Chapter 7: Modern Periodic Table
In 1869, a Russian chemist named Dmitri Mendeleev was trying to bring order to a growing crowd of elements. Only 63 were known, and each one behaved in its own way. Mendeleev wrote their names and properties on cards and kept rearranging them, looking for a pattern the way one looks for a rhythm in a piece of music.
The pattern he found was this: when the elements were lined up in increasing order of atomic mass, similar properties kept reappearing at regular intervals. So he folded the line into rows, placing elements with similar behaviour one below the other.
But the folding did not work perfectly. To keep similar elements together, Mendeleev had to leave some boxes empty. Instead of treating these gaps as failures, he made a bold claim: the gaps belonged to elements nobody had discovered yet. He even described what they would be like — their approximate mass, their density, the kind of oxide they would form.
Within the next two decades, chemists isolated gallium and germanium. Their measured properties sat remarkably close to what Mendeleev had written for his empty boxes. A table built only from observed patterns had successfully predicted the unknown.
That table has since been rebuilt on a deeper foundation — not atomic mass, but atomic number — and it is this modern periodic table you are about to study.
Diagram
Fig. 7.0 – Mendeleev at his desk with element cards spread out, beside a simplified early periodic table showing two blank boxes labelled “eka-aluminium” and “eka-silicon” with arrows pointing to the modern symbols Ga and Ge
As you work through this chapter, you will gather exactly what you need to reason out each of these questions for yourself.
7.1 Introduction
In the early nineteenth century, only about 30 elements were known. Chemists studied their physical properties — appearance, lustre, hardness, ability to conduct heat and electricity — and sorted them into three types: metals, nonmetals and metalloids. Metals were shiny and good conductors, nonmetals were mostly dull and poor conductors, and metalloids sat awkwardly in between, showing some properties of each.
This first sorting was useful, but it was very coarse. Iron and sodium are both metals, yet they behave completely differently towards water. Clearly, a three-way split could not capture the real variety among elements.
As more elements were discovered, chemists needed a sharper organising idea. The property they turned to was atomic mass, because it was the one number that could be measured for every element and that clearly differed from element to element. Two noteworthy early attempts at classification were built on it: Dobereiner’s triads and Newlands’ octaves.
The First Patterns in Atomic Mass
Dobereiner noticed that certain elements could be grouped in sets of three — called triads — in which all three members showed similar chemical behaviour, and the atomic mass of the middle element was close to the average of the atomic masses of the other two. Lithium, sodium and potassium form such a set.
Newlands went further. On arranging the known elements in increasing order of atomic mass, he found that every eighth element resembled the first in its properties, rather as the eighth note of a musical scale repeats the first. He called this the law of octaves.
Both schemes worked only for small stretches of the element list, and both broke down as more elements were discovered. Their real importance is that they established a powerful idea: the properties of elements are not scattered at random, but repeat.
What “Periodic” Really Means
Before going further, it is worth being clear about the word periodic. An event is periodic if it repeats itself after a fixed interval. A full moon returns after a fixed number of days; a pendulum returns to the same position after a fixed time. The interval may be measured in days or in seconds, but the essential feature is regular repetition.
In chemistry, the “interval” is not time at all. It is the position of an element in an ordered list. When elements are arranged in a particular order, elements with similar properties reappear at regular intervals along that list. This repetition of properties is called periodicity, and it is the single idea on which the entire periodic table rests.
Mendeleev’s Periodic Table
In 1869, the Russian chemist Dmitri Mendeleev put forward a periodic table of the 63 elements known at that time, using both the atomic masses and the properties of the elements. His table was based on a statement now known as Mendeleev’s periodic law.
Mendeleev arranged the known elements in increasing order of their atomic masses. The serial or ordinal number of an element in this increasing order — simply its position number in the list — was referred to as its atomic number. At this stage the atomic number was only a counting label; it carried no deeper physical meaning.
He then folded this long list in accordance with the recurrence of properties. Wherever properties began to repeat, he started a new row and placed the repeating element directly below the one it resembled. The result was a table made of vertical groups and horizontal series, which we now call periods. The way he built it is summarised in FC 7.1.
Flowchart
FC 7.1 – Building Mendeleev’s periodic table: 63 known elements → arrange in increasing order of atomic mass → observe recurrence of similar properties → fold the list into rows so that similar elements fall one below the other → vertical columns = groups (similar properties), horizontal rows = series/periods (gradual variation) → leave gaps where no known element fits → predict properties of the missing elements; later branch: inert gases discovered afterwards → accommodated by adding an extra group
In Mendeleev’s table, elements belonging to the same group showed similar properties. Within a series or period, the properties of elements showed a gradual variation from left to right rather than an abrupt change.
To keep this pattern intact, Mendeleev left some boxes empty, corresponding to certain atomic numbers. He did this deliberately, so that the periodicity of properties would not be disturbed by forcing a mismatched element into a group where it did not belong.
Mendeleev’s periodic table was accepted by the scientific community precisely because newly discovered elements fitted well into the gaps, with properties matching those predicted by his periodic law. The inert gases, which Mendeleev had not predicted and which were discovered in later years, could also be accommodated in this table by creating an additional group for them.
From Atomic Mass to Atomic Number
Mendeleev’s law worked remarkably well, but it had an unresolved difficulty: for a small number of neighbouring pairs, ordering strictly by atomic mass placed an element in a group whose chemistry it clearly did not share. Mendeleev had to reverse such pairs and let chemical properties overrule atomic mass. This hinted that atomic mass, though useful, was not the true organising quantity.
The answer came only after the structure of the atom was understood. The atomic number, which had been just an ordinal label in Mendeleev’s table, was recognised as the proton number, \(Z\), of the element — the number of protons in its nucleus, and hence the size of the positive nuclear charge.
This recognition was the outcome of Henry Moseley’s work (1913) on the X-ray spectroscopic study of a large number of elements. When a solid element is bombarded with high-energy electrons, it emits X-rays whose frequency is characteristic of that element. Moseley showed that this frequency is related to the atomic number, \(Z\), rather than to the atomic mass. The frequency changed in a simple, regular way as he moved from element to element, and the quantity it tracked was the nuclear charge. The route from Mendeleev’s law to the modern law is set out in FC 7.2.
Flowchart
FC 7.2 – From atomic mass to atomic number: elements ordered by atomic mass → a few neighbouring pairs fall out of chemical order → structure of the atom established → Moseley’s X-ray study (1913) → X-ray frequency found to depend on Z, not on atomic mass → Z identified as the proton number and as the more fundamental property → periodic law restated in terms of Z → modern (long form) periodic table
Diagram
Fig. 7.1b – Moseley’s plot: √(X-ray frequency) on the y-axis against atomic number Z on the x-axis, showing a straight line with several elements (Ca, Ti, Cr, Fe, Cu, Zn) marked as points on it; alongside, a second sketch of the same quantity plotted against atomic mass giving a scattered, non-linear set of points
Because the atomic number, \(Z\), was now seen as a more fundamental property of the atom than the atomic mass, Mendeleev’s periodic law was modified accordingly. The revised statement is called the modern periodic law.
Mendeleev’s periodic table was then revised in accordance with the modern periodic law. We will look into the final revised version, known as the modern periodic table and also called the long form of the periodic table, in the sections that follow.
7.2 Structure of the Modern Periodic Table
Over a long period of time, many scientists have come up with different forms of the periodic table — spirals, circles, staircases and tables of several shapes. However, the so-called ‘long form of the periodic table’, or ‘the modern periodic table’, which is a revised version of Mendeleev’s periodic table, is the most convenient and the most widely used form of the periodic table of elements today.
Its convenience comes from its layout. Every element has one fixed address in it, and that address by itself tells a chemist a great deal about how the element will behave.
Periods and Groups
The modern periodic table has horizontal rows intersecting the vertical columns, giving rise to a number of boxes. The horizontal rows are called periods — these are what Mendeleev called series — and the vertical columns are called groups.
There are seven periods, numbered 1 to 7, and eighteen groups, numbered 1 to 18, in the modern periodic table. Each element sits in exactly one period and one group, so naming both fixes its position completely. The full table is shown in Fig. 7.1.
Diagram
Fig. 7.1 – Modern Periodic Table: all 118 elements arranged in 7 periods and 18 groups, each box showing atomic number, symbol, name and atomic mass; the two f-block series of 14 elements each placed below the main table; colour key for alkali metal, alkaline earth, transition metal, basic metal, metalloid, nonmetal, halogen, noble gas, lanthanoid and actinoid
This numbering of the periods and groups is recommended by the International Union of Pure and Applied Chemistry, IUPAC — the international body that settles the conventions, names and symbols that chemists everywhere agree to use. Older books number the groups differently, using Roman numerals along with the letters A and B, which is exactly the confusion that the IUPAC scheme of 1 to 18 was designed to remove.
The 118 Boxes
The boxes formed at the intersection of the periods and the groups are the places for individual elements. Below the main table are placed two series containing fourteen elements each, drawn separately so that the table stays a manageable width on a page.
Diagram
Fig. 7.2a – Two outlines of the periodic table side by side: (i) the usual layout, 18 columns wide, with the two 14-element series detached below; (ii) the fully extended layout, 32 columns wide, with the same two series inserted in their proper positions in periods 6 and 7 — arrows connecting the detached series to their true positions
There are in all 118 boxes to accommodate 118 elements in the modern periodic table. As on today, all the 118 boxes are filled, as a result of the discovery of manmade elements. IUPAC has approved the names and symbols of all the 118 elements. (Refer to Fig. 7.1.)
The heavier of these manmade elements do not occur in nature at all. They are made a few atoms at a time in particle accelerators, by firing one nucleus at another so that the two may fuse, and many of them survive for only a fraction of a second before decaying. Even so, once such an atom has been made and identified, it earns its own permanent box in the table.
The Four Blocks
The overall shape of the modern periodic table shows that it is divided into four blocks. Two groups on the left form the s-block, six groups on the right constitute the p-block, ten groups in the centre form the d-block, and the two series at the bottom constitute the f-block. These divisions are shown in Fig. 7.2.
