce (\(Z = 58\)), to lutetium, Lu (\(Z = 71\)), and the Actinoids, from thorium, Th (\(Z = 90\)), to lawrencium, Lr (\(Z = 103\)).

They are characterised by the outer electronic configuration \((n-2)f^{1-14}(n-1)d^{0-1}ns^2\). Because the last electron added to each of these elements enters an \(f\)-orbital, these two series are called the Inner-Transition Elements or f-Block Elements.

They are all metals. Within each series, the properties of the elements are quite similar to one another. The chemistry of the early actinoids is more complicated than that of the corresponding lanthanoids, owing to the large number of oxidation states possible for these actinoid elements.

The actinoid elements are radioactive. Many of them have been made only in nanogram quantities, or even less, through nuclear reactions, and their chemistry is therefore not yet fully studied. The elements that come after uranium are called the Transuranium Elements.

The elements \(Z = 117\) and \(Z = 120\) have not yet been discovered. In which family/group would you place these elements, and give the electronic configuration in each case.

From Fig. 3.3, the element with \(Z = 117\) would belong to the halogen family (Group 17), and its electronic configuration would be \([\mathrm{Rn}]\,5f^{14}6d^{10}7s^27p^5\). The element with \(Z = 120\) would be placed in Group 2 (the alkaline earth metals), and would have the electronic configuration \([\mathrm{Uuo}]\,8s^2\).

3.6.5 Metals, Non-metals and Metalloids

In addition to the classification into s-, p-, d-, and f-blocks, Fig. 3.4 shows another broad classification of elements based on their properties: they can be divided into Metals and Non-Metals.

Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. Metals are usually solids at room temperature — mercury is an exception (it is a liquid), while gallium and caesium also have very low melting points (\(303\ \text{K}\) and \(302\ \text{K}\), respectively). Metals usually have high melting and boiling points, are good conductors of heat and electricity, and are both malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires).

Non-metals, in contrast, are located at the top right-hand side of the Periodic Table. In fact, along any horizontal row, the character of the elements changes from metallic on the left to non-metallic on the right. Non-metals are usually solids or gases at room temperature, with low melting and boiling points (boron and carbon are exceptions). They are poor conductors of heat and electricity, and most non-metallic solids are brittle — neither malleable nor ductile.

The elements become more metallic as we go down a group, whereas the non-metallic character increases as we move from left to right across the Periodic Table.

metallic vs non-metallic character (across a period →: metallic character decreases and non-metallic character increases, because the effective nuclear charge rises and atoms hold their outer electrons more tightly, making electron loss harder; down a group ↓: metallic character increases and non-metallic character decreases, because atoms grow larger and their outer electrons are lost more easily.)

This change from metallic to non-metallic character is not abrupt. It is marked by the thick zig-zag line shown in Fig. 3.4. The elements that border this line and run diagonally across the table — such as silicon, germanium, arsenic, antimony, and tellurium — show properties characteristic of both metals and non-metals. These elements are called Semi-metals or Metalloids.

Considering the atomic number and position in the Periodic Table, arrange the following elements in the increasing order of metallic character: Si, Be, Mg, Na, P.

Metallic character increases down a group and decreases along a period as we move from left to right. Applying this, the increasing order of metallic character is: P < Si < Be < Mg < Na.

An element has the ground-state electronic configuration \([\mathrm{Ar}]\,3d^54s^1\). Identify its block, its period, and its group in the Periodic Table.

The last electrons are entering an inner \(d\) orbital, so the element is a d-block element. The highest principal quantum number present is \(n = 4\), so it lies in the 4th period. For a d-block element, the group number equals the total number of \((n-1)d\) and \(ns\) electrons: here \(5 + 1 = 6\), so the element belongs to Group 6. (This configuration is that of chromium, \(Z = 24\).)

3.8 Summary

In this Chapter, you studied the development of the Periodic Law and the Periodic Table, and the periodic trends that follow from the electronic configurations of atoms. The key ideas are recapped below.

