Edited By Shivani Poonia | Updated on Jul 02, 2025 05:58 PM IST
The arrangement of elements in the periodic table is based on their electronic configurations, which determine their positions within specific periods and groups. This structure reflects the atomic number of each element, indicating the number of protons and electrons in a neutral atom. Elements within the same period show a gradual change in their outer electron shells, resulting in patterns in chemical behavior. Within a group, elements exhibit similar valence electron configurations, especially in the outermost shell, leading to comparable chemical properties.
This Story also Contains
The Period in the Modern Periodic Table-
Solved Examples Based On Electronic configuration in periods and groups-
Conclusion
Electronic Configuration in Periods and Groups
For example, all elements in Group 1, known as alkali metals, possess a single electron in their outermost shell. This makes them highly reactive, particularly with water and halogens. On the other hand, Group 18 elements—also called noble gases—have a complete set of electrons in their outermost shell, making them stable and largely unreactive under standard conditions.
This topic falls under the broader subject of "Classification of Elements and the Periodic Table," a fundamental chapter in Class 11 Chemistry. It is crucial not only for board examinations but also for competitive entrance exams such as JEE Main, NEET, SRMJEE, BITSAT, WBJEE, BCECE, and others. From 2013 to 2023, a total of four questions based on this concept appeared in competitive exams—three in JEE between 2020 and 2023 and one in NEET in 2017.
The Period in the Modern Periodic Table-
A horizontal row in the periodic table is called a period. The periodic table is divided into seven periods, each of which starts at the far left. When electrons start to fill a new primary energy level, a new period starts. There are eight elements in periods 2 and 3, compared to only two in period 1 (helium and hydrogen). This is the term defined for the modern periodic table. let us see some examples based on the same topic.
NEET Highest Scoring Chapters & Topics
This ebook serves as a valuable study guide for NEET exams, specifically designed to assist students in light of recent changes and the removal of certain topics from the NEET exam.
The period of any element is determined by the last shell in which the last electrons enter. For example, Fe's atomic number is 26. The electronic configuration can be written as:
1s22s22p63s23p64s23d6
Now, the last electron enters into the d-subshell but electrons are also present in the 4th subshell. Therefore, Fe belongs to the fourth period.
The Block in the Modern Periodic Table-
The atomic orbitals in which an element's valence electrons or vacancies are located unite the elements in a block of the periodic table. Charles Janet appears to have coined the phrase initially. A collection of elements found in nearby element groups is called an element block. The term was first used (in French) by Charles Janet. The spectroscopic lines of atomic orbitals that were described as sharp, primary, diffuse, and fundamental gave rise to the block names s, p, d, and f. This phrase "block" appears in the current periodic table.
The block of any element is determined by the last subshell in which the last electron enters. For example, Na has atomic number 11, thus its electronic configuration can be written as:
1s22s22p63s1
Now, its last electron enters into the s-subshell, therefore, Na belongs to the s-block.
Group
The group of any element is determined in different ways.
For s-block
If the last electron of any element enters the s-subshell, then the group number is equal to the number of electrons in the last s-subshell
For p-block
If the last electron of any element enters into the p-subshell, then the group number is equal to (12 + the number of electrons in the last p-subshell).
For d-block
If the last electron of any element enters into the d-subshell, then the group number is equal to (2 + the number of electrons in (n-1)d-subshell.
For f-block
There are only two series of f-block i.e., lanthanide and actinide. If the last electron of any element enters into the f-subshell and if the atomic number is between 57-71, then the element belongs to the lanthanide series i.e., 6th period. Further, if the last electron of any element enters into the f-subshell and if the atomic number is between 89-103, then the element belongs to the actinide series i.e., 7th period. All the elements from both these series belong to group 3.
Recommended topic video on ( Electronic configuration in Periods and Groups)
Solved Examples Based On Electronic configuration in periods and groups-
Example 1:Choose the correct option:
1) The period of the element is determined by its highest shell
2) The period of the element is determined by its last orbital
3) The period of the element is determined by its valence shell electrons
4) The period of the element is determined by its valency
Solution: The period of the element is determined by its highest shell.
Hence, the answer is the option (1).
Example 2: An element with atomic number 23 belongs to which period?
1) 3rd period
2) 2nd period
3) 5th period
4) 4th period
Solution: we have to check the last shell. The shell with a higher value of the principal quantum number(n) is the period for that element. The electronic configuration of an element having atomic number 23, is = [Ar] 3d3 4s2. The shell has a higher value of n is 4s. Therefore the period is = 4.
Hence, the answer is the option (4).
Example 3: K has atomic number 19, thus its electronic configuration can be written as:
1s22s22p63s23p64s1
Now it has 1 electron in the s-subshell, therefore K belongs to Group 1.
Example 4: Ge has atomic number 32, thus its electronic configuration can be written as:
1s22s22p63s23p64s23d104p2
Now, it has 2 electrons in the last p-subshell, therefore its group number is:
12 + 2 = 14
Thus, Ge belongs to Group 14
Example 5: Mn has the atomic number 25, thus its electronic configuration can be written as:
1s22s22p63s23p64s23d5
Now it has 5 electrons in the d-subshell, therefore its group number is:
2 + 5 = 7
Thus, Mn belongs to Group 7.
