Why Elements Are Arranged by Atomic Number, Not Atomic Mass

More than a hundred elements are known, each with its own properties. Arranged in the right order, they fall into families with similar behaviour, so an element's position tells you a great deal about how it will react.

This lesson covers how periodic classification developed into the modern periodic law, Moseley's relation, and the structure of the modern periodic table.

Early attempts grouped elements by atomic mass, Mendeleev's periodic law stated that properties are a periodic function of atomic mass, and the modern periodic law corrected this by stating that properties are a periodic function of atomic number.

Early groupings:

- Dobereiner's triads — groups of three similar elements in which the middle atomic mass is about the average of the other two; lithium (7), sodium (23) and potassium (39) give $\frac{7+39}{2}=23$
- Newlands' law of octaves — arranged by increasing atomic mass, every eighth element resembles the first, but the pattern held only for lighter elements, up to calcium

Mendeleev's periodic law. The properties of elements are a periodic function of their atomic masses.

Strengths of Mendeleev's table:

- Grouped elements with similar properties together
- Left gaps for elements not yet known and predicted their properties; the predicted eka-aluminium and eka-silicon closely match gallium and germanium
- Corrected doubtful atomic masses of some elements

Limitations:

- Some elements had to be placed out of mass order, such as tellurium (127.6) before iodine (126.9), to keep similar elements together
- Isotopes of one element have different masses but could not be given separate places
- The position of hydrogen was uncertain

Modern periodic law. The physical and chemical properties of elements are a periodic function of their atomic numbers. Ordering by atomic number removes the tellurium-iodine problem and gives all isotopes of an element one place.

An everyday example. Listing students in a register by roll number rather than by height gives each student one fixed place, just as atomic number gives each element one fixed position.

The substance. Periodicity comes from repeating electronic configurations — similar outer electron arrangements recur at regular intervals, which is why atomic number, not mass, is the right basis.

The modern periodic table has 7 horizontal periods and 18 vertical groups, and it is divided into s, p, d and f blocks according to the orbital that receives the last electron.

Periods:

- Each period corresponds to filling a new shell, and the period number equals the principal quantum number of the outermost shell
- Periods 1 to 7 hold 2, 8, 8, 18, 18, 32 and 32 elements

Groups. Elements in the same group have the same number of valence electrons and similar properties; groups are numbered 1 to 18.

Blocks:

- s-block — groups 1 and 2; configuration $n{s}^{1-2}$; reactive metals such as sodium and calcium
- p-block — groups 13 to 18; configuration $n{s}^{2}n{p}^{1-6}$; non-metals, metalloids, some metals and noble gases
- d-block — groups 3 to 12; transition elements with configuration $(n-1)d^{1-10}n{s}^{0-2}$
- f-block — lanthanoids and actinoids, placed separately at the bottom

Finding an element's position:

- Period = the highest principal quantum number
- s-block group = number of valence s electrons; p-block group = 10 + number of valence electrons; d-block group = number of $(n-1)d$ plus $ns$ electrons

Worked example 1. Chlorine (Z = 17), $[\mathrm{Ne}]3s^{2}3p^5$: period 3, p-block, group $10+7=17$.

Worked example 2. Iron (Z = 26), $[\mathrm{Ar}]3d^{6}4s^2$: period 4, d-block, group $6+2=8$.

An everyday example. Orange sodium street lamps and red strontium fireworks both rely on s-block metals whose behaviour follows from their place in the table.

The substance. Helium sits in group 18 although it is an s-block element — its $1s^2$ configuration is a complete shell, so it behaves like the other noble gases.