How the Periodic Table developed

  1. The classification of elements grew from the work of many scientists. The overall journey — from the first hints of a pattern to the modern law — is condensed in FC 3.3.

Flowchart

FC 3.3 – Development of the Periodic Table: Döbereiner’s Triads (middle element ≈ average of the other two) → Newlands’ Law of Octaves (every 8th element similar, valid up to Ca) → Meyer & Mendeleev (properties recur at regular intervals with atomic weight; Mendeleev leaves gaps and predicts new elements) → Moseley (atomic number is more fundamental than atomic weight) → Modern Periodic Law (properties are periodic functions of atomic number).

  1. Mendeleev’s Periodic Table was based on atomic masses.
  2. The Modern Periodic Table arranges the elements in order of their atomic numbers, in seven horizontal rows (periods) and eighteen vertical columns (groups or families).
  3. Atomic numbers increase consecutively along a period, and in a regular pattern down a group.

Groups, periods, and blocks

  1. Elements of the same group have similar valence shell electronic configurations, and therefore show similar chemical properties.
  2. Elements of the same period have a steadily increasing number of electrons from left to right, and therefore have different valencies.
  3. On the basis of the orbital being filled, four types of elements are recognised: s-block, p-block, d-block, and f-block elements.
  4. Hydrogen, with a single electron in the \(1s\) orbital, occupies a unique position and is placed separately at the top of the table.

Metals, non-metals, and metalloids

  1. Metals make up more than 78% of the known elements.
  2. Non-metals, located towards the top right of the table, number fewer than twenty.
  3. Elements at the borderline between metals and non-metals (e.g., Si, Ge, As) are called metalloids or semi-metals.
  4. Metallic character increases down a group and decreases from left to right across a period.
  5. Overall, the physical and chemical properties of elements vary periodically with their atomic numbers.

Periodic trends in properties

  1. Periodic trends are seen in atomic size, ionization enthalpy, electron gain enthalpy, electronegativity, and valence.
  2. Atomic radii decrease from left to right across a period and increase down a group.
  3. Ionization enthalpies generally increase across a period and decrease down a group; electronegativity follows a similar trend.
  4. Electron gain enthalpies, in general, become more negative across a period and less negative down a group.
  5. Valence shows periodicity: for representative elements, the valence is either equal to the number of outermost electrons or to eight minus that number.

Chemical reactivity and the nature of oxides

  1. Chemical reactivity is highest at the two extremes of a period and lowest in the centre.
  2. Reactivity at the left extreme comes from the ease of electron loss (low ionization enthalpy); at the right extreme it comes from the ease of electron gain.
  3. Highly reactive elements do not occur free in nature; they are usually found in the combined form.
  4. Oxides of the elements on the left are basic, those on the right are acidic, and those in the centre are amphoteric or neutral.

3.9 NCERT Questions

  1. What is the basic theme of organisation in the periodic table?

  2. Which important property did Mendeleev use to classify the elements in his periodic table, and did he stick to that?

  3. What is the basic difference in approach between Mendeleev’s Periodic Law and the Modern Periodic Law?

  4. On the basis of quantum numbers, justify that the sixth period of the periodic table should have 32 elements.

  5. In terms of period and group, where would you locate the element with \(Z = 114\)?

  6. Write the atomic number of the element present in the third period and seventeenth group of the periodic table.

  7. Which element do you think would have been named by

    1. Lawrence Berkeley Laboratory,
    2. Seaborg’s group?

  8. Why do elements in the same group have similar physical and chemical properties?

  9. What does atomic radius and ionic radius really mean to you?

  10. How do atomic radii vary in a period and in a group? How do you explain the variation?

  11. What do you understand by isoelectronic species? Name a species that will be isoelectronic with each of the following atoms or ions:

    1. \(\mathrm{F^-}\)
    2. Ar
    3. \(\mathrm{Mg^{2+}}\)
    4. \(\mathrm{Rb^+}\).