Example 6: The element with atomic number 35 belongs to which group:
1) 13th group
2) 17th group
3) 16th group
4)2nd group
Solution: The electronic configuration of the element with atomic number 35 is: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
Now, the group of the element is determined by the following formulae:
(i) For s-block: ns1 and ns2
(ii) For p-block: 12 + number of electrons in the P-subshell of the last shell.
=12 + 5 =17th group
(iii) For d-block: 2 + number of electrons in (n-1)d subshell.
Hence, the answer is the option (2).
Practice more Questions from the link given below:
Understanding how electrons are arranged in atoms reveals why elements behave chemically as they do. The distribution of electrons across shells and subshells—shaped by principles like Aufbau, Pauli, and Hund—determines an element’s tendency to lose, gain, or share electrons, thereby influencing bond formation and stability. As electrons fill successive orbitals, predictable trends emerge: atomic size decreases and ionization energy increases across a period, while the opposite occurs down a group . This foundational concept not only explains chemical behavior but also guides technological advances—such as semiconductor engineering, catalysis, and the development of novel magnetic and pharmaceutical materials .
Frequently Asked Questions (FAQs)
1.What is electronic configuration?
It’s the arrangement of an element’s electrons across its atomic orbitals. This layout determines the element's position on the periodic table—period, group, and block—and explains its chemical reactivity and properties.
2.What is electronic configuration?
Electronic configuration is the arrangement of electrons in an atom's orbitals. It describes how electrons are distributed in different energy levels and sublevels around the nucleus, following specific rules based on quantum mechanics.
3.How to write electronic configuration for p-block elements.
If the last electron of any element enters into the p-subshell, then the group number is equal to (12 + the number of electrons in the last p-subshell).
4.How does the period number relate to electron configuration?
The period number equals the principal quantum number (n) of the outermost shell.
Period 1 has n = 1 (fills 1s).
Period 2 has n = 2 (fills 2s and 2p), and so on .
5.Why do periods contain different numbers of elements?
because each shell and subshell holds a specific maximum:
Shell n can hold up to 2n² electrons (e.g. n=2 → 8 electrons: 2s+2p)
6.Why do elements in the same group behave similarly?
They share the same number of valence electrons, resulting in matching outer-shell electron configurations—hence similar chemistry.
7.Why do periods contain different numbers of elements?
because each shell and subshell holds a specific maximum:
Shell n can hold up to 2n² electrons (e.g. n=2 → 8 electrons: 2s+2p)
8.Why do elements in the same group behave similarly?
They share the same number of valence electrons, resulting in matching outer-shell electron configurations—hence similar chemistry.
9.What is 7th period of periodic table?
It spans 32 radioactive elements from Francium (Fr, Z=87) to Oganesson (Og, Z=118), covering the s-, f-, d-, and p-blocks.
All elements are radioactive, with many being synthetic and extremely short-lived.
The period was completed in 2016 with the official recognition of superheavy elements Nihonium, Moscovium, Tennessine, and Oganesson
10.What is 2 8 8 18 18 32 rule?
The numbers 2, 8, 8, 18, 18, 32 describe the typical increase in total electron count when moving down a group, not the absolute capacity of each shell. Each jump corresponds to filling subshells in this sequence: s, p, d, f.
Essentially, electrons add in these increments (+2, +8, +8, +18, +18, +32) due to the Aufbau principle, which fills subshells in order of increasing energy (n + ℓ rule).
So reading down a column: you go from one element to the next by adding roughly these numbers of electrons in order, reflecting how shells and subshells actually fill in real atoms.
11.What is 7th period of periodic table?
It spans 32 radioactive elements from Francium (Fr, Z=87) to Oganesson (Og, Z=118), covering the s-, f-, d-, and p-blocks.
All elements are radioactive, with many being synthetic and extremely short-lived.
The period was completed in 2016 with the official recognition of superheavy elements Nihonium, Moscovium, Tennessine, and Oganesson
12.What is 2 8 8 18 18 32 rule?
The numbers 2, 8, 8, 18, 18, 32 describe the typical increase in total electron count when moving down a group, not the absolute capacity of each shell. Each jump corresponds to filling subshells in this sequence: s, p, d, f.
Essentially, electrons add in these increments (+2, +8, +8, +18, +18, +32) due to the Aufbau principle, which fills subshells in order of increasing energy (n + ℓ rule).
So reading down a column: you go from one element to the next by adding roughly these numbers of electrons in order, reflecting how shells and subshells actually fill in real atoms.
13.How does electronic configuration relate to an element's position in the periodic table?
An element's position in the periodic table is directly related to its electronic configuration. The period number corresponds to the highest occupied energy level, while the group number is determined by the number of valence electrons in the outermost shell.
14.Why do elements in the same group have similar chemical properties?
Elements in the same group have similar chemical properties because they have the same number of valence electrons in their outermost shell. These valence electrons largely determine an element's chemical behavior and reactivity.
15.What is the Aufbau principle, and how does it apply to electronic configuration?
