Classification of Elements and Periodicity in Properties

Written and reviewed by Akanksha Tiwari, Physics & Chemistry faculty at UniDrill (M.Sc. Physics, Banaras Hindu University · 8 years teaching Class XI–XII science). Based on NCERT Chemistry Class XI, Chapter 3.

Could you predict an element nobody had seen?

In 1869 Dmitri Mendeleev arranged the known elements in a table, and where a pattern broke he left a gap and said an element would be found there. He even gave its properties. For the gap under silicon he predicted an atomic weight near 72, a density of 5.5 g cm⁻³ and an oxide of formula EO₂. In 1886 germanium was discovered with atomic weight 72.6, density 5.36 and the oxide GeO₂. A table that predicts the future is a powerful tool, and this chapter shows why it works. The answer is the electron configurations you met in Structure of Atom. Elements repeat their properties because they repeat their outer electron arrangement.

Why classify, and the first attempts

By 1800 only 31 elements were known, by 1865 there were 63, and today 118 have been named. Studying each element and its compounds one by one is hopeless, so chemists looked for patterns. Every attempt used atomic weight, because atoms of different elements were first compared by mass.

Check a triad. Lithium (7), sodium (23) and potassium (39): the mean of the outer two is $\dfrac{7 + 39}{2} = 23$, exactly sodium's weight. Calcium, strontium, barium (40, 88, 137) give $\dfrac{40 + 137}{2} = 88.5$, and chlorine, bromine, iodine (35.5, 80, 127) give $81.3$, both close.

Mendeleev's table: gaps that predicted the future

Mendeleev put elements in order of increasing atomic weight but grouped them by chemical similarity, and when the two clashed, similarity won. Iodine (126.9) is lighter than tellurium (127.6), yet he put iodine after tellurium so that it sat in the same column as chlorine and bromine. He also left gaps, naming the missing elements eka-aluminium, eka-silicon and eka-boron (later found as gallium, germanium and scandium). The predictions were remarkably accurate.

The table had real defects, though. In order of atomic weight, potassium should come before argon, nickel before cobalt and iodine before tellurium, but their properties demanded the opposite order. Hydrogen had to be placed in two groups. There was no place for isotopes, which have different weights but the same chemistry, and no way to fit the lanthanoids. The cause was unknown until atomic structure was understood.

Moseley and the Modern Periodic Law

In 1913 Henry Moseley bombarded different metals with electrons and measured the frequency $\nu$ of the characteristic X-rays they gave off. A plot of $\sqrt{\nu}$ against atomic number $Z$ was a straight line, which means the relation is

$$\sqrt{\nu} = a\,(Z - b)$$

for constants $a$ and $b$ specific to the type of X-ray line. Plotted against atomic weight, the same data scatter. So atomic number, the number of protons, is the real ordering quantity. This gives the Modern Periodic Law: the physical and chemical properties of elements are periodic functions of their atomic numbers. Mendeleev's anomalies vanish at once, because Ar ($Z = 18$) does come before K ($Z = 19$).

How the long form of the table is built

The long form has 7 periods (rows) and 18 groups (columns, numbered 1 to 18 under the IUPAC system). The period number is the highest principal quantum number $n$ occupied. The number of elements in a period comes straight from the filling order: count the orbitals filled between one noble gas and the next, and double.

Period 4, for example, has $1 + 5 + 3 = 9$ orbitals, and 9 orbitals hold 18 electrons, which is why the fourth period has 18 elements. The 14 lanthanoids (Ce to Lu) and 14 actinoids (Th to Lr) belong to periods 6 and 7 but are printed in two rows at the bottom to keep the table narrow.

To place an element, write its configuration, and then use these rules.

Worked example: find the position. Locate the elements with $Z$ = 35, 26, 56, 82 and 29.

The common wrong step: giving copper group 1 because it ends in 4s¹. The 3d electrons count towards the group number, so Cu is in group 11.

Beyond the textbook. For the exam: hydrogen ($1s^1$) is placed at the top of group 1, and helium ($1s^2$), although its configuration is s block, is placed in group 18 with the other noble gases because its shell is completely filled. In reality: hydrogen fits no group well. It gains an electron like the halogens and loses one like the alkali metals, yet it is a non metal. Some tables place it above the halogens and others show it on its own. Why it matters: it shows that a periodic table is a convenient arrangement of repeating properties, and a few elements will never fit perfectly.

Naming the elements above Z = 100

New heavy elements are made in laboratories and exist for fractions of a second. While discovery is disputed, IUPAC gives each a temporary systematic name built from its atomic number, using a root for each digit and ending in "ium".

Worked example: systematic names. Name the elements with $Z$ = 120, 119 and 102, and give their three letter symbols.