Diagram
Fig. 7.2 – Outer electronic configuration of elements in the four blocks of the periodic table: the s-block (groups 1–2) on the left, the p-block (groups 13–18) on the right, the d-block (groups 3–12) in the centre and the f-block (lanthanoid and actinoid series) below, each block shaded differently and each element box carrying its outer electronic configuration
Notice that the four blocks are not scattered patches. Each one is a solid rectangle of the table, and together they account for every one of the 118 boxes. The reason behind this clean division, and the letters s, p, d and f themselves, will become clear in the next section.
7.3 Periodic Table and Electronic Configuration
When Mendeleev put forth his periodic table in 1869, the structure of the atom was not known. He observed periodicity in the properties of elements simply by arranging them in an increasing order of atomic mass. He could see the pattern, but he could not say what caused it.
Later, with the advent of the quantum mechanical model of the atom, the properties of elements were correlated to electronic configuration. The repetition that Mendeleev had spotted in the laboratory turned out to be a repetition in the way electrons arrange themselves around the nucleus.
Rules That Decide Where an Electron Goes
You have learnt in Chapter 4 that the electrons in an atom are distributed in shells and subshells in accordance with the aufbau principle, which includes the increasing order of energy, the Pauli exclusion principle and Hund’s rule of maximum multiplicity. Since the whole of this section rests on these three ideas, it is worth restating them before going on.
The principal quantum number \(n\) fixes the shell, and therefore the main energy level and the average distance of the electron from the nucleus. The azimuthal quantum number \(l\) fixes the subshell within that shell, and hence the shape of the orbital: \(l = 0\) is an s-subshell, \(l = 1\) a p-subshell, \(l = 2\) a d-subshell and \(l = 3\) an f-subshell.
These three rules taken together fix the electron capacity of every subshell. Since an s-subshell has one orbital, a p-subshell three, a d-subshell five and an f-subshell seven, and since each orbital can hold two electrons, the capacities are 2, 6, 10 and 14 respectively. You will use these four numbers again and again in this chapter.
When elements are arranged in an increasing order of atomic number (\(Z\)), periodicity is observed in their electronic configuration, and this in turn is reflected in the characteristic structure of the modern periodic table. The location of an element in the modern periodic table is correlated to the quantum numbers of the last filled orbital. Let us now have a deeper look into the electronic configuration of the elements and the structure of the modern periodic table.
7.3.1 Electronic Configuration in Periods
We noted earlier that the periods in the modern periodic table are numbered 1 to 7. On inspection of the electronic configurations of elements in various periods (see Fig. 7.3), we understand that the period number is the same as the principal quantum number ‘\(n\)’ of the outermost or valence shell of the elements.
Diagram
Fig. 7.3 – Electronic configurations of the elements arranged period by period, with the outermost shell highlighted in each row so that the period number can be seen to equal the principal quantum number n of the valence shell
This one statement is worth memorising, because it converts a question about position into a question about configuration. If the valence shell of an element is the fourth shell, the element must lie in the fourth period — no exceptions.
How a Period Begins and Ends
Along a period, the atomic number increases by one and one electron is added to the outermost shell, which forms the neutral atom of the next element. Every period ends with a complete octet configuration of the valence shell — or a duplet in the case of the first period — and the next period begins with the addition of an electron to the next shell of higher energy compared to the previous period.
The first shell, thus, gets filled along the first period. As the first shell can accommodate only two electrons, there are two elements in the first period, namely, H (\(Z=1\)) : \(1\mathrm{s}^1\) and He (\(Z=2\)) : \(1\mathrm{s}^2\). The first period ends at ‘He’ because ‘He’ has a complete duplet.
Electrons are filled in the second shell along the second period. The second period, thus, begins with Li (\(Z=3\)) : \(1\mathrm{s}^2\,2\mathrm{s}^1\) and ends up with Ne (\(Z=10\)) : \(1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6\). ‘Ne’, with 8 electrons in its outermost second shell, has a complete octet. The second shell has an electron capacity of 8. It gets filled along the second period as the atomic number increases. Thus there are eight elements in the second period.
Similarly there are eight elements, Na (\(Z=11\)) to Ar (\(Z=18\)), having the condensed electronic configurations described together as \([\mathrm{Ne}]3\mathrm{s}^{1\text{-}2}3\mathrm{p}^{1\text{-}6}\), in the third period, as a result of completion of the third shell. The square-bracket notation is simply a shorthand: \([\mathrm{Ne}]\) stands for the entire configuration of neon, \(1\mathrm{s}^2 2\mathrm{s}^2 2\mathrm{p}^6\), so that only the electrons added beyond the previous noble gas need to be written out.
Why the Fourth Period Is Longer
The fourth period begins with filling of the 4s subshell. The first two elements of the fourth period are K (\(Z=19\)) : \([\mathrm{Ar}]\,4\mathrm{s}^1\) and Ca (\(Z=20\)) : \([\mathrm{Ar}]4\mathrm{s}^2\).
According to the aufbau principle, the next higher energy subshell is 3d, which can accommodate up to 10 electrons. Filling of the 3d-subshell results in the next 10 elements of the fourth period, from Sc (\(Z=21\)) : \([\mathrm{Ar}]\,4\mathrm{s}^2 3\mathrm{d}^1\) to Zn (\(Z=30\)) : \([\mathrm{Ar}]\,4\mathrm{s}^2 3\mathrm{d}^{10}\).
After this, the electrons enter the subshell 4p for the next six elements : Ga (\(Z = 31\)) : \([\mathrm{Ar}]\,4\mathrm{s}^2 3\mathrm{d}^{10}4\mathrm{p}^1\) to Kr (\(Z=36\)) : \([\mathrm{Ar}]\,4\mathrm{s}^2 3\mathrm{d}^{10}4\mathrm{p}^6\). The fourth period, thus, contains in all 18 elements (\(2+10+6=18\)).
The fifth period accommodates 18 elements as a result of successive filling of electrons in the 5s, 4d and 5p subshells — the same pattern as the fourth period, one shell higher.
Flowchart
FC 7.3 – Aufbau order of filling of subshells: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p; note beneath each arrow the electron capacity (s = 2, p = 6, d = 10, f = 14)
Diagram
Fig. 7.3a – Bar chart of the number of elements in each period: period number 1 to 7 on the x-axis, number of elements on the y-axis, bars of height 2, 8, 8, 18, 18, 32, 32, with each bar split and labelled to show the contribution of the s, p, d and f subshells
Summing Up the Pattern in a Period
In short, a period begins by filling of one electron to the ‘s’ subshell of a new shell and ends with an element having a complete octet (or duplet) corresponding to the same shell. Between these two ends, corresponding to the ‘s’ and ‘p’ subshells of the valence shell, the inner subshells ‘d’ and ‘f’ are filled successively, following the aufbau principle.
This is why the d-block and f-block elements always appear in the middle of a long period, never at its two ends: they represent the filling of inner subshells while the outermost shell stays unchanged.
7.3.2 Electronic Configuration in Groups
A new shell is added down a group. The general outer electronic configuration, therefore, is expected to be the same down any particular group. Indeed it is found to be so for the groups 1, 2 and 3. (See Fig. 7.2 and Table 7.1.)
Look at group 1 in Table 7.1. Lithium has \(2\mathrm{s}^1\) and sodium has \(3\mathrm{s}^1\) — a different shell each time, but the same single electron in an s-subshell. Written in a general form, that is \(n\mathrm{s}^1\) for every member of the group.
In the groups 13 to 18 the appropriate inner ‘d’ and ‘f’ subshells are completely filled, and the general outer electronic configuration is the same down the groups 13 to 18 (see Fig. 7.2 and Table 7.1). A completely filled inner subshell does not disturb the outer configuration, because it is complete for every member of the group alike.
Table 7.1 : General outer electronic configuration in groups 1 to 3 and 13 to 18
| Group number | General outer configuration | Examples |
|---|---|---|
| Group 1 | \(n\mathrm{s}^1\) | \(_{3}\mathrm{Li}:2\mathrm{s}^1\), \(_{11}\mathrm{Na}:3\mathrm{s}^1\) |
| Group 2 | \(n\mathrm{s}^2\) | \(_{4}\mathrm{Be}:2\mathrm{s}^2\), \(_{12}\mathrm{Mg}:3\mathrm{s}^2\) |
| Group 3 | \(n\mathrm{s}^2\,(n-1)\mathrm{d}^1\) | \(_{21}\mathrm{Sc}:4\mathrm{s}^2 3\mathrm{d}^1\), \(_{39}\mathrm{Y}:5\mathrm{s}^2 4\mathrm{d}^1\) |
| Group 13 | \(n\mathrm{s}^2\,n\mathrm{p}^1\) | \(_{5}\mathrm{B}:2\mathrm{s}^2 2\mathrm{p}^1\), \(_{13}\mathrm{Al}:3\mathrm{s}^2 3\mathrm{p}^1\) |
| Group 14 | \(n\mathrm{s}^2\,n\mathrm{p}^2\) | \(_{6}\mathrm{C}:2\mathrm{s}^2 2\mathrm{p}^2\), \(_{14}\mathrm{Si}:3\mathrm{s}^2 3\mathrm{p}^2\) |
| Group 15 | \(n\mathrm{s}^2\,n\mathrm{p}^3\) | \(_{7}\mathrm{N}:2\mathrm{s}^2 2\mathrm{p}^3\), \(_{15}\mathrm{P}:3\mathrm{s}^2 3\mathrm{p}^3\) |
| Group 16 | \(n\mathrm{s}^2\,n\mathrm{p}^4\) | \(_{8}\mathrm{O}:2\mathrm{s}^2 2\mathrm{p}^4\), \(_{16}\mathrm{S}:3\mathrm{s}^2 3\mathrm{p}^4\) |
| Group 17 | \(n\mathrm{s}^2\,n\mathrm{p}^5\) | \(_{9}\mathrm{F}:2\mathrm{s}^2 2\mathrm{p}^5\), \(_{17}\mathrm{Cl}:3\mathrm{s}^2 3\mathrm{p}^5\) |
| Group 18 | \(n\mathrm{s}^2\,n\mathrm{p}^6\) | \(_{10}\mathrm{Ne}:2\mathrm{s}^2 2\mathrm{p}^6\), \(_{18}\mathrm{Ar}:3\mathrm{s}^2 3\mathrm{p}^6\) |
In the groups 4 to 12, however, the ‘d’ and ‘f’ subshells are introduced at a later stage — the 4th period for ‘d’ and the 6th period for ‘f’ — down the group. As a result, variation in the general outer configuration is introduced only at the later stage down the groups 4 to 12.
