IMAT Chemistry · Periodic Table

Groups, Periods, Blocks and Typical Charges

The table orders elements by atomic number; rows (periods) give the outer shell, columns (groups) give the valence electrons, and the group predicts an element's usual charge.

All Periodic Table notes

Why this matters

The 2013 paper asked which of five formulas breaks the usual charges of main-group elements, and 2011 asked about the oxide and sulfate of a group 13 element. Both need the charge an element takes from its group.

Concept 1 of 3: Groups, periods and blocks of the periodic table

The table is built from electron configurations. A new row starts each time electrons begin a new shell, so the period number is the outer shell. Elements in one column have the same number of outer electrons, which is why they react alike. The block is simply the subshell that the last electron enters.

Definition

Elements are arranged in order of atomic number $Z$ . (Writing configurations is taught in the Atomic Structure chapter.)

A period is a row. The period number equals the number of the outer shell: sodium, $[\mathrm{Ne}]\,3s^1$ , is in period 3.
A group is a column, numbered 1 to 18. Older books number the main groups I to VIII, so group 13 is III and group 17 is VII.
Main-group valence electrons: groups 1 and 2 have 1 and 2; groups 13 to 18 have the group number minus 10.
Family names: group 1 alkali metals, group 2 alkaline earth metals, group 17 halogens, group 18 noble gases; groups 3 to 12 are the transition metals.

Block	Groups	Subshell being filled	Examples
s-block	Groups 1 and 2 (and helium)	ns	Na, Mg, Ca
p-block	Groups 13 to 18	np	C, O, Cl, Ne
d-block	Groups 3 to 12	(n−1)d	Fe, Cu, Zn
f-block	Lanthanides and actinides, drawn as two rows below	(n−2)f	Ce, U

Practice this conceptself-check · 3 quick reps

Group 13 to 18 elements have the group number minus 10 valence electrons

Chlorine is in group 17 but has 7 valence electrons, not 17. With the old numbering it is group VII, which gives 7 directly. Check which numbering the question uses.

Concept 2 of 3: Metals, non-metals and metalloids

Metal atoms hold their outer electrons loosely, so they lose them easily, form positive ions and share a sea of mobile electrons that conducts electricity. Non-metal atoms hold electrons tightly and tend to gain or share them. Along the boundary sit the metalloids, which are in between in both looks and behaviour.

Definition

A staircase line runs from boron down to astatine. Metals lie to the left of it (most elements), non-metals to the right, plus hydrogen.
Metalloids lie along the line: boron, silicon, germanium, arsenic, antimony and tellurium.
Metallic character increases down a group and from right to left along a period, so the most metallic elements are at the bottom left.
Metal oxides are basic (they react with acids). Non-metal oxides are acidic or neutral. A few oxides are amphoteric (react with both acids and bases): $\mathrm{Al_2O_3}$ , $\mathrm{ZnO}$ .

Class	Where in the table	Typical physical properties	Oxides	Examples
Metals	Left and centre, below the staircase	Shiny, conduct heat and electricity, malleable; mostly solids (mercury is a liquid)	Basic; a few amphoteric	Na, Mg, Fe, Cu
Non-metals	Right, above the staircase, plus hydrogen	Dull, poor conductors (graphite conducts), brittle as solids; many are gases	Acidic ( $\mathrm{CO_2}$ , $\mathrm{SO_2}$ ) or neutral (CO, $\mathrm{H_2O}$ )	C, N, O, S, Cl
Metalloids	Along the staircase	Look metallic but are brittle; semiconductors	Weakly acidic or amphoteric	B, Si, Ge, As

Practice this conceptself-check · 3 quick reps

Aluminium is a metal, not a metalloid

Aluminium touches the staircase line but is a good conductor and a typical metal. Its oxide is amphoteric, which can mislead. The metalloids are B, Si, Ge, As, Sb and Te.

Concept 3 of 3: Typical ion charges from the group, and writing ionic formulas

Main-group atoms gain or lose just enough electrons to reach the arrangement of the nearest noble gas. Metals on the left lose their few valence electrons; non-metals on the right gain the few they lack. Once you know the two charges, the formula is the smallest whole-number ratio that makes the total charge zero.

Definition

Typical charges (and oxidation states) of main-group elements:

Group 1: +1. Group 2: +2. Group 13: +3 (aluminium, gallium, indium; boron forms covalent compounds).
Group 14: +4 in compounds such as $\mathrm{SnS_2}$ ; tin and lead also show +2.
Group 15: −3 in nitrides and phosphides. Group 16: −2. Group 17: −1 in halides. Group 18: no ions.
Common polyatomic ions: $\mathrm{OH^-}$ , $\mathrm{NO_3^-}$ , $\mathrm{HCO_3^-}$ , $\mathrm{SO_4^{2-}}$ , $\mathrm{CO_3^{2-}}$ , $\mathrm{PO_4^{3-}}$ , $\mathrm{NH_4^+}$ .
In metal hydrides such as $\mathrm{NaH}$ , hydrogen is −1.

Charge balance in an ionic formula

\sum (\text{ion charge} \times \text{number of ions}) = 0

Worked example

Write the formulas of aluminium sulfate and calcium nitride.

Practice this conceptself-check · 4 quick reps

Group 13 metals form 3+ ions

Aluminium, gallium and indium form

\mathrm{M^{3+}}

, so their oxides are

\mathrm{M_2O_3}

, their sulfates

\mathrm{M_2(SO_4)_3}

and their carbonates (where they exist)

\mathrm{M_2(CO_3)_3}

. A formula like

\mathrm{AlSO_4}

treats the metal as 2+ and is wrong.

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Formulas (1)

Typical ion charges from the group, and writing ionic formulas
Charge balance in an ionic formula
$\sum (\text{ion charge} \times \text{number of ions}) = 0$

Reference tables (2)

Groups, periods and blocks of the periodic table4 rows

Block	Groups	Subshell being filled	Examples
s-block	Groups 1 and 2 (and helium)	ns	Na, Mg, Ca
p-block	Groups 13 to 18	np	C, O, Cl, Ne
d-block	Groups 3 to 12	(n−1)d	Fe, Cu, Zn
f-block	Lanthanides and actinides, drawn as two rows below	(n−2)f	Ce, U

Metals, non-metals and metalloids3 rows

Class	Where in the table	Typical physical properties	Oxides	Examples
Metals	Left and centre, below the staircase	Shiny, conduct heat and electricity, malleable; mostly solids (mercury is a liquid)	Basic; a few amphoteric	Na, Mg, Fe, Cu
Non-metals	Right, above the staircase, plus hydrogen	Dull, poor conductors (graphite conducts), brittle as solids; many are gases	Acidic ( $\mathrm{CO_2}$ , $\mathrm{SO_2}$ ) or neutral (CO, $\mathrm{H_2O}$ )	C, N, O, S, Cl
Metalloids	Along the staircase	Look metallic but are brittle; semiconductors	Weakly acidic or amphoteric	B, Si, Ge, As

Watch out for (3)

Group 13 to 18 elements have the group number minus 10 valence electrons
→ Groups, periods and blocks of the periodic table
Aluminium is a metal, not a metalloid
→ Metals, non-metals and metalloids
Group 13 metals form 3+ ions
→ Typical ion charges from the group, and writing ionic formulas

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