The common wrong step: dropping the "nil" for zero, which turns 120 into "Unbium". When a root ending in "i" (bi, tri) meets "ium", only one "i" is kept, as in Unnilbium.

After the discoveries were confirmed, permanent names were approved.

The four blocks and the metal to non metal divide

Where an element sits depends on which subshell receives its last electron, and the block tells you how it behaves.

More than 78% of elements are metals, found on the left and centre. They are usually solid with high melting and boiling points, and are malleable, ductile and good conductors. Non metals sit at the top right, are often gases or low melting solids, and are poor conductors. A zig zag line separates them, and along it lie the metalloids (Si, Ge, As, Sb, Te) with intermediate properties. Metallic character falls from left to right across a period and rises down a group.

Beyond the textbook. For the exam: group 12 elements Zn, Cd and Hg are d block elements. In reality: the IUPAC definition of a transition element is one with an incomplete d subshell in the atom or in a common ion. Zn, Cd and Hg have full $d^{10}$ shells in the atom and in their usual +2 ions, so they do not behave as typical transition metals. Why it matters: you will meet this again in the d and f block chapter, where colour and magnetism need partly filled d orbitals.

The master idea: effective nuclear charge

Every periodic trend rests on one question: how strongly does the nucleus hold the outermost electron? The answer is the effective nuclear charge, the pull of the nucleus after the inner electrons have shielded part of it.

$$Z_{\text{eff}} = Z - S$$

where $S$ measures the shielding. Across a period, the electron added goes into the same shell, where electrons shield one another poorly, so $Z_{\text{eff}}$ rises and the atom shrinks. Down a group, a new shell is added, which is far from the nucleus and well shielded by all the inner shells, so the outer electron is held more loosely despite the higher $Z$. With these two facts you can predict almost every trend below.

Atomic radius

An atom has no sharp edge, so its size is defined from distances between nuclei, and the definition depends on how it is bonded.

These four tables hold most of the data you will need for the rest of the chapter. Radii fall steadily across a period and rise down a group. The covalent and metallic radii are listed together here, and the noble gas values (van der Waals) are left out because they are measured differently.

Worked example: radius from bond length. The Br–Br bond length is 228 pm. Estimate the covalent radius of bromine, and then the C–Br bond length (covalent radius of C is 77 pm).

The common wrong step: comparing a noble gas radius with its neighbours without noticing it is a van der Waals radius. Neon looks bigger than fluorine only because it is measured differently.

Ionic radius

An ion's size depends on its charge, and the rule is simple. A cation is smaller than its atom because it has lost electrons (often a whole shell), while the same nuclear charge now pulls fewer electrons. An anion is larger than its atom because the extra electrons repel each other and $Z_{\text{eff}}$ per electron falls. Na is 186 pm but Na⁺ is about 100 pm. Cl is 99 pm but Cl⁻ is 181 pm. The more positive the charge, the smaller the cation, so Fe³⁺ is smaller than Fe²⁺.

Species with the same number of electrons are isoelectronic. Among these, size is controlled by nuclear charge alone, and more protons pull the same electrons closer.

Worked example: rank isoelectronic ions. Arrange O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ in order of decreasing radius.

Pause and predict: Which is larger, Fe²⁺ or Fe³⁺?

Check your answer: Fe²⁺. Both have the same nuclear charge of 26 protons, but Fe³⁺ has one electron fewer, so the same pull acts on fewer electrons that repel each other less, and the ion shrinks. The more positive the charge on a cation of one element, the smaller it is.

Ionisation enthalpy

The first ionisation enthalpy is the energy needed to remove the most loosely held electron from an isolated gaseous atom. It is always positive, and it is the measure of how well an element can form a cation.

X(g) → X⁺(g) + e⁻ X⁺(g) → X²⁺(g) + e⁻ (second)

Removing each further electron is harder than the one before, because the electron is leaving an ion of higher positive charge: IE₃ > IE₂ > IE₁. A sudden huge jump shows that the next electron comes from an inner, tightly held shell.

Worked example: group from successive ionisation enthalpies. An element has successive ionisation enthalpies (kJ mol⁻¹) of 738, 1451, 7733 and 10543. Which group is it in?

Trend. Across a period, ionisation enthalpy rises because $Z_{\text{eff}}$ rises and the radius shrinks. The peaks are at noble gases (full shells) and the lows at alkali metals (one loosely held electron). Down a group it falls, because the outer electron is farther and better shielded. Group 1 shows this clearly, with IE₁ falling from 520 kJ mol⁻¹ for Li to 376 kJ mol⁻¹ for Cs.