In other words, the members of these middle groups do not all start from the same point. The first member of such a group has no f-electrons behind it at all, while a later member does, and that difference shows up in the outer configuration.
7.3.3 Electronic Configuration in the Four Blocks
We noted in section 7.2 that the structure of the modern periodic table shows four blocks. These blocks are formed in accordance with the subshell in which the last electron enters. Accordingly the four blocks are named as s-block, p-block, d-block and f-block.
The rule is a simple one, and it is worth stating plainly before the details: find the subshell that receives the last electron in the aufbau sequence, and that letter names the block.
Flowchart
FC 7.4 – Locating an element from its electronic configuration: write the configuration by the aufbau order → identify the subshell receiving the last electron → that letter gives the BLOCK; branch: last electron in s → s-block, in p → p-block, in (n−1)d → d-block, in (n−2)f → f-block → read the highest principal quantum number n of the valence shell → that number gives the PERIOD → count the electrons in the outermost s and p subshells (for s- and p-block) or in ns + (n−1)d (for d-block) → that count gives the GROUP
s-Block
The last electron in the s-block elements is filled in an s-subshell. There being only one orbital in an s-subshell, the general outer electronic configuration of s-block elements is \(n\mathrm{s}^{1\text{-}2}\). Thus elements of the group 1 and group 2 belong to the s-block. The s-block is present on the left extreme of the modern periodic table.
p-Block
The last electron in the p-block elements is filled in a p-subshell. There being three degenerate p-orbitals in a p-subshell, up to 6 electrons can be filled. Therefore, the elements belonging to six groups, namely, group 13, 14, 15, 16, 17 and 18, constitute the p-block. The p-block appears on the right in the modern periodic table.
The p-block ends with the group 18, which represents the family of inert gases. Remarkably, the first element of the group 18, helium (He), does not have the p subshell, as its valence shell has \(n =1\) and its configuration is shown as \(1\mathrm{s}^2\). Yet ‘He’ is placed in the 18th group of the p-block, because its valence shell is completely filled (which is a complete duplet), similar to the complete valence shell of the other elements belonging to group 18 (which have complete octets).
The general electronic configuration for the p-block, from the second to the seventh period, is \(n\mathrm{s}^2\,n\mathrm{p}^{1\text{-}6}\).
d-Block
The d-block in the modern periodic table is formed as a result of filling the last electron in a d-orbital. A d-subshell is present in the shells with \(n \geq 3\), and according to the \((n+l)\) rule (refer to Chapter 4) the energy of the \(n\)s orbital is less than that of the \((n-1)\)d orbital. As a result, the last electron enters an \((n-1)\)d-subshell only after the \(n\)s subshell is completely filled.
The \((n+l)\) rule is the tool that settles such comparisons, so it is worth recalling in its exact form.
Applying it to 4s and 3d makes the point at once: for 4s, \(n+l = 4+0 = 4\); for 3d, \(n+l = 3+2 = 5\). The lower sum belongs to 4s, so 4s is filled first — which is exactly what we saw happening at the start of the fourth period.
There being five orbitals in a d-subshell, 10 electrons can successively be accommodated. There are ten groups, namely, groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 in the d-block, which appears in the centre of the modern periodic table.
The general outer electronic configuration of the d-block elements is \(n\mathrm{s}^{0\text{-}2}\,(n-1)\mathrm{d}^{1\text{-}10}\). Some variations in the configuration, consequent to the extra stability associated with a half-filled and a fully filled subshell, are readily observed. For example, the outer electronic configuration of Cr (\(Z = 24\)) is \(4\mathrm{s}^1 3\mathrm{d}^5\) instead of \(4\mathrm{s}^2 3\mathrm{d}^4\). This is because both the 4s and 3d subshells are half-filled.
f-Block
In the f-block elements the last electron is filled in an f-orbital. As there are seven orbitals in an f-subshell, the general outer electronic configuration of the f-block is \(n\mathrm{s}^2\,(n-1)\mathrm{d}^{0\text{-}1}\,(n-2)\mathrm{f}^{1\text{-}14}\).
The variations as a result of the extra stability of half-filled and fully filled subshells need to be accounted for here as well. For example, the outer electronic configuration of \(_{63}\mathrm{Eu}\) is \(6\mathrm{s}^2 4\mathrm{f}^7 5\mathrm{d}^0\), because the 4f subshell is half-filled.
The f-block constitutes two series of 14 elements called the lanthanoid and the actinoid series, put one below the other. The f-block is placed separately at the bottom of the periodic table.
7.4 Blockwise Characteristics of Elements
We have seen in the previous section that the 118 elements in the modern periodic table are distributed in four blocks, having a general electronic configuration according to their block. The elemental properties are characteristic of the block to which they belong.
This is the practical pay-off of everything in section 7.3. Once you know an element’s block, you already know a good deal about how it will behave — whether it is a metal, how readily it will give up or take in electrons, and what kind of compounds it is likely to form.
Flowchart
FC 7.5 – Identifying the block of an element from its observed properties: is the element a metal? → if no, it is a nonmetal or metalloid → p-block. If yes → does it react vigorously with cold water and form colourless, predominantly ionic compounds? → yes → s-block. If no → does it form coloured compounds and show more than one oxidation state? → yes → d-block. If it is a metal placed in one of the two separate series at the bottom of the table → f-block
7.4.1 Characteristics of s-Block Elements
The s-block contains the elements of group 1 (alkali metals) and group 2 (alkaline earth metals). All these elements are reactive metals, and occur in nature only in the combined state — that is, always as compounds, never as the free element.
Their compounds, with the exception of Li and Be, are predominantly ionic. This is because they have only one or two valence electrons, which they can lose readily, forming \(\mathrm{M^{+}}\) or \(\mathrm{M^{2+}}\) ions.
They have low ionization enthalpies, which decrease down the group, resulting in increased reactivity. Ionization enthalpy is simply the energy needed to pull an electron away from a gaseous atom; a low value means the atom parts with its outer electron easily. This quantity is taken up properly, with data, in section 7.5.
7.4.2 Characteristics of p-Block Elements
The p-block contains elements of groups 13 to 18. The p-block elements together with the s-block elements are called main group elements or representative elements, because between them they display the full range of behaviour that the periodic law was framed to describe.
The Noble Gases and the Reactive Nonmetals
The last group of the p-block, namely, the 18th group, is the family of noble/inert gases. These have closed valence shells — a complete duplet in the case of ‘He’ and a complete octet in the case of the other noble gases — and therefore very low chemical reactivity. An atom with a closed valence shell has nothing to gain, energetically, from either losing or acquiring an electron.
Elements of group 17 (the halogen family) and group 16 (the chalcogens) include reactive nonmetals. Their electron gain enthalpies are highly negative, which means that a large amount of energy is released when the atom accepts an electron. As a result, they gain one or two electrons readily and form anions (\(\mathrm{X^{-}}\) or \(\mathrm{X^{2-}}\)) which have a complete octet. Electron gain enthalpy, like ionization enthalpy, is treated quantitatively in section 7.5.
The Only Block with All Three Types of Element
The p-block contains all the three traditional types of elements. The metals lie on the left of the block, the nonmetals on the right, and the metalloids along a zig-zag line (see Fig. 7.1) which separates the metals from the nonmetals. No other block does this — the s-, d- and f-blocks are made up entirely of metals. The dividing line is drawn again, in isolation, in Fig. 7.4a.
Diagram
Fig. 7.4a – The p-block drawn alone, groups 13 to 18, with the zig-zag staircase line running from boron down to astatine; metals shaded on the left of the line, nonmetals on the right, and the metalloids (B, Si, Ge, As, Sb, Te) marked on the line itself
The nonmetallic character increases as we move from the left to the right, whereas it decreases as we go down a group.
7.4.3 Characteristics of d-Block Elements
The d-block contains elements of the groups 3 to 12. They are all metals. The d-block elements are known as transition elements or transition metals. They form a bridge between the chemically reactive s-block elements and the less reactive elements of groups 13 and 14 — which is exactly what the word “transition” is meant to convey about their position in the table.
Most of the d-block elements possess partially filled inner d-orbitals. As a result, the d-block elements have properties such as variable oxidation state, paramagnetism and the ability to form coloured ions. They can be used as catalysts.
Each of these four properties traces back to the same structural feature. Because the \(n\)s and \((n-1)\)d subshells are close in energy, electrons can be lost from either, so more than one oxidation state becomes available. Unpaired electrons in the partially filled d-subshell make the substance paramagnetic, that is, weakly attracted into a magnetic field. Those same partially filled d-orbitals allow the absorption of visible light that gives many of their ions their characteristic colours. And the availability of more than one oxidation state, together with vacant d-orbitals that can bind reacting molecules, is what allows these metals to act as catalysts.
Zn, Cd and Hg, with the configuration \(n\mathrm{s}^2\,(n-1)\mathrm{d}^{10}\) — that is, completely filled s- and d-subshells — do not show the properties typical of transition metals.
7.4.4 Characteristics of f-Block Elements
The f-block contains elements all of which are metals and are placed in the two rows called the lanthanoid series (\(_{58}\mathrm{Ce}\) to \(_{71}\mathrm{Lu}\)) and the actinoid series (\(_{90}\mathrm{Th}\) to \(_{103}\mathrm{Lr}\)).
