The Structure of the Periodic Table
- The periodic table is organized into periods, groups, and blocks.
- Each of these helps us understand how an element’s electronic structure links to its properties.
Periods: Horizontal Rows
The horizontal rows of the periodic table are called periods.
- Each period number corresponds (for the main-group elements) to the highest main energy level (principal quantum number, n) occupied by electrons in atoms of those elements.
- As you move across a period from left to right:
- The number of protons in the nucleus increases.
- This stronger nuclear attraction pulls electrons closer to the nucleus.
- This leads to trends such as decreasing atomic radius and increasing ionization energy across the period.
Period Number and Principal Energy Levels
Definition
Period number
The period number of an element tells you the highest principal energy level (n) that contains electrons.
- Period 1: Elements like hydrogen (H) and helium (He) have electrons only in the n = 1 level (1s).
- Period 2: Elements like lithium (Li) and neon (Ne) have electrons in the n = 1 and n = 2 levels (1s, 2s, 2p).
- Period 4: Elements like potassium (K) and krypton (Kr) have electrons in levels from n = 1 up to n = 4.
Example
- Let’s analyse bromine (Br), which is in Period 4 and Group 17:
- Period 4 → its outermost electrons are in the 4th energy level (n = 4).
- Group 17 → it has 7 valence electrons.
- Electron configuration of bromine: $$\text { Br: } 1 s^2 2 s^2 2 p^6 3 s^2 3 p^6 3 d^{10} 4 s^2 4 p^5$$
- The valence electrons are the 4s² 4p⁵ electrons (7 total).
Groups: Vertical Columns
- The vertical columns of the periodic table are called groups.
- Elements in the same group have the same number of valence electrons.
- This explains why they have similar chemical properties and show similar patterns of reactivity.
Group Numbers and Valence Electrons
Definition
Group number
The group number of a main-group element indicates the number of valence electrons (electrons in the outermost energy level).
- Valence electrons are crucial because they determine an element’s chemical reactivity and bonding behaviour.
- Group 1 (alkali metals): 1 valence electron → ns¹
- Group 2 (alkaline earth metals): 2 valence electrons → ns²
- Groups 13–18 (p-block):
- Number of valence electrons = group number – 10
- Group 15 → 5 valence electrons → ns² np³
- Group 17 → 7 valence electrons → ns² np⁵
Example
- Chlorine is in Group 17, so it has 7 valence electrons.
- Its electron configuration is: $$\mathrm{Cl}: 1 s^2 2 s^2 2 p^6 3 s^2 3 p^5$$
- The valence electrons are the 3s² 3p⁵ electrons (7 in total).
- Chlorine is highly reactive because it needs only 1 more electron to achieve a stable octet (8 electrons) in its outer shell.
Note
- For transition metals (discussed in the later articles), the group number (in IB or modern notation) corresponds to the total number of $(n–1)d$ and $ns$ electrons in the outermost two energy levels.
- This becomes more important in IB when discussing variable oxidation states.
Chemical Similarity In a Group
Because elements in a group have the same number of valence electrons, they often:
- Form ions with the same charge (e.g. Group 1 always forms +1 ions).
- Show similar reactions with other substances.
Example
- All Group 1 elements (alkali metals) have 1 valence electron and react vigorously with water to form a metal hydroxide and hydrogen gas.
- All Group 17 elements (halogens) have 7 valence electrons and typically form –1 ions (e.g. Cl⁻, Br⁻, I⁻).
Example
Consider phosphorus (P), which is in Group 15 and Period 3.
- Group 15 → 5 valence electrons.
- Period 3 → valence electrons are in the third energy level (n = 3).
Blocks: s, p, d, and f
The periodic table can also be divided into blocks, based on which type of atomic orbital is being filled with electrons.
s-block
- Groups 1 and 2 (plus hydrogen) make up the s-block.
- The outermost electrons are in an s-subshell.
- Their electron configurations end in ns¹ or ns².
p-block
- Groups 13 to 18 form the p-block.
- This block includes non-metals, metalloids, and some metals.
- The outermost electrons are in a p-subshell.
- Their configurations end in ns² np¹–np⁶.
d-block
- Groups 3 to 12 are the d-block, containing the transition metals.
- Here, the (n–1)d subshell is being filled (e.g. 3d, 4d).
- These elements show typical transition metal behaviour (variable oxidation states, coloured compounds, complex ion formation).
f-block
- The f-block contains the lanthanoids (lanthanides) and actinoids (actinides), shown as the two rows at the bottom of the table.
- The (n–2)f subshell is being filled (e.g. 4f, 5f).
- Many of these elements are used in magnets, lasers, and nuclear applications.
What is Electron Configuration?
Note
- Note that electron configurations will be fully revisited in the Matter section.
- In the given article, it is discussed in relation to periodic table, namely, how it can be derived based on the structure of the periodic table.
Definition
Electron configuration
It describes the arrangement of an atom’s electrons in its energy levels, sublevels and orbitals.
- Electrons occupy regions around the nucleus called orbitals.
- Orbitals are grouped into sublevels: s, p, d, f.
- Sublevels belong to energy levels, labelled by principal quantum numbers: n = 1, 2, 3, …
Key rule – Aufbau principle
- Electrons fill orbitals in order of increasing energy, following the Aufbau principle (“build up”):
- The 1s orbital is filled first, then 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.
- Maximum number of electrons per sublevel:
- s sublevel → 1 orbital → 2 electrons max
- p sublevel → 3 orbitals → 6 electrons max
- d sublevel → 5 orbitals → 10 electrons max
- f sublevel → 7 orbitals → 14 electrons max
Tip
- The period number tells you the principal energy level (n) of the valence electrons.
- The block tells you which sublevel (s, p, d, f) is being filled.
- The group (for main-group elements) tells you the number of valence electrons.