Look again at the period 2 and period 3 columns of the data tables. Each period follows the rule except at two places. The full period 2 order is Li < B < Be < C < O < N < F < Ne. Two dips explain the exceptions. Boron's electron comes from a 2p orbital, which is farther out and shielded by the 2s electrons, so it leaves more easily than the 2s electron of beryllium. In oxygen, the fourth 2p electron must pair with another in the same orbital, and the repulsion helps it escape. Nitrogen's $2p^3$ is half filled, with three unpaired electrons, which is a stable arrangement.

Pause and predict: Why is the second ionisation enthalpy of sodium (4562 kJ mol⁻¹) so much higher than that of magnesium (1451 kJ mol⁻¹)?

Check your answer: Removing the second electron from Na⁺ means breaking into the stable $2p^6$ core. In Mg⁺ the second electron is still the outer 3s electron. A noble gas core is far more tightly held.

Electron gain enthalpy

Electron gain enthalpy is the enthalpy change when a gaseous atom takes up an electron.

X(g) + e⁻ → X⁻(g)

It is negative when energy is released. The more negative it is, the more readily the atom accepts an electron. Halogens are the most negative, because one electron completes their octet.

In the group 17 table above, chlorine, not fluorine, is the most negative. Fluorine is so small that its 2p subshell is crowded, and the incoming electron is repelled strongly by the electrons already there. In the larger 3p subshell of chlorine the repulsion is less. The same effect makes O (−141 kJ mol⁻¹) less negative than S (−200 kJ mol⁻¹). Noble gases have positive values (He +48, Ne +116, Ar +96 kJ mol⁻¹), because their shells are full and the new electron must go into the next shell. Nitrogen, with its stable half filled $2p^3$, has a value close to zero, much less negative than carbon (−122 kJ mol⁻¹).

The common wrong step: treating electron gain enthalpy and ionisation enthalpy as opposites that must follow the same order. They share trends but the exceptions are different.

Electronegativity

Electronegativity is the tendency of an atom bonded to another to attract the shared electrons towards itself. Unlike ionisation enthalpy, it is a relative scale and cannot be measured directly. The most widely used is the Pauling scale, in which fluorine, the most electronegative element, is fixed at 4.0.

The data tables show it rising steadily across each period (Li 1.0 to F 4.0, Na 0.9 to Cl 3.0) and falling down a group (F 4.0 to At 2.2, Li 1.0 to Cs 0.7). Electronegativity is directly related to non metallic character and inversely to metallic character. The three similar looking terms are different quantities.

Valence and oxidation state

Valence is the combining capacity of an element. For groups 1, 2, 13 and 14 it equals the number of valence electrons, and for groups 15, 16 and 17 it is $8 - \text{valence electrons}$, so it is 3, 2 and 1. Transition elements show variable valence. The oxidation state is the charge an atom would have if the shared electrons went to the more electronegative atom. In OF₂, fluorine is more electronegative, so F is −1 and O is +2. In Na₂O, Na is +1 and O is −2.

Second period anomalies and diagonal relationships

The first member of each group (Li, Be, B, C, N, O, F) differs from the rest of its group. The reasons are small size, high charge to radius ratio, high electronegativity and only four valence orbitals (2s and three 2p), with no d orbitals. This puts a ceiling on covalency.

That is why phosphorus makes PCl₃ and PCl₅ while nitrogen makes only NCl₃. The first members also form strong multiple bonds ($p\pi$–$p\pi$) such as C=C, N≡N and C=O.

Moving one step down and one step right, elements in a diagonal pair behave alike, because the increase in charge is offset by the increase in size. The two key pairs are lithium and magnesium, and beryllium and aluminium.

Reactivity and the nature of oxides

Reactivity is highest at the two ends of a period: at the left, metals lose electrons easily (Na reacts with water, 2Na + 2H₂O → 2NaOH + H₂), and at the right, non metals gain them easily. Reactivity is lowest at the centre. The nature of the oxide changes steadily across period 3, from strongly basic to strongly acidic.

Basic oxide with water: Na₂O + H₂O → 2NaOH. Acidic oxide with water: Cl₂O₇ + H₂O → 2HClO₄. An amphoteric oxide reacts with both acid and base:

Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O and Al₂O₃ + 2NaOH + 3H₂O → 2Na[Al(OH)₄]

Neutral oxides, such as CO, NO and N₂O, react with neither. Metallic character grows down a group, so oxides become more basic as you descend.

Exam traps

Trend and reaction sheet

60-second revision

Where this chapter connects

• Before this: Structure of Atom gave the electron configurations and filling rules that decide every position in the table.
• Next: Chemical Bonding and Molecular Structure uses ionisation enthalpy, electron gain enthalpy and electronegativity to decide whether atoms form ionic or covalent bonds.
• Class XII link: The d and f Block Elements explains variable oxidation states, colour and the lanthanoid trends that begin here.