Periodic Classification of Elements — Notes

This chapter is assessed through internal (school-based) assessment, not the board theory paper.

As more elements were discovered, chemists looked for a way to group similar elements so that their properties could be studied and predicted together.

Döbereiner’s triads (early 1800s): groups of three elements with similar properties. When the three are written in order of increasing atomic mass, the atomic mass of the middle element is roughly the average of the other two.

• Example: lithium (6.9), sodium (23.0), potassium (39.1): mean of 6.9 and 39.1 = 23.0.
• Achievement: the first hint that properties are linked to atomic mass.
• Limitation: only a few such triads could be found; most elements did not fit.

Newlands (1866) arranged the known elements in order of increasing atomic mass and found that every eighth element had properties like the first, like the eighth note of a musical scale (sa, re, ga, ma, pa, dha, ni, sa).

• Lithium, sodium and potassium fall in one column; fluorine and chlorine in another.

Limitations

• Worked well only up to calcium; for heavier elements the pattern broke down.
• Assumed only the elements known then existed, so later discoveries did not fit (for example, the noble gases).
• To keep the pattern, two elements were put in one place (cobalt and nickel), even in the same column as very different elements (fluorine, chlorine, bromine), while iron was placed away from cobalt and nickel.

Mendeléev (1869) arranged the 63 elements then known by increasing atomic mass, and placed elements with similar properties in the same vertical column (group). His main test of similarity was the formulas of oxides and hydrides.

Mendeléev’s periodic law: the properties of elements are a periodic function of their atomic masses.

Achievements

• Left gaps for undiscovered elements and predicted their properties: eka-boron (scandium), eka-aluminium (gallium), eka-silicon (germanium). ("Eka" is Sanskrit for "one".)
• Placed elements by properties even when this broke the mass order (tellurium before iodine), so families stayed together.
• The noble gases, found later, fitted in as a new group without disturbing the table.

• Hydrogen: resembles both alkali metals (one electron, forms positive ions) and halogens (forms molecules of two atoms), so it had no fixed place.
• Isotopes: isotopes of an element have different masses but the same chemical properties. A table based on mass would have to place them apart.
• Irregular steps: atomic mass does not increase in regular steps, so it was impossible to say how many elements might be missing between two known ones.
• Mass order broken: some heavier elements had to come before lighter ones to keep families together.

In 1913, Henry Moseley showed that the atomic number (number of protons) is a more fundamental property than atomic mass.

Modern periodic law: the properties of elements are a periodic function of their atomic numbers.

• 18 groups (vertical columns, numbered 1 to 18) and 7 periods (horizontal rows).
• Isotopes have the same atomic number, so they share one place.
• Atomic-number order automatically puts tellurium (52) before iodine (53), cobalt (27) before nickel (28) and argon (18) before potassium (19).
• Metals are on the left and in the middle; non-metals on the upper right; metalloids (boron, silicon, germanium, arsenic, antimony, tellurium, polonium) on a zig-zag line between them.
• Groups 3 to 12 are the transition elements; the lanthanoids and actinoids are shown in two rows below.

The position of an element follows from its electronic configuration.

• Same group → same number of valence electrons → similar properties.
• Period number = number of occupied shells.
• Among the first 20 elements: 1 or 2 valence electrons → group 1 or 2; 3 to 8 valence electrons → group = 10 + valence electrons (helium, with 2, is in group 18).
• Number of elements per period: 2, 8, 8, 18, 18, 32, 32. A shell can hold at most 2n² electrons.

Example: atomic number 13 → 2, 8, 3 → period 3, group 13 (aluminium).

Valency: across a period it rises from 1 to 4 and falls to 0; down a group it stays the same. Valency = valence electrons (1 to 4) or 8 − valence electrons (5 to 8).

Atomic size (atomic radius):

• Across a period: decreases. Electrons go into the same shell while the nuclear charge increases, so they are pulled closer.
• Down a group: increases. A new shell is added each time, so the outer electrons are farther from the nucleus.

Metallic and non-metallic character:

• Metals tend to lose electrons; non-metals tend to gain them.
• Across a period: metallic character decreases, non-metallic character increases.
• Down a group: metallic character increases, non-metallic character decreases.
• Oxides: metal oxides are basic; non-metal oxides are acidic.