These series are named after their preceding elements, lanthanum (\(_{57}\mathrm{La}\)) and actinium (\(_{89}\mathrm{Ac}\)), in the third group of the d-block of the 6th and 7th period respectively. The lanthanides are also known as rare earth elements.
The last electron of the elements of these series is filled in the \((n-2)\)f subshell, and therefore these are called inner-transition elements — the electron is entering a subshell two shells inside the valence shell, deeper still than the \((n-1)\)d subshell of the transition elements. These elements have very similar properties within each series.
Their close similarity follows directly from where that electron goes. Since the added electrons are buried two shells below the outermost one, the valence shell scarcely changes from one member of a series to the next, and chemical behaviour is governed by the valence shell. This is also why separating the individual lanthanoids from one another is a notoriously difficult task.
The actinide elements beyond \(_{92}\mathrm{U}\) are called transuranium elements. All the transuranium elements are manmade and radioactive.
7.5 Periodic Trends in Elemental Properties
The original structure of the periodic table was based on empirically observed periodicity in the elemental properties. Thus, elemental properties show similarity in a group and show gradual variation across a period. The quantum mechanical model of the atom explains the observed periodic trends of elemental properties.
We will now discuss in this section the periodic trends in some physical and chemical properties of elements, correlating them with electronic configuration and the nuclear charge.
These trends are explained in terms of two fundamental factors, namely, the attraction of the extranuclear electrons towards the nucleus, and the repulsion between electrons belonging to the same atom. These attractive and repulsive forces operate simultaneously in an atom. This results in two interrelated phenomena called effective nuclear charge and screening effect.
Everything that follows in this section — sizes, enthalpies, electronegativity, reactivity — comes back to these two ideas, so it is worth getting them clear first.
7.5.1 Effective Nuclear Charge and Screening Effect
In a multi-electron atom the positively charged nucleus attracts the negatively charged electrons around it, and there is mutual repulsion amongst the negatively charged extranuclear electrons.
The repulsion by inner-shell electrons is particularly important. This results in pushing the outer-shell electrons further away from the nucleus. The outer-shell electrons are, thus, held less tightly by the nucleus. In other words, the attraction of the nucleus for an outer electron is partially cancelled.
It means that an outer-shell electron does not experience the actual positive charge present on the nucleus. The net nuclear charge actually experienced by an electron is called the effective nuclear charge, \(Z_{\text{eff}}\). The effective nuclear charge is lower than the actual nuclear charge, \(Z\).
In other words, the inner electrons shield the outer electrons from the nucleus to a certain extent. This effect of the inner electrons on the outer electrons is called the screening effect or shielding effect of the inner/core electrons. The situation is drawn in Fig. 7.5c.
Diagram
Fig. 7.5c – A shell diagram of a sodium atom: nucleus marked +11 at the centre, ten core electrons drawn in the inner two shells, one valence electron in the outer shell; a bold arrow of attraction from nucleus to valence electron partly cancelled by repulsion arrows from the core electrons, with the resultant labelled Z(eff)
The relationship between the two is written as:
\[\text{Effective nuclear charge} = Z_{\text{eff}} = Z - \text{electron shielding} = Z - \sigma\]
Here \(\sigma\) (sigma) is called the shielding constant or screening constant, and the value of \(\sigma\) depends upon the type of the orbital that the electron occupies. An electron in an orbital that penetrates close to the nucleus is shielded less, and so it experiences a larger \(Z_{\text{eff}}\) than an electron in a less penetrating orbital of the same shell.
Across a Period
As we move across a period, the actual nuclear charge increases by +1 at a time, the valence shell remains the same, and the newly added electron gets accommodated in the same shell. There is no addition of electrons to the core.
Thus, shielding due to core electrons remains the same, though the actual nuclear charge, \(Z\), increases. The net result is that the effective nuclear charge, \(Z_{\text{eff}}\), goes on increasing across a period.
Down a Group
On the other hand, the \(Z_{\text{eff}}\) decreases down a group. This is because, as we move down a group, a new larger valence shell is added. As a result, there is an additional shell in the core.
The shielding effect of the increased number of core electrons outweighs the effect of the increased nuclear charge; and thereby the effective nuclear charge felt by the outer electrons decreases largely down a group.
7.5.2 Periodic Trends in Physical Properties
Many physical properties of elements, such as melting point, boiling point and density, show periodicity. In this section we are going to consider the periodic trends in physical properties with reference to atomic radius, ionic radius, ionization enthalpy, electron gain enthalpy and electronegativity.
a. Atomic Radius
You have learnt in Chapter 4 that the quantum mechanical model of the atom describes the extranuclear part of the atom as the electron cloud. As a direct implication of this, the atom has no definite boundary. There is no surface at which the electron density suddenly stops; it merely fades away with distance.
The atomic size or atomic radius, therefore, can be estimated from the internuclear distance under different circumstances. What can be measured accurately is not the edge of an atom but the distance between the centres of two atoms bonded together, and the radius is worked back from that.
In the case of nonmetals (except noble gases), the atoms of an element are bonded to each other by covalent bonds (refer to Chapter 5). The bond length of a single bond is taken as the sum of the radii of the two single bonded atoms. This is called the covalent radius of the atom. For example:
The bond length of the C–C bond in diamond is 154 pm. Therefore, the atomic radius of carbon is estimated to be 77 pm.
The bond length of the Cl–Cl bond in \(\mathrm{Cl_2}\) is measured as 198 pm. Therefore, the atomic radius of Cl is estimated to be 99 pm.
Both estimates are illustrated in Fig. 7.4.
Diagram
Fig. 7.4 – Estimation of atomic radii of metals and nonmetals: (upper) nine touching circles labelled Be in a close-packed arrangement with the centre-to-centre distance of adjacent atoms marked 224 pm and the working “atomic radius of Be = 224 pm ÷ 2 = 112 pm”; (lower) two overlapping circles labelled Cl forming the Cl₂ molecule with the internuclear distance marked 198 pm and the working “atomic radius of Cl = 198 pm ÷ 2 = 99 pm”
In the case of metals, the distance between the adjacent atoms in a metallic sample is measured. One half of this distance is taken as the metallic radius. Thus, atomic radius is one half of the internuclear distance between two adjacent atoms of a metal or two single bonded atoms of a nonmetal.
Note that the unit used throughout is the picometre, \(1\ \text{pm} = 10^{-12}\ \text{m}\), which is the convenient size scale for atoms.
Table 7.2 : Atomic radii of some elements
Element symbol (atomic radius / pm)
| Period Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| 2 | Li (152) | Be (111) | B (88) | C (77) | N (74) | O (66) | F (64) |
| 3 | Na (186) | Mg (160) | Al (143) | Si (117) | P (110) | S (104) | Cl (99) |
| 4 | K (231) | Br (114) | |||||
| 5 | Rb (244) | I (133) | |||||
| 6 | Cs (262) | At (140) |
It can be seen from Table 7.2 that atomic radius decreases across a period (upto group 17) and increases down a group.
Greater the effective nuclear charge, stronger is the attraction of the nucleus for the outer electrons, and smaller is the atomic radius.
As we move across a period, the screening effect caused by the core electrons remains the same; on the other hand the effective nuclear charge goes on increasing (see section 7.5.1). The valence electrons are, therefore, more tightly bound, and in turn the atomic radius goes on decreasing along a period.
The effective nuclear charge decreases and the shielding effect increases down a group (see section 7.5.1). The valence electrons are, thus, held by a weaker attractive force, and the atomic radius increases down a group.
b. Ionic Radius
An atom forms a positively charged ion, a cation, on the removal of one or more electrons, whereas a negatively charged ion, an anion, is formed with the gain of one or more electrons. Measurements of distances between neighbouring cations and anions in ionic crystals have been useful for the estimation of ionic radii. The ionic radii show the same trends as those of atomic radii.
A cation is smaller than the atom from which it is formed, because it contains fewer electrons than the atom, though the nuclear charge is the same. As a result the shielding effect is less and the effective nuclear charge is larger within a cation.
An anion has a larger radius than the corresponding atom, as it has a greater number of electrons than the atom. These additional electrons result in increased electron repulsion, decreased effective nuclear charge and, in turn, the increased size.
Some atoms and ions contain the same number of electrons and are called isoelectronic species (see Chapter 4). The actual nuclear charge of the isoelectronic species is, however, different.
The radii of isoelectronic species vary according to the actual nuclear charge. A larger nuclear charge exerts greater attraction for the electrons, and the radius of that isoelectronic species becomes smaller.
For example, \(\mathrm{F^{-}}\) and \(\mathrm{Na^{+}}\) both have 10 electrons; their radii are 133 pm and 98 pm respectively, as the nuclear charge of \(\mathrm{F^{-}}\) is \(+9\), which is smaller than that of \(\mathrm{Na^{+}}\), which is \(+11\).
c. Ionization Enthalpy
Removal of an electron from the neutral atom X results in the formation of the cation \(\mathrm{X^{+}}\). The energy required to remove an electron from the isolated gaseous atom in its ground state is called ionization enthalpy (\(\Delta_i H\)).
Ionization enthalpy is the quantitative measure of the tendency of an element to lose an electron, and is expressed in \(\text{kJ mol}^{-1}\). Since electrons are lost one at a time, we have first ionization enthalpy, second ionization enthalpy, and so on, for a given element:
\[\mathrm{X}\text{(g)} \rightarrow \mathrm{X^{+}}\text{(g)} + \mathrm{e^{-}} \qquad ; \ \Delta_i H_1\]
\[\mathrm{X^{+}}\text{(g)} \rightarrow \mathrm{X^{2+}}\text{(g)} + \mathrm{e^{-}} \qquad ; \ \Delta_i H_2\]
Ionization enthalpy is always positive, since energy always needs to be supplied to knock out an electron from an atom.
The second ionization enthalpy, \(\Delta_i H_2\), is larger than the first ionization enthalpy, \(\Delta_i H_1\), as it involves removal of an electron from the positively charged species. Pulling a negative electron away from a species that is already positive requires more work than pulling it from a neutral atom.
