Electronic Configuration and Types of Elements: Definition, Configuration, Anomalies and Examples

• Electronic Configuration: In group 1 elements, the last electron enters an ns¹ orbital. For those in group 2, it enters an ns² orbital. The letter n stands for the principal energy level, which grows larger as you move down the group.
• Valency: Group 1 metals lose a single electron to form a +1 ion, whereas Group 2 metals shed two electrons to yield a +2 ion. Because ejecting electrons from their outer shell requires less energy than acquiring additional ones, these elements consistently generate positively charged ions.
• Physical Properties: s-block metals are usually soft, have low melting and boiling points because their metallic bonds aren't very strong. For example, sodium melts at just 98°C, which is quite low when compared to something like iron.
• Chemical Properties: These elements are generally very reactive, especially those in group 1. Since they have low ionization energies, they lose electrons easily. Most form ionic compounds, but beryllium and magnesium sometimes form covalent bonds. They’re also known as strong reducing agents. Plus, many s-block elements produce flame colors—sodium glows yellow, and potassium burns with a violet flame.
• Conductivity: s-block metals are great conductors of heat and electricity. That is because their outer electrons move freely, allowing energy to pass through the metal without much resistance.

Elements belonging to group 13-18 are in the p-block. The last electron of these electrons enters the p-orbital of valence shell. This block includes metals (such as Al, Sn), non-metals (e.g., C, N, O), metalloids (such as Si, As), as well as noble gases (e.g., Ne, Ar). JEE Main and CUET aspirants need to learn about these electronic configurations to answer conceptual questions in the exam.

Characteristics of p-Block Elements

The following are the characteristics of p-block elements:

• Electronic Configuration: The general configuration for p-block elements is ns² np¹–⁶, indicating that they have between 2 and 6 electrons in the p-orbital of their outer shell.
• Diversity: This block includes a wide range of element types  such as metals, non-metals, and metalloids. The metallic character decreases as you move across a period but increases down a group.
• Ionization Energy: p-block elements generally have higher ionization energies than those in the s-block. This energy increases across a period due to the stronger nuclear charge and decreases down a group because of the larger atomic size.
• Bonding: Most non-metals in the p-block form covalent bonds. Some p-block metals, however, are capable of forming ionic compounds. Noble gases tend to be unreactive because their outermost np⁶ orbitals are completely filled, giving them a naturally stable electron configuration that resists change. Some metals in this block, however, can also create ionic compounds. As for the noble gases, they hardly react at all. That’s because their outermost np⁶ orbitals are already full, making their electron setups very stable and leaving little drive for chemical change.
• Oxidation States: p-block elements can exhibit both positive and negative oxidation states, unlike Fluorine, which only shows a negative state. A good example is Nitrogen, which shows –3 in ammonia (NH₃) and +5 in nitric acid (HNO₃).
• Oxidizing Power: The ability to act as oxidizing agents increases across a period. This is due to rising electron affinity. A good example is the halogens, such as Chlorine, which readily gain electrons in reactions.

Example

Chlorine (atomic number 17) has the configuration ${1s}^2{2s}^2{2p}^6{3s}^2{3p}^5$ . Its last electron enters 3 p orbital. This is why Chlorine is in p-block, group 17, and period 3.

The d-block includes those elements whose last electron enters the (n-1)d orbital from groups 3 to 12. These are known as the transition elements, with four series: $3\mathrm{d}(\mathrm{S}\mathrm{c}$ to Zn to Cd, 5 d ( $\mathrm{L}\mathrm{a},\mathrm{H}\mathrm{f}$ to Hg ), and incomplete 6d series. In total, there are 39 d -block elements. NCERT solutions of this chapter covers questions related to d-block elements and students must practice them for better performance in the CBSE board exams.

Example: The atomic number of Copper is 29 and has the configuration $\left[\mathrm{Ar}\right]{3d}^{10}{4s}^1$ . Its last electron enters the 3d orbital due to which it belongs to d-block, group 11, period 4.

The f-block elements are also known as the inner transition elements. These include elements where the last electron enters the ( $\mathrm{n}-2$  )f orbital. It consists of two series including lanthanides ( Ce to $\mathrm{L}\mathrm{u},4\mathrm{f}$  ) and actinides ( Th to $\mathrm{L}\mathrm{r},5\mathrm{f}$ ). Both of the series consist of 14 elements each. IISER and NEET students must practice conceptual questions based on different types of elements and their electronic configuration.

Characteristics of f-Block Elements

• Electronic Configuration: General form is $(n-2)f^{1-14}(n-1)d^{0-2}n{s}^2$ . For example, Cerium ( $Z=58$ ) is [ $Xe$ ] $4f^1\mathrm{}5d^1\mathrm{}6s^2$ .
• Physical Properties: All are metals, with lanthanides called rare earths and most actinides being transuranic (synthetic). They have high densities (e.g., Iridium at $22.61{\mathrm{g}\mathrm{c}\mathrm{m}}^{-3}$ .
• Chemical Properties: Show variable valency, form colored ions, and actinides are radioactive. They form complexes due to high charge and small ionic radii.
• Lanthanide Contraction: Poor shielding by 4 f electrons causes a decrease in atomic radii across the lanthanide series, impacting d-block elements in period 6.

Example

The atomic number of Cerium id 58 and its configuration is [Xe] $4f^{1}5d^{1}6s^2$ . Due to this configuration, it is in the f block, group 3 and period 6.

To determine an elements block, group, and period the following must be done:

• Period: This is correspondent to the principal quantum number ( n ) of valence shell.
• Block: It is determined by the subshell (s, p, d, or f) in which the last electron enters.
• Group Number:
  • s-block: Group number = number of s-electrons (1 or 2).
  • p-block: Group number $=10+(\mathrm{s}+\mathrm{p}$ electrons in the valence shell).
  • d-block: Group number $=(\mathrm{n}-1)\mathrm{d}+\mathrm{n}\mathrm{s}$ electrons.
  • f-block: In most of the elements, this is group 3, as lanthanides and actinides are placed separately.

The following trends and anamolies have been observed in the elements of different blocks:

• s-Block: These elements react quickly because their outermost electrons are loosely held. You’ll notice that group 1 metals—like sodium and potassium—tend to be even more reactive than those in group 2, such as magnesium or calcium.
  
• p-Block: As you move across a row in the periodic table, the elements shift from metallic to non-metallic in nature. Near the end of each row, the noble gases appear. They don’t react much because their electron shells are already full, leaving them with no reason to bond.
  
• d-Block: This group is full of metals that can show more than one charge in chemical reactions. That’s what we call variable oxidation states. They’re also common in reactions as catalysts, mainly because of how their d-orbitals work with electrons.
  
• f-Block: The lanthanide contraction is a quiet but important trend here. It causes these elements to stay similar in size, even as their atomic number increases. Because of that, many of them behave almost alike in chemical reactions.

This topic aims to explain students why any element falls into either s, p, d or f block. By understanding electronic configurations and the nature of the element you can easily understand the chemical behavior and bonding type. Entrance examinations in the country include conceptual questions based on your knowledge of the topic. It is therefore, suggested to first understand the concept and then practice questions.