Classification of Elements and Periodicity in Properties — Study Notes

1. Introduction

The periodic table is one of the most powerful organising tools in chemistry. It arranges the 118 known elements so that elements with similar properties fall in vertical columns, revealing patterns in atomic structure and reactivity. The modern periodic table was developed from the work of Dobereiner, Newlands, Mendeleev and Moseley, and it allows chemists to predict the properties of elements and their compounds with remarkable accuracy.

Early attempts at classification were based on atomic mass. Dobereiner grouped elements into triads, Newlands proposed the law of octaves, and Mendeleev in 1869 arranged elements in order of increasing atomic mass while leaving gaps for undiscovered elements. However, Mendeleev's table had anomalies, and it was Henry Moseley who showed in 1913 that the fundamental property determining an element's position is its atomic number.

The modern periodic law states that the physical and chemical properties of elements are a periodic function of their atomic number. This chapter explores how the periodic table is structured, how electronic configurations repeat periodically, and how properties such as atomic radius, ionisation enthalpy, electron gain enthalpy and electronegativity vary down groups and across periods.

2. Development of the Periodic Table

Johann Dobereiner noted that some elements could be arranged in triads where the atomic mass of the middle element was the average of the other two, as in lithium, sodium and potassium. John Newlands proposed the law of octaves, observing that every eighth element resembled the first, similar to musical octaves, but this law failed beyond calcium.

Mendeleev's periodic table arranged elements by increasing atomic mass and grouped them by similar properties, leaving gaps for elements yet to be discovered. It successfully predicted the properties of gallium, scandium and germanium. However, it could not explain why tellurium (128) must precede iodine (127), nor could it place isotopes and lanthanides sensibly.

The modern periodic table solved these problems. Moseley discovered that the atomic number, or the number of protons, is the fundamental property of an element. The modern periodic law states that properties of elements are a periodic function of atomic number, and the table is now arranged in 18 vertical columns called groups and 7 horizontal rows called periods.

3. Nomenclature of Elements with Atomic Number Above 100

Elements with atomic numbers above 100 have temporary systematic names derived from Latin and Greek numerals. Each digit is assigned a root: 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn. The element with Z = 101 is unnilunium (Unu), Z = 102 is unnilbium (Unb), Z = 103 is unniltrium (Unt), Z = 104 is unnilquadium (Unq) and Z = 105 is unnilpentium (Unp).

Temporary symbols consist of three letters, the first in capital and the following two in lowercase, and the name ends with the suffix -ium. For example, element 106 is unnilhexium (Unh), 107 is unnilseptium (Uns) and 108 is unniloctium (Uno). Once the element is officially recognised, it receives a permanent name, such as Rutherfordium (Rf) for 104, Dubnium (Db) for 105 and Seaborgium (Sg) for 106.

4. Electronic Configuration and the Periodic Table

The periodic table is intimately linked to electronic configuration. Elements in the same group have the same outer-shell electronic configuration, which is why they show similar chemical behaviour. Elements in the same period have the same number of shells but different outer electrons.

• s-block elements have their outermost electron in an s orbital; groups 1 and 2, including hydrogen and helium.
• p-block elements have their outermost electron in a p orbital; groups 13 to 18.
• d-block elements have partially filled d orbitals; groups 3 to 12, the transition elements.
• f-block elements have partially filled f orbitals; the lanthanoids (Z = 58-71) and actinoids (Z = 90-103), placed at the bottom of the table.

The position of an element reveals its configuration. For example, an element with configuration [Ar] 3d^5 4s^2 belongs to the d-block, group 7 and period 4. The group number for s and p block elements is given by the number of electrons in the outermost shell (for the main groups), while for d-block elements it is related to the total electrons in (n-1)d and ns orbitals.

5. Periodicity of Properties

Periodicity refers to the recurrence of properties at regular intervals when elements are arranged in order of atomic number. Several fundamental properties show clear periodic trends.

Atomic radius is the distance from the centre of the nucleus to the outermost shell of an electron. Within a period, atomic radius decreases from left to right because the effective nuclear charge increases while the shells remain the same, pulling electrons closer. Down a group, the radius increases because additional shells are added. Covalent radius, metallic radius and van der Waals radius describe the size in different bonding situations.

Ionisation enthalpy is the minimum energy required to remove the most loosely bound electron from a gaseous isolated atom. It increases across a period and decreases down a group. Elements with stable configurations, such as half-filled or fully filled subshells, have unusually high ionisation enthalpies. The successive ionisation enthalpies always increase because removing an electron from a positive ion becomes harder.