Tables 7.3 and 7.4 show the values of first ionization enthalpy down group 1 and across period 2 respectively.
Table 7.3 : First ionization enthalpy values of elements of group 1
| Elements of Group 1 | Atomic Number \(Z\) | Outer electronic configuration | \(\Delta_i H_1\) / kJ mol\(^{-1}\) |
|---|---|---|---|
| Li | 3 | \(2\mathrm{s}^1\) | 520 |
| Na | 11 | \(3\mathrm{s}^1\) | 496 |
| K | 19 | \(4\mathrm{s}^1\) | 419 |
| Rb | 37 | \(5\mathrm{s}^1\) | 403 |
| Cs | 55 | \(6\mathrm{s}^1\) | 374 |
It is seen from Table 7.3 that on moving down the group the first ionization enthalpy decreases. This is because the electron is to be removed from the larger valence shell. Screening due to core electrons goes on increasing and the effective nuclear charge decreases down the group (see section 7.5.1). The removal of the outer electron, therefore, becomes easier.
Table 7.4 : First ionization enthalpy values of elements of period 2
| Element of period 2 | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| Atomic number \(Z\) | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 |
| Outer electronic configuration | \(2\mathrm{s}^1\) | \(2\mathrm{s}^2\) | \(2\mathrm{s}^2 2\mathrm{p}^1\) | \(2\mathrm{s}^2 2\mathrm{p}^2\) | \(2\mathrm{s}^2 2\mathrm{p}^3\) | \(2\mathrm{s}^2 2\mathrm{p}^4\) | \(2\mathrm{s}^2 2\mathrm{p}^5\) | \(2\mathrm{s}^2 2\mathrm{p}^6\) |
| \(\Delta_i H_1\) / kJ mol\(^{-1}\) | 520 | 899 | 801 | 1086 | 1402 | 1314 | 1681 | 2080 |
An overall increase of the first ionization enthalpy across period 2 (see Table 7.4) can be noticed. This is because the screening is the same and the effective nuclear charge increases across a period (see section 7.5.1). As a result the outer electron is held more tightly. The first ionization enthalpy, therefore, increases across a period. The alkali metal displays the lowest first ionization enthalpy. The inert gas shows the highest first ionization enthalpy across a period.
The Two Irregularities in Period 2
Some irregularities are noticed for the first ionization enthalpies as we move across a period. Reading Table 7.4 carefully, the values do not climb smoothly from Li to Ne — they dip twice.
For example, the first ionization enthalpy of ‘B’ is smaller than that of ‘Be’. This is because ‘Be’ loses the electron from a 2s orbital while ‘B’ loses the electron from a 2p orbital, which has less penetration (see shapes of orbitals in Chapter 4) than the 2s orbital, and therefore it is easier to remove a 2p electron than a 2s electron. A less penetrating orbital keeps its electron farther from the nucleus on average and is better shielded, so that electron is held more loosely.
Similarly, the first ionization enthalpy of ‘O’ is smaller than that of ‘N’. This is because ‘O’ loses the electron from a doubly occupied ‘2p’ orbital. Due to electron–electron repulsion, it is easier to lose this electron than an electron from the singly occupied 2p orbital in the nitrogen atom. (This is a consequence of Hund’s rule of maximum multiplicity, described in Chapter 4.)
Diagram
Fig. 7.5a – Graph of first ionization enthalpy (kJ mol⁻¹) on the y-axis against atomic number Z from 3 to 10 on the x-axis, points joined by a line, each point labelled with its element symbol Li, Be, B, C, N, O, F, Ne; the two dips at B and at O circled and marked “irregularity”
d. Electron Gain Enthalpy
Addition of an electron to a neutral atom (X) results in the formation of an anion (\(\mathrm{X^{-}}\)). The enthalpy change that takes place when an electron is added to an isolated gaseous atom in its ground state is called the electron gain enthalpy, \(\Delta_{eg}H\).
Electron gain enthalpy is a quantitative measure of the ease with which an atom adds an electron, forming the anion, and is expressed in units of \(\text{kJ mol}^{-1}\):
\[\mathrm{X}\text{(g)} + \mathrm{e^{-}} \rightarrow \mathrm{X^{-}}\text{(g)} \qquad ; \ \Delta_{eg}H\]
\(\Delta_{eg}H\) may be positive or negative, depending upon whether the process of adding an electron is endothermic or exothermic. This is the crucial difference from ionization enthalpy, which has only one possible sign.
Elements of group 17 have very high negative values of electron gain enthalpy. This is because they attain a stable noble gas electronic configuration on the addition of one electron.
Elements of group 18, the noble gases, on the other hand, exhibit high positive values for electron gain enthalpy. Since the added electron has to enter the next higher shell with a larger principal quantum number, this is a very unstable electronic configuration, due to very low effective nuclear charge and high shielding from the core electrons.
The alkali metals have very low negative electron gain enthalpy values. In general, electron gain enthalpies are large negative for elements of the upper right of the periodic table, excluding the group 18 of noble gases.
The trends for electron gain enthalpy values of the elements in the periodic table are less regular than the ionization enthalpies. An overall trend reveals that electron gain enthalpy becomes more negative with increase of atomic number along the period upto the second last element, because the effective nuclear charge increases across the period and it is easier to add an electron to a smaller atom.
Down the group, the electron has to be added to a farther shell, and the electron gain enthalpy, thus, becomes less negative.
e. Electronegativity
When two atoms of different elements form a covalent bond, the electron pair is shared unequally. The ability of a covalently bonded atom to attract the shared electrons toward itself is called electronegativity (\(EN\)).
It is not an experimentally measurable quantity. Unlike ionization enthalpy and electron gain enthalpy, which are energies that can be measured directly, electronegativity has to be assigned on a scale. A number of numerical scales of electronegativity were developed by many scientists. The Pauling scale of electronegativity is the one used most widely.
Linus Pauling assigned (1922) arbitrarily a value of 4.0 for fluorine, which is expected to have the highest value of electronegativity. Table 7.5 shows the values of electronegativity assigned to some other elements.
Table 7.5 : Electronegativity (\(EN\)) values of some elements (Pauling scale)
Element symbol (\(EN\))
| Period Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| 1 | H (2.1) | ||||||
| 2 | Li (1.0) | Be (1.5) | B (2.0) | C (2.5) | N (3.0) | O (3.5) | F (4.0) |
| 3 | Na (0.9) | Mg (1.2) | Al (1.5) | Si (1.8) | P (2.1) | S (2.5) | Cl (3.0) |
| 4 | K (0.8) | Br (2.8) | |||||
| 5 | Rb (0.8) | I (2.5) | |||||
| 6 | Cs (0.7) | At (2.2) |
Electronegativity represents the attractive force exerted by the nucleus on shared electrons. Electron sharing between covalently bonded atoms takes place using the valence electron. The electronegativity depends upon the effective nuclear charge experienced by the electron involved in the formation of the covalent bond.
Electronegativity increases as we move across the period. This is because the effective nuclear charge increases steadily across the period.
The \(EN\) decreases down the group. The size of the valence shell goes on increasing, the shielding effect of the core electrons goes on increasing, and in turn the effective nuclear charge decreases down the group.
Electronegativity predicts the nature of the bond, or, how strong is the force of attraction that holds two atoms together (refer to Chapter 5). A large difference in electronegativity between two bonded atoms means the shared pair sits much closer to one of them, giving the bond a strongly polar, ionic character; a small difference gives an essentially nonpolar covalent bond.
7.5.3 Periodic Trends in Chemical Properties
The most fundamental chemical property of an element is its combining power. This property is numerically expressed in terms of valency or valence. Valency of an element indicates the number of chemical bonds that the atom can form, giving a molecule.
Another frequently used term related to valence is the oxidation state or oxidation number. Valence does not have any sign associated with it, but oxidation number does, and can be either + or –, which is decided by the electronegativities of the atoms that are bonded. The more electronegative atom of a bonded pair is assigned the negative oxidation number, since it is the one drawing the shared electrons towards itself.
The second aspect of chemical property is the chemical reactivity. The chemical reactivity is related to the ease with which an element loses or gains the electrons.
The chemical properties of elements are related to electronic configuration. This single statement is what ties the whole of section 7.5 together, and the chain of cause and effect is set out in FC 7.6.
