Periodic table and periodicity: Class 9 Chemistry

Key points

Chemists needed a way to arrange the known elements so that their properties could be studied and predicted. Dobereiner grouped elements in triads, Newlands arranged them in octaves, and Mendeleev (1869) published a table in which elements were placed in order of increasing atomic mass.

Mendeleev's periodic law said that the properties of elements are a periodic function of their atomic masses. He left gaps for undiscovered elements and predicted the properties of some, such as eka-silicon, which was later found as germanium. His table had faults: a few elements, such as tellurium and iodine, had to be placed out of mass order to fit their properties, and hydrogen had no clear place.

Moseley (1913) studied X-rays from different elements and found that atomic number, not atomic mass, is the basic property. The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. This removed the faults in Mendeleev's order.

The modern periodic table has 7 horizontal rows called periods and 18 vertical columns called groups. Elements in a group have the same number of valence electrons, so they behave in a similar way. The period number equals the number of occupied shells.

The lengths of the periods are:

• Period 1: 2 elements (H, He)
• Periods 2 and 3: 8 elements each (short periods)
• Periods 4 and 5: 18 elements each (long periods)
• Period 6: 32 elements, including the lanthanides
• Period 7: the actinides, placed with period 6 below the main table

The table is divided into blocks according to the subshell that is being filled:

• s-block: groups 1 and 2, reactive metals (helium is also placed in s).
• p-block: groups 13 to 18, which contain metals, metalloids and non-metals.
• d-block: groups 3 to 12, the transition metals.
• f-block: lanthanides and actinides, placed at the bottom.

The position of an element is found from its electronic configuration. The number of shells gives the period. For groups 1 and 2 the group number equals the valence electrons. For groups 13 to 18 the group number is 10 plus the valence electrons. Chlorine (2,8,7) has 3 shells and 7 valence electrons, so it is in period 3, group 17.

Some important groups:

• Group 1, alkali metals (Li, Na, K): soft, very reactive, react with water to give hydrogen and an alkali, stored under oil.
• Group 2, alkaline earth metals (Mg, Ca): harder and less reactive than group 1.
• Group 17, halogens (F, Cl, Br, I): reactive non-metals with 7 valence electrons. Reactivity falls going down the group.
• Group 18, noble gases (He, Ne, Ar): full outer shell, almost unreactive, exist as single atoms.

Periodicity means that properties repeat at regular intervals. Atomic radius is half the distance between the nuclei of two bonded atoms of the same element. Across a period, the nuclear charge rises while electrons enter the same shell, so the nucleus pulls them closer and the radius decreases. Down a group, new shells are added, so the radius increases.

The shielding effect is the reduction in the pull of the nucleus on outer electrons caused by the inner electrons. Down a group, more inner shells give more shielding, so the outer electron is held more loosely.

Ionisation energy is the energy needed to remove one electron from a gaseous atom. It rises across a period because the radius gets smaller, and it falls down a group because of larger size and more shielding. Electronegativity is the power of an atom to attract the shared electrons in a bond. It rises across a period and falls down a group. Fluorine is the most electronegative element.

Metallic character is the tendency to lose electrons. It decreases across a period and increases down a group. This is why sodium is a metal, silicon is a metalloid and chlorine is a non-metal, and why caesium is more reactive than lithium.

For a main-group atom, read its electron arrangement in two ways: occupied shells identify the period, while the outer-shell pattern helps identify the group. Use modern Groups 1 to 18 and do not confuse a group number with the number of protons.

Study tip: Learn the three trends (radius, ionisation energy, electronegativity) as a table showing the direction across a period and down a group, and always attach a reason using nuclear charge, shells and shielding. Practise finding the period and group from atomic number, since this is a regular short question.

Definitions and flashcards

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Short questions with model answers

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Numericals with worked steps

Common mistakes

• Saying that Mendeleev arranged elements by atomic number. He used atomic mass.
• Writing the group of a p-block element as just its number of valence electrons without adding 10.
• Stating that atomic radius increases across a period.
• Confusing a group (vertical) with a period (horizontal).
• Giving hydrogen as an alkali metal. It is placed above group 1 but is a non-metal.

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