Electron gain enthalpy is the energy change when an electron is added to a neutral gaseous atom. It becomes more negative across a period because of increasing effective nuclear charge, and less negative down a group as atomic size increases. Group 17 elements have very negative electron gain enthalpies, while noble gases have positive values. Chlorine has a more negative electron gain enthalpy than fluorine despite fluorine being smaller, because of greater electron-electron repulsion in the compact fluorine atom.

Electronegativity is the tendency of an atom to attract the shared pair of electrons in a covalent bond. It increases across a period and decreases down a group. Fluorine is the most electronegative element with a value of 4.0 on the Pauling scale, while caesium and francium have the lowest values.

Valency of an element in a period first increases and then decreases. Across a period, valency increases from 1 (for alkali metals) to a maximum and then decreases to zero for noble gases. Metallic character decreases across a period and increases down a group, while the reverse holds for non-metallic character.

6. Anomalous Properties of Second Period Elements

The elements of the second period (Li, Be, B, C, N, O, F) differ remarkably from their heavier group members. They have small atomic sizes, high ionisation enthalpies, high electronegativity and limited valency (maximum of four). They have no d orbitals, so they cannot expand their valency beyond four, unlike elements of the third period.

Lithium differs from sodium in its diagonal relationship with magnesium, and beryllium resembles aluminium. The diagonal relationship arises because the small size and high charge of the second period elements make them behave like elements diagonally below and to the right. This explains why Li forms a chloride that is covalent and why Be(OH)2 is amphoteric.

The anomalous behaviour means that the first element of each group is often significantly different from its successors, and this must be remembered when comparing properties such as basicity, solubility and the stability of compounds.

7. Trends in Chemical Reactivity

The chemical reactivity of elements is a consequence of all the periodic properties discussed above. Alkali metals are highly reactive because they have low ionisation enthalpies and easily lose their single valence electron. Halogens are highly reactive non-metals because they have very negative electron gain enthalpies and readily gain an electron.

The reactivity of metals depends on their tendency to lose electrons, which increases down a group. The reactivity of non-metals depends on their tendency to gain electrons, which decreases down a group. The oxides of elements show a periodic trend: metallic oxides are basic, non-metallic oxides are acidic, and amphoteric oxides behave both ways. The acidic character of oxides increases across a period, while the basic character increases down a group.

Quick Revision Tables

Table 1: Blocks of the Periodic Table

Table 2: Trends of Periodic Properties

Table 3: Group Numbers and Valencies

Mind Map

Important Diagrams (SVG)

Diagram 1: Trends of Atomic Radius in the Periodic Table

Diagram 2: Blocks of the Periodic Table

Common Mistakes

1. Comparing atomic radii across a period and down a group incorrectly; remember the decrease across a period and increase down a group.
2. Writing the group number for d-block elements incorrectly; it relates to the sum of (n-1)d and ns electrons, not just the outer s electrons.
3. Assuming fluorine has the most negative electron gain enthalpy; chlorine does, because of electron repulsion in the small F atom.
4. Forgetting that noble gases have positive electron gain enthalpy since adding an electron requires energy.
5. Using atomic mass instead of atomic number when locating an element; the modern periodic law is based on atomic number.
6. Confusing ionisation enthalpy and electron gain enthalpy; one measures energy to remove an electron, the other the energy change on adding an electron.
7. Overlooking the anomalous behaviour of second period elements, especially Li and Be, which show diagonal relationships.

Exam Tips

1. Memorise the order of elements in the first three periods, since trends are frequently tested using these elements.
2. Know that ionisation enthalpy of nitrogen is higher than oxygen because of the stable half-filled 2p^3 configuration.
3. Remember that chlorine has the highest (most negative) electron gain enthalpy in the entire periodic table.
4. Quote the Pauling scale value of 4.0 for fluorine and 0.7 for caesium; electronegativity questions are common.
5. For group identification, use the valence shell configuration: s-block groups 1-2, p-block groups 13-18, d-block groups 3-12.
6. Practise predicting whether an element is metallic, non-metallic or metalloid based on its position and electronegativity.
7. Be able to write the systematic nomenclature names for Z = 101 to 108, as this is a favourite direct question.

Conclusion

The modern periodic table is the logical culmination of nearly a century of attempts to organise the elements. Arranged by atomic number, its rows and columns reflect the periodic repetition of electronic configurations, which in turn controls atomic radius, ionisation enthalpy, electron gain enthalpy and electronegativity. These trends allow chemists to predict reactivity, the nature of oxides and the positions of unknown elements. Understanding the relationship between electronic configuration and periodic properties, together with the anomalous behaviour of second period elements, gives students a powerful framework for explaining the chemistry of all the elements covered in the s-block, p-block and transition series later in the syllabus.