  12. Consider the following species: \(\mathrm{N^{3-}}\), \(\mathrm{O^{2-}}\), \(\mathrm{F^-}\), \(\mathrm{Na^+}\), \(\mathrm{Mg^{2+}}\) and \(\mathrm{Al^{3+}}\).

    1. What is common in them?
    2. Arrange them in the order of increasing ionic radii.

  13. Explain why cations are smaller and anions larger in radii than their parent atoms.

  14. What is the significance of the terms ‘isolated gaseous atom’ and ‘ground state’ while defining ionization enthalpy and electron gain enthalpy? (Hint: requirements for comparison purposes.)

  15. The energy of an electron in the ground state of the hydrogen atom is \(-2.18\times10^{-18}\ \text{J}\). Calculate the ionization enthalpy of atomic hydrogen in terms of \(\text{J mol}^{-1}\). (Hint: apply the idea of the mole concept to derive the answer.)

  16. Among the second-period elements, the actual ionization enthalpies are in the order Li < B < Be < C < O < N < F < Ne. Explain why

    1. Be has higher \(\Delta_i H\) than B, and
    2. O has lower \(\Delta_i H\) than N and F.

  17. How would you explain the fact that the first ionization enthalpy of sodium is lower than that of magnesium, but its second ionization enthalpy is higher than that of magnesium?

  18. What are the various factors due to which the ionization enthalpy of the main group elements tends to decrease down a group?

  19. The first ionization enthalpy values (in \(\text{kJ mol}^{-1}\)) of Group 13 elements are:

B Al Ga In Tl
801 577 579 558 589

How would you explain this deviation from the general trend?

  1. Which of the following pairs of elements would have a more negative electron gain enthalpy?

    1. O or F
    2. F or Cl.

  2. Would you expect the second electron gain enthalpy of O to be positive, more negative, or less negative than the first? Justify your answer.

  3. What is the basic difference between the terms electron gain enthalpy and electronegativity?

  4. How would you react to the statement that the electronegativity of N on the Pauling scale is 3.0 in all nitrogen compounds?

  5. Describe the theory associated with the radius of an atom as it

    1. gains an electron,
    2. loses an electron.

  6. Would you expect the first ionization enthalpies for two isotopes of the same element to be the same or different? Justify your answer.

  7. What are the major differences between metals and non-metals?

  8. Use the periodic table to answer the following questions.

    1. Identify an element with five electrons in the outer subshell.
    2. Identify an element that would tend to lose two electrons.
    3. Identify an element that would tend to gain two electrons.
    4. Identify the group having metal, non-metal, liquid as well as gas at room temperature.

  9. The increasing order of reactivity among Group 1 elements is Li < Na < K < Rb < Cs, whereas that among Group 17 elements is F > Cl > Br > I. Explain.

  10. Write the general outer electronic configuration of s-, p-, d-, and f-block elements.

  11. Assign the position of the element having outer electronic configuration

    1. \(ns^2np^4\) for \(n = 3\),
    2. \((n-1)d^2ns^2\) for \(n = 4\), and
    3. \((n-2)f^7(n-1)d^1ns^2\) for \(n = 6\), in the periodic table.

  12. The first (\(\Delta_i H_1\)) and the second (\(\Delta_i H_2\)) ionization enthalpies (in \(\text{kJ mol}^{-1}\)) and the electron gain enthalpy (\(\Delta_{eg}H\), in \(\text{kJ mol}^{-1}\)) of a few elements are given below:

Element \(\Delta_i H_1\) \(\Delta_i H_2\) \(\Delta_{eg}H\)
I 520 7300 – 60
II 419 3051 – 48
III 1681 3374 – 328
IV 1008 1846 – 295
V 2372 5251 + 48
VI 738 1451 – 40

Which of the above elements is likely to be: a. the least reactive element; b. the most reactive metal; c. the most reactive non-metal; d. the least reactive non-metal; e. the metal which can form a stable binary halide of the formula \(\mathrm{MX_2}\) (X = halogen); f. the metal which can form a predominantly stable covalent halide of the formula \(\mathrm{MX}\) (X = halogen)?

