Electronic Configurations & Types of Elements: s, p, d & f-block Elements

Table of Contents
1. Classification of the Elements
2. s-block elements
3. d-block elements (Transition elements)
4. f-block elements
5. Prediction of Period, Group and Block of an Element
6. Classification by Physical and Chemical Properties
7. Typical Elements
8. Diagonal Relationship
View more

Classification of the Elements

The classification of elements in the modern periodic table is based on the subshell in which the differentiating (last) electron is placed. On this basis, elements are classified into four blocks: s-block, p-block, d-block and f-block elements. Elements that have similar valence shell electronic configurations occur in the same group and therefore show similar chemical properties.

• More than 70-80% of the known elements are metals.
• Elements located at the top right of the periodic table are non-metals and are fewer in number (less than 20 in the long-form table among the main-group elements).
• In a group, metallic character increases down the group as atomic number increases.
• In a period, metallic character decreases from left to right.
• Chemical and physical properties vary regularly with atomic number (periodic law).
• Elements of the same group exhibit similar chemical properties because of similar valence shell electronic configuration.

s-block elements

When shells up to (n - 1) are completely filled and the last electron enters the s-orbital of the outermost (nth) shell, the elements are called s-block elements.

• Groups 1 and 2 (I A and II A) constitute the s-block.
• General electronic configuration: [inert gas] ns^1-2.
• s-block elements lie on the extreme left of the periodic table and are almost exclusively metals.
• They show typical metallic properties: high electropositivity, good electrical and thermal conductivity, and metallic lustre (except few exceptions like francium behavior is largely predicted).

Group I A (Alkali metals)

• These elements have configuration [inert gas] ns¹.
• They are highly reactive metals, form M+ ions, and react vigorously with water (forming hydroxides and hydrogen).
• They have low ionisation enthalpy and show increasing reactivity down the group.

Group II A (Alkaline earth metals)

Summary of various trends in Periodic Table.

• These have configuration [inert gas] ns².
• They are less reactive than alkali metals, form M2+ ions, and have higher ionisation enthalpies and higher melting points (generally) than alkali metals.

p-block elements

When shells up to (n - 1) are completely filled and the differentiating electron enters the p-orbital of the nth shell, the elements are called p-block elements.

• Groups 13 to 18 constitute the p-block.
• General electronic configuration: [inert gas] ns² np^1-6.
• p-block elements lie on the extreme right of the periodic table (except helium which is placed with the noble gases).
• This block includes metals, metalloids and non-metals; together with s-block elements they are called the representative (normal) elements.

Group III A (13) - Boron family

Group IV A (14) - Carbon family

Group V A (15) - Nitrogen family

Group VI A (16) - Oxygen family

d-block elements (Transition elements)

When the outermost (nth) and penultimate ((n - 1)th) shells are incompletely filled and the differentiating electron enters the (n - 1) d-orbitals, the elements are called d-block elements or transition elements.

• Groups 3 to 12 constitute the d-block.
• General electronic configuration: [inert gas] (n - 1)d^1-10 ns^0-2.
• d-block elements are commonly referred to as transition elements when they have partially filled d-orbitals in their neutral state or in any of their stable oxidation states.
• d-block elements are arranged in four series corresponding to filling of (n - 1)d orbitals:
  • 3d series: Sc (21) to Zn (30)
  • 4d series: Y (39) to Cd (48)
  • 5d series: La/Hf region to Hg (depends on inclusion of lanthanides)
  • 6d series: includes heavier homologues (partly theoretical/less common)

f-block elements 

• When the n, (n - 1) and (n - 2) shells are incompletely filled and the last electron enters into an f-orbital of the antepenultimate ((n - 2)th) shell, the elements are called f-block elements.
• General electronic configuration: (n - 2)f^1-14 (n - 1)d^0-1 ns².
• These elements are metals and are sometimes referred to collectively as inner-transition elements because their differentiating electrons occupy inner f-orbitals rather than the outermost orbital.
• They are placed separately (as two rows) below the main body of the periodic table to avoid undue expansion of the table; they are often associated with group 3 in placement.
• Properties within each series are quite similar, hence the lanthanides were historically referred to as the "rare earths" because their oxides were rare in early mineral samples.

• The two series of f-block elements are:
  • Lanthanide series (4f-series): fourteen elements from ₅₈Ce to ₇₀Lu; electrons progressively fill the 4f subshell.
  • Actinide series (5f-series): fourteen elements from ₉₀Th to ₁₀₃Lr; electrons progressively fill the 5f subshell.