Flowchart
FC 7.6 – Chain of causation behind every periodic trend: change in atomic number → change in nuclear charge Z and in the number of core shells → change in effective nuclear charge Z(eff) → change in atomic and ionic radius → change in ionization enthalpy, electron gain enthalpy and electronegativity → change in metallic/nonmetallic character → change in chemical reactivity and in the nature of the oxides formed
Table 7.6 : Periodic trends in valency of main group elements
| Group | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| General outer electronic configuration | \(n\mathrm{s}^1\) | \(n\mathrm{s}^2\) | \(n\mathrm{s}^2 n\mathrm{p}^1\) | \(n\mathrm{s}^2 n\mathrm{p}^2\) | \(n\mathrm{s}^2 n\mathrm{p}^3\) | \(n\mathrm{s}^2 n\mathrm{p}^4\) | \(n\mathrm{s}^2 n\mathrm{p}^5\) | \(n\mathrm{s}^2 n\mathrm{p}^6\) |
| Number of valence electrons | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 |
| Valency | 1 | 2 | 3 | 4 | 3, 5 | 2, 6 | 1, 7 | 0 |
| Formula of hydride | \(\mathrm{LiH}\) | \(\mathrm{BeH_2}\) | \(\mathrm{B_2H_6}\) | \(\mathrm{CH_4}\) | \(\mathrm{NH_3}\) | \(\mathrm{H_2O}\) | \(\mathrm{HF}\) | |
| \(\mathrm{NaH}\) | \(\mathrm{MgH_2}\) | \(\mathrm{AlH_3}\) | \(\mathrm{SiH_4}\) | \(\mathrm{PH_3}\) | \(\mathrm{H_2S}\) | \(\mathrm{HCl}\) | ||
| \(\mathrm{KH}\) | \(\mathrm{CaH_2}\) | \(\mathrm{GaH_3}\) | \(\mathrm{GeH_4}\) | \(\mathrm{AsH_3}\) | \(\mathrm{H_2Se}\) | \(\mathrm{HBr}\) | ||
| Formula of oxide | \(\mathrm{Li_2O}\) | \(\mathrm{BeO}\) | \(\mathrm{B_2O_3}\) | \(\mathrm{CO_2}\) | \(\mathrm{N_2O_3}\), \(\mathrm{N_2O_5}\) | \(\mathrm{OF_2}\) | ||
| \(\mathrm{Na_2O}\) | \(\mathrm{MgO}\) | \(\mathrm{Al_2O_3}\) | \(\mathrm{SiO_2}\) | \(\mathrm{P_4O_6}\), \(\mathrm{P_2O_5}\) | \(\mathrm{SO_2}\), \(\mathrm{SO_3}\) | \(\mathrm{Cl_2O_7}\) | ||
| \(\mathrm{K_2O}\) | \(\mathrm{CaO}\) | \(\mathrm{Ga_2O_2}\) | \(\mathrm{GeO_2}\) | \(\mathrm{As_2O_3}\), \(\mathrm{As_2O_5}\) | \(\mathrm{SeO_2}\), \(\mathrm{SO_3}\) | \(\mathrm{Br_2O}\) |
a. Periodic Trends in Valency
Valency of the main group elements is usually equal to the number of valence electrons (outer electrons) and/or equal to the difference between 8 and the number of valence electrons. Table 7.6 shows the periodic trends in the valency of main group elements, by taking examples of hydrides and some oxides.
The two rules correspond to the two ways an atom can complete an octet. An atom with few valence electrons finds it easier to use or lose all of them, so its valency equals the number of valence electrons. An atom close to a complete octet finds it easier to gain or share the few electrons still missing, so its valency equals \(8 -\) (number of valence electrons). This is why elements in the middle and right of a period often show two valencies, as Table 7.6 records for groups 15, 16 and 17.
As may be noticed from Table 7.6 that the valency remains the same down the group, and shows a gradual variation across the period as the atomic number increases from left to right.
b. Periodic Trends in Metallic–Nonmetallic Character
The metals, nonmetals and metalloids appear in separate regions of the modern periodic table: metals on the left, nonmetals on the right and metalloids along a zig-zag line separating the two.
Metals are characterized by good electrical conductivity and the ability to form compounds by loss of valence electrons. Nonmetals are characterized by their poor electrical conductivities and the ability to form compounds by gain of valence electrons in the valence shell. This can be explained in terms of ionization enthalpy and electron gain enthalpy.
The ionization enthalpy decreases down the group. The tendency to lose valence electrons, thus, increases down the group, and the metallic character increases down a group. The ionization enthalpy increases across the period, and consequently metallic character decreases across a period.
Electron gain enthalpy becomes more and more negative across the period, and it tends to be less negative down the group. Thus, nonmetallic character increases across the period and decreases down the group.
Diagram
Fig. 7.5b – A schematic periodic table outline with four labelled arrows summarising the trends: across a period → atomic radius decreases, ionization enthalpy increases, electronegativity increases, metallic character decreases; down a group ↓ atomic radius increases, ionization enthalpy decreases, electronegativity decreases, metallic character increases
c. Periodic Trends in Chemical Reactivity
Chemical reactivity of elements is decided by how easily it attains the electronic configuration of the nearest inert gas by gaining or losing electrons.
The elements preceding an inert gas react by gaining electrons in the outermost shell, whereas the elements which follow an inert gas in the periodic table react by loss of valence electrons.
Thus the chemical reactivity is decided by the electron gain enthalpy and ionization enthalpy values, which in turn are decided by effective nuclear charge and finally by the atomic size.
The ionization enthalpy is smallest for the element on the extreme left in a period, whereas the electron gain enthalpy is most negative for the second last element on the extreme right (preceding the inert gas, which is the last element of a period). Thus, the most reactive elements lie on the extreme left and the extreme right (excluding inert gases) of the periodic table.
The Nature of the Oxides
Apart from the metallic–nonmetallic character, the chemical reactivity can be illustrated by comparing the reaction of elements with oxygen to form oxides, and the nature of the oxides.
The reactive elements on the extreme left (that is, alkali metals) react vigorously with oxygen to form oxides (for example \(\mathrm{Na_2O}\)), which react with water to form strong bases (like \(\mathrm{NaOH}\)).
The reactive elements on the right (that is, halogens) react with oxygen to form oxides (for example \(\mathrm{Cl_2O_7}\)), which on reaction with water form strong acids (like \(\mathrm{HClO_4}\)).
The oxides of the elements in the centre of the main group elements are either amphoteric (for example \(\mathrm{Al_2O_3}\)), neutral (for example \(\mathrm{CO}\), \(\mathrm{NO}\)) or weakly acidic (for example \(\mathrm{CO_2}\)). An amphoteric oxide is one that reacts with an acid as well as with a base, behaving as a base towards the one and as an acid towards the other.
Flowchart
FC 7.7 – Nature of the oxides across a period: extreme left (alkali metals) → basic oxides, e.g. Na₂O → with water give strong bases, e.g. NaOH → moving right, centre of the period → amphoteric oxides, e.g. Al₂O₃ (react with both acid and base), and neutral oxides, e.g. CO, NO → further right → weakly acidic oxides, e.g. CO₂ → extreme right (halogens) → strongly acidic oxides, e.g. Cl₂O₇ → with water give strong acids, e.g. HClO₄
The Middle of the Table
The change in atomic radii of transition metals and inner transition elements is rather small. Therefore, the transition metals and inner transition elements belonging to the individual series have similar chemical properties. Their ionization enthalpies are intermediate between those of the s-block and the p-block.
The reason for the small change in radius is the one already met in section 7.3: along these series the electrons are being added to an inner d or f subshell, not to the valence shell. The added inner electrons shield the outer ones effectively, so the net pull on the valence shell hardly changes from one member of the series to the next — and neither, therefore, does the size or the chemistry.
7.6 Summary
In the early nineteenth century about 30 elements were known, and they were classified as metals, nonmetals and metalloids on the basis of their physical properties. As the number of elements grew, classification was attempted on the basis of atomic mass through Dobereiner’s triads and Newlands’ octaves.
In 1869, Dmitri Mendeleev arranged the 63 known elements in increasing order of atomic mass and folded the list wherever properties recurred, producing a table of vertical groups and horizontal series (now called periods). Mendeleev’s periodic law: the physical and chemical properties of elements are a periodic function of their atomic masses.
Elements of the same group in Mendeleev’s table showed similar properties, while properties varied gradually along a series. He deliberately left gaps for undiscovered elements and predicted their properties; the later discovery of those elements, and the accommodation of the inert gases in an additional group, won his table wide acceptance.
Henry Moseley’s work (1913) on X-ray spectra showed that the frequency of X-rays emitted by an element depends on its atomic number, \(Z\), and not on its atomic mass. The atomic number was recognised as the proton number and as the more fundamental property of the atom. Modern periodic law: the physical and chemical properties of elements are a periodic function of their atomic numbers.
The revised table is the modern periodic table, also called the long form of the periodic table. It has seven periods (1 to 7) and eighteen groups (1 to 18), numbered as recommended by IUPAC, with two series of fourteen elements each placed below the main table. It has 118 boxes for 118 elements, all of which are now filled, partly through the discovery of manmade elements, and all of whose names and symbols have been approved by IUPAC.
The table is divided into four blocks: the s-block (two groups on the left), the p-block (six groups on the right), the d-block (ten groups in the centre) and the f-block (the two series at the bottom).
Electrons are distributed among shells and subshells according to the aufbau principle (increasing order of energy), the Pauli exclusion principle and Hund’s rule of maximum multiplicity. Because periodicity in electronic configuration follows from increasing \(Z\), the position of an element in the table is fixed by the quantum numbers of its last filled orbital.
The period number equals the principal quantum number \(n\) of the valence shell. A period begins by filling one electron into the s-subshell of a new shell and ends with an element having a complete octet (a duplet in the first period) of that shell; the inner d and f subshells are filled in between, following the aufbau principle. The number of elements in a period equals the total electron capacity of the subshells filled along it — 2, 8, 8, 18, 18 and 32 for periods 1 to 6.
Down a group, a new shell is added, so the general outer electronic configuration stays the same for groups 1, 2 and 3 and for groups 13 to 18 (Table 7.1). In groups 4 to 12 the d and f subshells enter only at a later stage — d from the 4th period and f from the 6th period — so variation in the outer configuration appears only lower down these groups.
The four blocks are named after the subshell receiving the last electron. Their general outer electronic configurations are: s-block \(n\mathrm{s}^{1\text{-}2}\); p-block \(n\mathrm{s}^2\,n\mathrm{p}^{1\text{-}6}\) (from the second to the seventh period); d-block \(n\mathrm{s}^{0\text{-}2}\,(n-1)\mathrm{d}^{1\text{-}10}\); f-block \(n\mathrm{s}^2\,(n-1)\mathrm{d}^{0\text{-}1}\,(n-2)\mathrm{f}^{1\text{-}14}\).
Helium, with the configuration \(1\mathrm{s}^2\), has no p-subshell but is placed in group 18 of the p-block, because its valence shell is completely filled as a complete duplet, like the complete octets of the other noble gases.