  1. Predict the formulas of the stable binary compounds that would be formed by the combination of the following pairs of elements.
    1. Lithium and oxygen
    2. Magnesium and nitrogen
    3. Aluminium and iodine
    4. Silicon and oxygen
    5. Phosphorus and fluorine
    6. Element 71 and fluorine.
  2. In the modern periodic table, the period indicates the value of:
    1. atomic number
    2. atomic mass
    3. principal quantum number
    4. azimuthal quantum number.
  3. Which of the following statements related to the modern periodic table is incorrect?
    1. The p-block has 6 columns, because a maximum of 6 electrons can occupy all the orbitals in a p-shell.
    2. The d-block has 8 columns, because a maximum of 8 electrons can occupy all the orbitals in a d-subshell.
    3. Each block contains a number of columns equal to the number of electrons that can occupy that subshell.
    4. The block indicates the value of the azimuthal quantum number (\(l\)) for the last subshell that received electrons in building up the electronic configuration.
  4. Anything that influences the valence electrons will affect the chemistry of the element. Which one of the following factors does not affect the valence shell?
    1. Valence principal quantum number (\(n\))
    2. Nuclear charge (\(Z\))
    3. Nuclear mass
    4. Number of core electrons.
  5. The size of the isoelectronic species \(\mathrm{F^-}\), Ne, and \(\mathrm{Na^+}\) is affected by:
    1. nuclear charge (\(Z\))
    2. valence principal quantum number (\(n\))
    3. electron–electron interaction in the outer orbitals
    4. none of the factors, because their size is the same.
  6. Which one of the following statements is incorrect in relation to ionization enthalpy?
    1. Ionization enthalpy increases for each successive electron.
    2. The greatest increase in ionization enthalpy is experienced on removal of an electron from the core noble-gas configuration.
    3. The end of valence electrons is marked by a big jump in ionization enthalpy.
    4. Removal of an electron from an orbital bearing a lower \(n\) value is easier than from an orbital having a higher \(n\) value.
  7. Considering the elements B, Al, Mg, and K, the correct order of their metallic character is:
    1. B > Al > Mg > K
    2. Al > Mg > B > K
    3. Mg > Al > K > B
    4. K > Mg > Al > B.
  8. Considering the elements B, C, N, F, and Si, the correct order of their non-metallic character is:
    1. B > C > Si > N > F
    2. Si > C > B > N > F
    3. F > N > C > B > Si
    4. F > N > C > Si > B.
  9. Considering the elements F, Cl, O, and N, the correct order of their chemical reactivity in terms of oxidizing property is:
    1. F > Cl > O > N
    2. F > O > Cl > N
    3. Cl > F > O > N
    4. O > F > N > Cl.

3.10 Check Your Concepts

  1. Match each scientist in Column I with the correct contribution in Column II.
Column I Column II
a. Döbereiner i. Showed atomic number is more fundamental than atomic weight
b. Newlands ii. Law of Triads
c. Mendeleev iii. Placement of the actinoids below the lanthanoids
d. Moseley iv. Law of Octaves
e. Seaborg v. Left gaps and predicted eka-aluminium and eka-silicon

  1. Newlands’ Law of Octaves worked only for the lighter elements and broke down after calcium. Suggest a reason why an ordering rule based purely on a fixed repeat of eight could not hold for the heavier elements.

  2. Assertion: When elements are arranged by atomic number, tellurium is correctly placed before iodine, even though iodine is lighter. Reason: The atomic number, not the atomic weight, is the true basis of periodicity. State whether the assertion and reason are correct, and whether the reason correctly explains the assertion.