Prediction of Period, Group and Block of an Element

• The block of an element corresponds to the subshell which receives the last electron (s, p, d or f).
• The period of an element is the principal quantum number (n) of the outermost shell (highest occupied energy level).
• The groupnumber can be deduced from the valence electrons (general rules):
  • For s-block elements: Group number = number of valence electrons (for representative s-block elements, 1 or 2).
  • For p-block elements: Group number = 10 + number of valence electrons (for main-group labelling in the 18-group system).
  • For d-block elements: Group number = number of (n - 1)d electrons + number of ns electrons (counting outer-shell electrons contributing to chemistry).

Classification by Physical and Chemical Properties

Elements are also commonly grouped as metals, non-metals and metalloids (semimetals) according to characteristic physical and chemical properties.

Metals

• Metals characteristically lose electrons to form cations and exhibit metallic lustre.
• They constitute more than 75% of the elements and are placed on the left and centre of the periodic table.
• Most metals are solids at room temperature (exception: mercury) and have high melting and boiling points.
• Metals are good conductors of heat and electricity.
• Oxides of metals are generally basic (some high oxidation state metal oxides are amphoteric or acidic; e.g., CrO₃ is acidic).

Non-metals

• Non-metals tend to gain electrons to form anions and are located at the top right of the periodic table.
• They are usually gases or solids with low melting and boiling points and are poor conductors of heat and electricity.
• Oxides of non-metals are generally acidic in nature.

Metalloids (Semimetals)

• Some elements along the dividing line between metals and non-metals exhibit properties intermediate between the two and are called metalloids or semimetals.
• Typical metalloids are boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb) and tellurium (Te).
• Oxides of metalloids are generally amphoteric (showing both acidic and basic character).

Typical Elements

• Elements of the third period are often called typical elements of their respective groups because the third-period element generally shows trends similar to heavier members of the group; example: Na, Mg, Al, Si, P, S, Cl.
• For many chemical predictions, the properties of the third-period element are used as a reference for the group (for example, Na is often used to characterise properties of alkali metals rather than Li in general statements about the group), because the second-period elements show some special behaviour due to their small size and absence of d-orbitals.
• Second-period elements differ from their heavier congeners in many respects because of smaller atomic size, higher ionisation enthalpies and lack of available d-orbitals for bonding.

Diagonal Relationship

Certain elements of the second period resemble elements of the third period of the next group in properties; this is called the diagonal relationship. Common examples are the pairs Li-Mg and Be-Al. The diagonal relationship arises because of several comparable properties between the diagonally related elements.

• Reasons for diagonal relationship include:
  • Similar atomic and ionic sizes (e.g. atomic radii Li = 1.23 Å, Mg = 1.36 Å; ionic radii Li⁺ = 0.60 Å, Mg²⁺ = 0.65 Å).
  • Comparable electropositive/electronegative character.
  • Similar polarising power (charge/radius ratio).
  • Similar electronegativity values (e.g. Li = 1.0, Mg = 1.2; Be = 1.5, Al = 1.5 on Pauling scale approximations).

Similarities between Li and Mg

• Both react directly with nitrogen to form nitrides: Li₃N and Mg₃N₂, whereas other alkali metals do not readily form stable nitrides.
• Fluorides, carbonates and phosphates of Li and Mg are relatively insoluble in water, while corresponding salts of heavier alkali metals are more soluble.
• Both Li and Mg are comparatively hard metals, whereas other alkali metals are soft.
• LiOH and Mg(OH)₂ are weak bases compared to the strong hydroxides of the heavier alkali metals.
• Metallic bonding in Li and Mg is stronger compared to other alkali metals, reflecting higher melting and boiling points.
• Thermal decomposition of lithium and magnesium nitrates yields the corresponding oxides: LiNO₃ → Li₂O and Mg(NO₃)₂ → MgO on suitable heating pathways.
• Thermal stability of Li₂CO₃ and MgCO₃ is lower compared to carbonates of heavier alkali metals and they decompose more readily to give CO₂.

Similarities between Be and Al

• Both Be and Al do not impart a coloured flame in a Bunsen burner test, unlike many other metals that give coloured flames due to d-orbital transitions.
• Both are relatively resistant to oxidation in air (form protective oxide layers), making them comparatively stable.
• Both are insoluble in liquid ammonia and therefore do not form the typical deep-blue solvated electron solutions seen with alkali metals in liquid ammonia.
• Neither readily forms peroxides or superoxides under ordinary conditions.
• Both have relatively low reducing power compared to more electropositive metals (consistent with their standard electrode potentials).
• Be and Al form bridged halide structures (halogen bridging) in certain covalent halide complexes.