Use adaptive quiz-based learning to study this topic faster and more effectively.

# Periodic trends

## Groups and ionic charges

The charge of an ion in an ionic bond can be determined from the group number of the element.

Atoms in groups 1, 2 and 3 tend to form cations with charges of $+1$, $+2$ and $+3$ respectively. These atoms lose their valence electrons to gain full outer shells.

Atoms in groups 5, 6 and 7 form anions with charges of $-3$, $-2$ and $-1$ respectively. These atoms gain electrons.

Atoms in group 4 tend not to form ions directly. They can form ions when bonded to other atoms (for example, the carbonate ion contains carbon).

In the ionic compound, sodium chloride ($\ce{NaCl}$), sodium loses one electron while chlorine gains this electron.

The donated electron (indicated in green) found in chlorine comes from the valence shell of sodium (indicated by the dotted circle).

The atomic radius (plural: atomic radii) is the distance between the centre of an atom and its valence electrons (in the outermost shell).

Atoms with more electron shells have larger atomic radii than atoms with fewer shells.

An atom with three shells has valence electrons that are much further away from the nucleus than the valence electrons in an atom with just one shell.

For groups 1, 2, 3 to 0, the atomic radius increases as one moves down a group in the periodic table because elements further down in a group have more electron shells.

Beryllium ($\ce{Be}$), magnesium ($\ce{Mg}$) and calcium ($\ce{Ca}$) are elements of group 2. As we move from beryllium to calcium, the atomic radius increases.

Nitrogen ($\ce{N}$), oxygen ($\ce{O}$) and fluorine ($\ce{F}$) each have two electron shells. The number of valence electrons (on the second electron shell) increases from nitrogen to fluorine.