  3. Helium has the electronic configuration \(1s^2\), which would place it in the s-block. Give the reason why it is nevertheless positioned along with the Group 18 elements in the p-block.

  4. Hydrogen is placed separately at the top of the Periodic Table rather than firmly inside Group 1. Give two reasons that justify this special treatment.

  5. An element has the ground-state electronic configuration \([\mathrm{Ar}]\,3d^{10}4s^24p^3\). Identify its block, period, and group, and state whether it is a metal, non-metal, or metalloid.

  6. Give reasons:

    1. transition metals commonly show variable oxidation states;
    2. many transition-metal compounds are coloured.

  7. Explain why the second ionization enthalpy of an element is always greater than its first ionization enthalpy.

  8. Arrange Na, Mg, and Al in order of increasing first ionization enthalpy, and explain why the value for Al does not simply continue the trend set by Na and Mg.

  9. Explain, using size and charge/radius ratio, why lithium resembles magnesium (a diagonal relationship) more closely than it resembles sodium.

  10. State two reasons why boron, the first member of Group 13, differs in its chemical behaviour from the other members of the group.

  11. Assertion: The electron gain enthalpy of fluorine is less negative than that of chlorine. Reason: The added electron in fluorine enters a small \(n = 2\) shell where it suffers greater electron–electron repulsion. State whether the assertion and reason are correct, and whether the reason correctly explains the assertion.

  12. Classify each of the following oxides as basic, acidic, amphoteric, or neutral, and justify each answer using the position of the element in the Periodic Table: \(\mathrm{Na_2O}\), \(\mathrm{CO_2}\), \(\mathrm{Al_2O_3}\), \(\mathrm{N_2O}\).

  13. Complete and balance the following reactions, and state what each tells you about the nature of the oxide:

    1. \(\mathrm{K_2O} + \mathrm{H_2O} \rightarrow\)
    2. \(\mathrm{SO_3} + \mathrm{H_2O} \rightarrow\)

  14. Explain why the electronegativity of an element is not a fixed constant, unlike its atomic number.

  15. Arrange F, Cl, Br, and I in decreasing order of the magnitude of their electron gain enthalpy, and point out the one place where the trend does not follow the simple “decreases down the group” expectation.

3.11 Practice With Numericals

  1. The energy required to remove the most loosely bound electron from a single gaseous lithium atom is \(8.64\times10^{-19}\ \text{J}\). Calculate the first ionization enthalpy of lithium in \(\text{kJ mol}^{-1}\). (Take \(N_A = 6.022\times10^{23}\ \text{mol}^{-1}\).)

  2. The atomic weights of calcium and barium are 40 and 137, respectively. Using Döbereiner’s law of triads, predict the atomic weight of strontium, the middle member of this triad.

  3. Calculate the percentage by mass of oxygen in aluminium oxide, \(\mathrm{Al_2O_3}\). (Atomic masses: Al = 27, O = 16.)

  4. The first ionization enthalpy of magnesium is \(737\ \text{kJ mol}^{-1}\). Express this energy as the ionization energy per single magnesium atom, in joules. (Take \(N_A = 6.022\times10^{23}\ \text{mol}^{-1}\).)

  5. Determine the number of electrons in each of the species \(\mathrm{S^{2-}}\), \(\mathrm{Cl^-}\), and \(\mathrm{K^+}\), and state whether the three are isoelectronic. (Atomic numbers: S = 16, Cl = 17, K = 19.)

  6. The atomic (covalent) radius of bromine is 114 pm. Estimate the Br–Br bond length in a bromine molecule, \(\mathrm{Br_2}\).

  7. Using the orbital-counting method, calculate the maximum number of elements that would be present in the seventh period of the Periodic Table if it were complete. (The orbitals filled across this period are \(7s\), \(5f\), \(6d\), and \(7p\).)

  8. Calculate the percentage by mass of sodium in sodium oxide, \(\mathrm{Na_2O}\). (Atomic masses: Na = 23, O = 16.)