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Inorganic chemistry (3.2)›Periodicity (3.2.1)

Periodicity (3.2.1)

Classification of elements in the periodic table, trends in ionisation energy, physical properties, and atomic radius.
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The periodic table is a fundamental tool in chemistry that organises the elements based on their atomic number, creating a systematic arrangement that reveals patterns in their properties and reactivity.

These repeating patterns are known as periodicity.

A detailed periodic table of elements displaying atomic numbers and symbols for each element, including groups and periods, with a key indicating the atomic number and symbol.
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The periodic table is divided into rows called periods.

A detailed periodic table of elements displaying atomic numbers and symbols for each element, organized into periods and groups, with a key indicating the atomic number and symbol.

There are a total of seven periods.

Elements in the same period have the same number of electron shells.

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Columns in the periodic table are known as groups.

A detailed periodic table of elements, displaying the chemical symbols, atomic numbers, and groups of various elements arranged in a grid format. The table includes sections for lanthanides and actinides at the bottom.

There are 18 groups in the periodic table.

Elements within the same group have similar chemical properties because they have the same number of outer electrons.

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All of the elements in the periodic table belong to blocks based on their electron configuration.

s-block: the outermost electron is in the s-subshell; for example, the electron configuration for potassium is

p-block: the outermost electron is in the p-subshell; for example, the electron configuration for aluminium is

d-block: the outermost electron is in the d-subshell; for example, the electron configuration for vanadium is

f-block: the outermost electron is in the f-subshell; for example, the electron configuration for cerium is

A color-coded periodic table of elements, displaying the s-block, p-block, d-block, and f-block. Each block is highlighted in different colors: pink for s-block, blue for d-block, yellow for p-block, and green for f-block. The table includes element symbols, atomic numbers, and their respective positions in the table.
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First ionisation energy decreases down the group.

The increasing atomic radius down the group means the outer electrons are further from the nucleus.

The increase in energy level means there is an increased amount of nuclear shielding from filled electron shells.

The net effect is that, despite the increased nuclear charge, the effective nuclear attraction to the outermost electron decreases. The outermost electron is easier to remove and the first ionisation energy decreases.

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Although across a period, the general trend is that the first ionisation energy increases, there are some exceptions.

A graph showing the first ionization energy (IE) in kJ mol^-1 plotted against the atomic number of elements from lithium (Li) to argon (Ar). The graph features a series of data points connected by a line, illustrating the trend in ionization energy across these elements.

There is a drop in first ionisation energy between Group 2 and 3, and . This is linked to the outermost electron sitting in the p-subshell, which is at a higher energy level than the s-subshell.

There is another drop in first ionisation energy between Group 5 and 6, and . This is linked to electron–electron repulsion in the half filled p-subshell. The fourth electron in the p-subshell is paired in an orbital and so experiences more electron-electron repulsion than the unpaired p-subshell electrons in group five elements.

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The general trend across periods is an increase in melting point from Group 1 to Group 4. There is then a sharp decrease in melting point between Group 4 and Group 5.

The melting points from Group 5 onwards are comparatively low.

The trend exists due to the transition from giant lattices to simple molecules, held together by weak intermolecular forces.

A graph comparing the melting points of elements in Period 2 and Period 3 of the periodic table. The x-axis represents atomic number (Z), while the y-axis shows melting point. Period 2 includes elements from Lithium (Li) to Neon (Ne), with a notable peak at Carbon (C). Period 3 includes elements from Sodium (Na) to Argon (Ar), with a peak at Silicon (Si). Labels indicate regions of 'Giant lattices' and 'Simple molecules' for both periods.

There is an increase in melting point between phosphorus () and sulfur () in period 3. This is due to the increased molecular size and therefore the number of electrons is higher in compared to , creating to more significant temporary dipoles. This leads to stronger intermolecular forces that require more energy to overcome.

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The atomic radius is taken from atoms in their most stable elemental form. It gives information about the distance between the nucleus and the outer shell electrons.

A chart displaying the atomic radius of various elements arranged in a grid format. The atomic radius decreases from left to right and increases from top to bottom. Each element is represented by a circle, with hydrogen (H) at the top left and cesium (Cs) at the bottom left.

Across periods in the periodic table, atomic radii decrease from left to right.

The increasing atomic number across the period, and therefore the increased nuclear charge, attracts the electrons more tightly to the nucleus.

Electrons are added to the same electron shell, so there is no significant increase in nuclear shielding across the period. This means the stronger nuclear attraction directly reduces the atomic radius.

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