Extra stability of half-filled and fully filled subshells produces variations in configuration — for example Cr (\(Z = 24\)) is \(4\mathrm{s}^1 3\mathrm{d}^5\) instead of \(4\mathrm{s}^2 3\mathrm{d}^4\), and \(_{63}\mathrm{Eu}\) is \(6\mathrm{s}^2 4\mathrm{f}^7 5\mathrm{d}^0\).
s-block elements — group 1 (alkali metals) and group 2 (alkaline earth metals) — are reactive metals found only in the combined state. They have low ionization enthalpies, which decrease down the group, and readily form \(\mathrm{M^{+}}\) or \(\mathrm{M^{2+}}\) ions; their compounds are predominantly ionic, except those of Li and Be.
p-block elements (groups 13 to 18), together with the s-block, are the main group or representative elements. Group 18 is the family of noble/inert gases with closed valence shells and very low reactivity; groups 16 (chalcogens) and 17 (halogens) include reactive nonmetals with highly negative electron gain enthalpies that readily form \(\mathrm{X^{-}}\) or \(\mathrm{X^{2-}}\) anions. The p-block alone contains metals, nonmetals and metalloids, the last lying along the zig-zag line. Nonmetallic character increases across a period and decreases down a group.
d-block elements (groups 3 to 12) are all metals, called transition elements, forming a bridge between the reactive s-block and the less reactive elements of groups 13 and 14. Their partially filled inner d-orbitals give them variable oxidation states, paramagnetism, coloured ions and catalytic ability. Zn, Cd and Hg, with \(n\mathrm{s}^2\,(n-1)\mathrm{d}^{10}\), do not show these typical transition-metal properties.
f-block elements are all metals, arranged in the lanthanoid series (\(_{58}\mathrm{Ce}\) to \(_{71}\mathrm{Lu}\)) and the actinoid series (\(_{90}\mathrm{Th}\) to \(_{103}\mathrm{Lr}\)), named after lanthanum (\(_{57}\mathrm{La}\)) and actinium (\(_{89}\mathrm{Ac}\)). The lanthanides are also called rare earth elements. Since the last electron enters the \((n-2)\)f subshell, they are called inner-transition elements and have very similar properties within a series. Elements beyond \(_{92}\mathrm{U}\) are the transuranium elements, all manmade and radioactive.
Periodic trends arise from two simultaneous factors: attraction of the extranuclear electrons by the nucleus, and repulsion among the electrons. Together these give the screening (shielding) effect of the core electrons and the effective nuclear charge:
\[Z_{\text{eff}} = Z - \sigma\]
where \(\sigma\) is the shielding or screening constant, whose value depends on the type of orbital the electron occupies. \(Z_{\text{eff}}\) increases across a period, because the core and hence the shielding stay unchanged while \(Z\) rises; it decreases largely down a group, because a new core shell is added and the extra shielding outweighs the rise in \(Z\).
Atomic radius is one half of the internuclear distance between two adjacent atoms of a metal (metallic radius) or two single-bonded atoms of a nonmetal (covalent radius). It decreases across a period (upto group 17) as \(Z_{\text{eff}}\) increases, and increases down a group as a new shell is added and shielding grows.
Ionic radii follow the same trends. A cation is smaller than its parent atom, because it has fewer electrons, less shielding and a larger \(Z_{\text{eff}}\); an anion is larger, because the extra electrons increase repulsion and lower \(Z_{\text{eff}}\). Among isoelectronic species, which have the same number of electrons, the one with the larger nuclear charge has the smaller radius.
Ionization enthalpy (\(\Delta_i H\)) is the energy required to remove an electron from an isolated gaseous atom in its ground state, expressed in \(\text{kJ mol}^{-1}\):
\[\mathrm{X}\text{(g)} \rightarrow \mathrm{X^{+}}\text{(g)} + \mathrm{e^{-}} \quad ; \ \Delta_i H_1 \qquad\qquad \mathrm{X^{+}}\text{(g)} \rightarrow \mathrm{X^{2+}}\text{(g)} + \mathrm{e^{-}} \quad ; \ \Delta_i H_2\]
It is always positive, and \(\Delta_i H_2 > \Delta_i H_1\) because the electron is removed from a positively charged species. It decreases down a group and increases across a period, being lowest for the alkali metal and highest for the inert gas of a period. Irregularities occur where the electron is removed from a less penetrating orbital (\(\Delta_i H_1\) of B is less than that of Be) or from a doubly occupied p-orbital (\(\Delta_i H_1\) of O is less than that of N).
- Electron gain enthalpy (\(\Delta_{eg}H\)) is the enthalpy change when an electron is added to an isolated gaseous atom in its ground state, in \(\text{kJ mol}^{-1}\):
\[\mathrm{X}\text{(g)} + \mathrm{e^{-}} \rightarrow \mathrm{X^{-}}\text{(g)} \quad ; \ \Delta_{eg}H\]
It may be negative (exothermic) or positive (endothermic). Group 17 elements have very high negative values, group 18 elements high positive values, and alkali metals very low negative values. Its trends are less regular than those of ionization enthalpy, but overall it becomes more negative across a period upto the second last element, and less negative down a group.
Electronegativity (\(EN\)) is the ability of a covalently bonded atom to attract the shared electrons towards itself. It cannot be measured experimentally; the Pauling scale is the most widely used, with Linus Pauling (1922) arbitrarily assigning 4.0 to fluorine. \(EN\) increases across a period as \(Z_{\text{eff}}\) increases, and decreases down a group as the valence shell grows and shielding increases. It predicts the nature of the bond between two atoms.
Among chemical properties, valency expresses combining power and is usually equal to the number of valence electrons or to \(8 -\) (number of valence electrons). Valency remains the same down a group and varies gradually across a period. Oxidation number is the related quantity that carries a \(+\) or \(-\) sign, decided by the electronegativities of the bonded atoms.
Metallic character increases down a group and decreases across a period, following the fall in ionization enthalpy; nonmetallic character increases across a period and decreases down a group, following the increasingly negative electron gain enthalpy.
Chemical reactivity depends on how easily an element attains the configuration of the nearest inert gas. Elements preceding an inert gas react by gaining electrons; those following one react by losing them. The most reactive elements therefore lie at the two extremes of a period — the extreme left, with the smallest ionization enthalpy, and the extreme right excluding the inert gases, with the most negative electron gain enthalpy.
Reactivity is also shown by the nature of the oxides. Alkali metals give basic oxides (\(\mathrm{Na_2O}\)), which with water give strong bases (\(\mathrm{NaOH}\)); halogens give acidic oxides (\(\mathrm{Cl_2O_7}\)), which with water give strong acids (\(\mathrm{HClO_4}\)); oxides of elements in the centre are amphoteric (\(\mathrm{Al_2O_3}\)), neutral (\(\mathrm{CO}\), \(\mathrm{NO}\)) or weakly acidic (\(\mathrm{CO_2}\)).
Because the change in atomic radii along the transition and inner-transition series is small, the members of each series have similar chemical properties, and their ionization enthalpies are intermediate between those of the s-block and the p-block.
The whole chapter can be read as a single chain of cause and effect, summarised in FC 7.8.
Flowchart
FC 7.8 – The chapter in one chain: elements ordered by atomic number Z → periodicity in electronic configuration → position in the modern periodic table (period = n of valence shell; block = subshell of last electron; group = count of outer electrons) → blockwise characteristics (s, p, d, f) → effective nuclear charge Z(eff) = Z − σ → atomic and ionic radius → ionization enthalpy, electron gain enthalpy, electronegativity → valency, metallic/nonmetallic character → chemical reactivity and nature of the oxides
7.7 Check Your Concepts
Mendeleev’s periodic law and the modern periodic law differ in only one word, yet that word changed the whole basis of classification. State both laws, identify the word, and explain why the change was necessary.
Mendeleev left blank boxes in his table rather than shifting elements to fill them. Explain how this decision strengthened, rather than weakened, the case for his periodic law.
Assertion (A): Atomic number is regarded as a more fundamental property of an atom than atomic mass. Reason (R): Moseley showed that the frequency of X-rays emitted by an element is related to its atomic number. State whether (A) and (R) are true, and whether (R) correctly explains (A).
The inert gases were not predicted by Mendeleev at all, yet his table could accommodate them once they were discovered. Explain how this was possible, and what it shows about the structure of a good classification.
Why is the period number of an element the same as the principal quantum number of its valence shell? Would this still hold if the 3d subshell were lower in energy than 4s? Justify your answer.
The third shell has a total electron capacity of 18, yet the third period contains only 8 elements. Explain this apparent contradiction.
Complete the following statement and justify it: “A period begins with the filling of the ____ subshell of a new shell and ends when the ____ subshell of the same shell is complete.” Where, in this scheme, do the d and f subshells get filled?
Explain why the general outer electronic configuration stays the same down groups 1, 2 and 3 and down groups 13 to 18, but varies at a later stage down groups 4 to 12.
Helium is placed in group 18 although it has no p-electrons, while beryllium with the configuration \(2\mathrm{s}^2\) is placed in group 2. Justify both placements.
The expected outer configuration of Cr (\(Z = 24\)) is \(4\mathrm{s}^2 3\mathrm{d}^4\), but the observed one is \(4\mathrm{s}^1 3\mathrm{d}^5\). Explain the observation, and name one f-block element that shows a comparable deviation.
Match the following:
| Column A | Column B |
|---|---|
| a. \(n\mathrm{s}^{1\text{-}2}\) | i. Inner-transition elements |
| b. \(n\mathrm{s}^2\,n\mathrm{p}^{1\text{-}6}\) | ii. s-block |
| c. \(n\mathrm{s}^{0\text{-}2}\,(n-1)\mathrm{d}^{1\text{-}10}\) | iii. p-block |
| d. \(n\mathrm{s}^2\,(n-1)\mathrm{d}^{0\text{-}1}\,(n-2)\mathrm{f}^{1\text{-}14}\) | iv. Transition elements |
Zn, Cd and Hg lie within the d-block but are not regarded as typical transition metals. State the configuration they share, and explain which three characteristic transition-metal properties they consequently fail to show.
Why are the lanthanoids and actinoids called inner-transition elements, and why do the members of each series resemble one another so closely?
The name “rare earth element” is often described as a misnomer. Explain why, referring to what actually made these elements difficult to obtain.
Assertion (A): The effective nuclear charge experienced by a valence electron increases across a period. Reason (R): The number of core electrons remains unchanged across a period while the nuclear charge increases by one at each step. State whether (A) and (R) are true, and whether (R) correctly explains (A).
Down a group, both the nuclear charge and the number of core electrons increase. Explain why the net effect is a decrease in the effective nuclear charge felt by the valence electrons.
“An atom has no definite boundary, yet chemists quote atomic radii to the nearest picometre.” Resolve this apparent conflict, and state how the atomic radius of a metal and of a nonmetal is each estimated.
Arrange the following in the increasing order of size and justify the order: \(\mathrm{Mg}\), \(\mathrm{Mg^{2+}}\), \(\mathrm{O}\), \(\mathrm{O^{2-}}\).
The first ionization enthalpy of boron is lower than that of beryllium, and that of oxygen is lower than that of nitrogen. Explain each of these irregularities separately, naming the orbital involved in every case.
Explain why the second ionization enthalpy of sodium is very much larger than its first, whereas for magnesium the two values are much closer together.
Electron gain enthalpy can be negative or positive, but ionization enthalpy is always positive. Explain the difference, and state what a large positive electron gain enthalpy tells you about an atom.
Fluorine is the most electronegative element, yet chlorine has the more negative electron gain enthalpy. Distinguish clearly between the two properties, and account for the observation.
Electronegativity cannot be measured experimentally, yet the Pauling scale assigns it definite numerical values. Explain how such a scale can be constructed and why the value 4.0 was fixed for fluorine.
Arrange the elements Na, Mg, Si and Cl in the increasing order of (a) atomic radius, (b) first ionization enthalpy and (c) electronegativity. State the single underlying factor responsible for all three orders.
Elements of group 15, 16 and 17 each show two valencies, while those of groups 1, 2 and 13 show only one. Explain, using the two rules for valency of main group elements.
Predict whether the oxide of each of the following will be basic, amphoteric or acidic, and give a reason in each case: \(\mathrm{K}\), \(\mathrm{Al}\), \(\mathrm{S}\).
Write the equations to show that \(\mathrm{Al_2O_3}\) is amphoteric. Explain what the position of aluminium in the periodic table has to do with this behaviour.
Why do the most reactive elements of a period lie at its two extremes, excluding the inert gases? State how the reactivity at the two ends differs in nature.
The elements of a transition series have very similar chemical properties, while those of a short period differ sharply from one another. Account for this contrast in terms of the subshell being filled.
An element X forms a colourless, predominantly ionic chloride and reacts vigorously with cold water; an element Y forms a coloured chloride and reacts only slowly with water. Identify the block to which each belongs and justify your assignment.
Explain the following
- The elements Li, B, C, Be and N have the electronegativities 1.0, 2.0, 1.5 and 3.0, respectively, on the Pauling scale.
- The atomic radii of Cl, I and Br are 99, 133 and 114 pm, respectively.
- The ionic radii of \(\mathrm{F^{-}}\) and \(\mathrm{Na^{+}}\) are 133 and 98 pm, respectively.
- \(_{13}\mathrm{Al}\) is a metal, \(_{14}\mathrm{Si}\) is a metalloid and \(_{15}\mathrm{P}\) is a nonmetal.
- Cu forms coloured salts while Zn forms colourless salts.
Write the outer electronic configuration of the following using orbital notation method. Justify.
- Ge (belongs to period 4 and group 14)
- Po (belongs to period 6 and group 16)
- Cu (belongs to period 4 and group 11)
La belongs to group 3 while Hg belongs to group 12 and both belong to period 6 of the periodic table. Write down the general outer electronic configuration of the ten elements from La to Hg together using orbital notation method.
Ionization enthalpy of Li is 520 kJ mol\(^{-1}\) while that of F is 1681 kJ mol\(^{-1}\). Explain.
Explain the screening effect with a suitable example.
Why is the second ionization enthalpy greater than the first ionization enthalpy?
Why do the elements belonging to the same group have similar chemical properties?
Explain: electronegativity and electron gain enthalpy. Which of the two can be measured experimentally?
The electronic configurations of some elements are given below: a. \(1\mathrm{s}^2\) b. \(1\mathrm{s}^2\,2\mathrm{s}^2\,2\mathrm{p}^6\). In which group and period of the periodic table are they placed?
For each of the following pairs, indicate which of the two species is of large size:
- \(\mathrm{Fe^{2+}}\) or \(\mathrm{Fe^{3+}}\)
- \(\mathrm{Mg^{2+}}\) or \(\mathrm{Ca^{2+}}\)
Select the smaller ion from each of the following pairs:
- \(\mathrm{K^{+}}\), \(\mathrm{Li^{+}}\)
- \(\mathrm{N^{3-}}\), \(\mathrm{F^{-}}\)
Define:
- Ionic radius
- Electronegativity
Compare the chemical properties of metals and nonmetals.
What are the valence electrons? For s-block and p-block elements show that the number of valence electrons is equal to its group number.
Define ionization enthalpy. Name the factors on which ionisation enthalpy depends. How does it vary down the group and across a period?
How does the atomic size vary in a group and across a period? Explain with a suitable example.
Give reasons.
- Alkali metals have low ionization enthalpies.
- Inert gases have exceptionally high ionization enthalpies.
- Fluorine has less electron affinity than chlorine.
- Noble gases possess relatively large atomic size.
Consider the oxides \(\mathrm{Li_2O}\), \(\mathrm{CO_2}\), \(\mathrm{B_2O_3}\).
- Which oxide would you expect to be the most basic?
- Which oxide would be the most acidic?
- Give the formula of an amphoteric oxide.
7.8 Practice With Numericals
Assuming that only the core electrons shield the valence electrons, and that each core electron shields completely, calculate the effective nuclear charge experienced by the valence electron of \(_{11}\mathrm{Na}\), \(_{12}\mathrm{Mg}\) and \(_{17}\mathrm{Cl}\). Comment on the order obtained.
Using the same approximation, calculate \(Z_{\text{eff}}\) for the valence electrons of \(_{4}\mathrm{Be}\) and \(_{20}\mathrm{Ca}\). Both belong to group 2; state what the calculation implies about their relative atomic sizes and what additional factor must be considered.
The I–I bond length in an iodine molecule is 266 pm. Calculate the atomic radius of iodine and compare your answer with the value listed in Table 7.2.
In a sample of metallic sodium, the distance between the centres of two adjacent atoms is 372 pm. Calculate the metallic radius of sodium.
Using the atomic radii in Table 7.2, calculate the percentage decrease in atomic radius on going from Na to Cl across period 3.
From Table 7.4, calculate (a) the difference between the first ionization enthalpies of Ne and Li, and (b) the ratio of the first ionization enthalpy of F to that of Li. Express the ratio to two decimal places.
From Table 7.3, calculate the percentage decrease in the first ionization enthalpy on going from Li to Cs down group 1.
The sixth period is filled in the order \(6\mathrm{s} < 4\mathrm{f} < 5\mathrm{d} < 6\mathrm{p}\). Write the electron capacity of each of these four subshells and hence calculate the number of elements in the sixth period. Repeat the calculation for the fifth period, which fills 5s, 4d and 5p.
An element lies in period 5 and group 15 of the modern periodic table. Write its complete electronic configuration in the aufbau order, and hence determine its atomic number.
Determine the atomic number of the last element of period 4 and of the last element of period 5, showing the arithmetic that leads to each. State what both these elements have in common in their valence shell.
Using the electronegativity values in Table 7.5, calculate the electronegativity difference for the bonds Na–Cl, Al–Cl, H–Cl and Cl–Cl, and arrange the four bonds in the decreasing order of polarity.
State the number of electrons present in each of the species \(\mathrm{O^{2-}}\), \(\mathrm{F^{-}}\), \(\mathrm{Ne}\), \(\mathrm{Na^{+}}\) and \(\mathrm{Al^{3+}}\), and the nuclear charge of each. Hence arrange them in the decreasing order of radius.
An element has the electronic configuration \([\mathrm{Ar}]\,4\mathrm{s}^2 3\mathrm{d}^{10} 4\mathrm{p}^3\). Calculate its atomic number, and state its period, group and block.
The valencies of an element A of group 13 and an element B of group 16 are to be used to predict the empirical formula of the compound they form. Determine the valency of each from its group number, and hence write the empirical formula by the method of cross multiplication.
7.9 Case-Based Questions
- With the help of the diagram, answer the questions given below:
- Which atom should have the smaller ionization enthalpy, oxygen or sulfur?
- The lithium forms +1 ions while beryllium forms +2 ions?
Diagram
Fig. 7.6 – Blank grid of the periodic table as printed in the NCERT exercise, with only four boxes filled in: Li and Be at the left of period 2, and O in period 2 and S in period 3 towards the right of the grid
7.10 Objective Questions
- Consider the elements B, Al, Mg and K; predict the correct order of metallic character:
- B > Al > Mg > K
- Al > Mg > B > K
- Mg > Al > K > B
- K > Mg > Al > B
- In the modern periodic table, the period number indicates the:
- atomic number
- atomic mass
- principal quantum number
- azimuthal quantum number
- The lanthanides are placed in the periodic table at
- left hand side
- right hand side
- middle
- bottom
- If the valence shell electronic configuration is \(n\mathrm{s}^2 n\mathrm{p}^5\), the element will belong to
- alkali metals
- halogens
- alkaline earth metals
- actinides
- In which group of elements of the modern periodic table are halogens placed?
- 17
- 6
- 4
- 2
- Which of the atomic numbers represent the s-block elements?
- 7, 15
- 3, 12
- 6, 14
- 9, 17
- Which of the following pairs is NOT isoelectronic?
- \(\mathrm{Na^{+}}\) and Na
- \(\mathrm{Mg^{2+}}\) and Ne
- \(\mathrm{Al^{3+}}\) and \(\mathrm{B^{3+}}\)
- \(\mathrm{P^{3-}}\) and \(\mathrm{N^{3-}}\)
- Which of the following pairs of elements has similar properties?
- 13, 31
- 11, 20
- 12, 10
- 21, 33