The Aufbau principle states that electrons fill orbitals from lowest to highest energy levels. It guides the order in which electrons are added to atomic orbitals, helping to determine an element's electronic configuration.
16.How does Hund's rule affect electronic configuration?
Hund's rule states that electrons in the same sublevel (e.g., p orbitals) will occupy separate orbitals with parallel spins before pairing up. This minimizes electron repulsion and results in a more stable configuration.
17.What is the Pauli exclusion principle, and how does it relate to electronic configuration?
The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers. This means that each orbital can hold a maximum of two electrons with opposite spins, influencing how electrons are distributed in an atom.
18.Why do transition elements have similar properties despite having different electronic configurations?
Transition elements have similar properties because they all have partially filled d-orbitals. These d-orbitals allow for various oxidation states and complex formation, leading to similar chemical and physical properties across the series.
19.How does the concept of effective nuclear charge affect electronic configuration across a period?
Effective nuclear charge increases across a period as protons are added to the nucleus. This stronger attraction to the nucleus causes a contraction in atomic size and influences the energy levels of electrons, affecting the electronic configuration.
20.What is electron shielding, and how does it impact electronic configuration?
Electron shielding occurs when inner electrons partially shield outer electrons from the full nuclear charge. This effect reduces the attraction between the nucleus and outer electrons, influencing their energy levels and the overall electronic configuration.
21.How does the octet rule relate to electronic configuration?
The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons (like noble gases). This rule is based on the electronic configuration of elements and guides many chemical reactions and bonding patterns.
22.Why do some elements, like copper and chromium, have unexpected electronic configurations?
Some elements, like copper and chromium, have unexpected configurations due to the extra stability provided by half-filled or fully filled d-orbitals. This results in one electron from the s-orbital being promoted to the d-orbital, creating a more stable arrangement.
23.How does electronic configuration explain the formation of ions?
Electronic configuration explains ion formation by showing how atoms gain or lose electrons to achieve a stable noble gas configuration. Metals typically lose electrons to form cations, while non-metals gain electrons to form anions, based on their valence electron arrangement.
24.What is the relationship between electronic configuration and atomic radius?
Electronic configuration influences atomic radius. As you move across a period, the atomic radius generally decreases due to increased nuclear charge and electron-electron repulsion. Down a group, atomic radius increases as new energy levels are added.
25.How does electronic configuration affect an element's electronegativity?
Electronic configuration impacts electronegativity by determining the attraction of an atom for shared electrons. Elements with nearly full valence shells tend to have higher electronegativity, as they strongly attract electrons to complete their octet.
26.Why do noble gases have such stable electronic configurations?
Noble gases have stable electronic configurations because their outermost s and p orbitals are completely filled (except for helium, which has a full 1s orbital). This full outer shell makes them chemically inert and resistant to forming compounds.
27.How does the concept of energy levels relate to electronic configuration?
Energy levels in an atom are discrete shells where electrons reside. The electronic configuration describes how electrons are distributed among these energy levels, with each level having a specific capacity and sublevels (s, p, d, f orbitals).
28.What is the significance of valence electrons in electronic configuration?
Valence electrons, the outermost electrons in an atom, are crucial in electronic configuration as they determine an element's chemical properties, bonding behavior, and reactivity. They are the electrons involved in chemical reactions and bond formation.
29.How does electronic configuration explain the periodic trends in ionization energy?
Electronic configuration explains ionization energy trends by showing how tightly electrons are held. As you move across a period, ionization energy generally increases due to increased nuclear charge. Down a group, it decreases as valence electrons are further from the nucleus.
30.What role does electronic configuration play in determining an element's metallic or non-metallic character?
Electronic configuration determines metallic or non-metallic character by the number and arrangement of valence electrons. Metals have few valence electrons and tend to lose them easily, while non-metals have nearly full valence shells and tend to gain electrons.
31.How does the concept of subshells (s, p, d, f) relate to electronic configuration?
Subshells (s, p, d, f) represent different types of orbitals within each energy level. The electronic configuration describes how electrons fill these subshells, following specific patterns and energy orders, which ultimately determine an element's properties.
32.Why do elements in the same period have different properties despite having the same number of electron shells?
Elements in the same period have different properties because, while they have the same number of electron shells, they differ in the number of valence electrons and protons. This leads to variations in electron-nucleus interactions and chemical behavior.
33.How does electronic configuration explain the formation of chemical bonds?
Electronic configuration explains bond formation by showing how atoms share, gain, or lose electrons to achieve stable configurations. The type and strength of bonds formed depend on the valence electron arrangements of the participating atoms.
34.What is the connection between electronic configuration and an element's oxidation states?
Electronic configuration determines possible oxidation states by showing how many electrons an atom can lose, gain, or share. Elements with multiple available electrons in their outer shells can often exhibit various oxidation states.
35.How does the Madelung rule (n + l rule) guide the filling of orbitals in electronic configuration?
The Madelung rule states that orbitals are filled in order of increasing n + l value (where n is the principal quantum number and l is the azimuthal quantum number). This rule helps predict the order in which electrons fill orbitals in multi-electron atoms.
36.Why do some elements have more than one possible electronic configuration?
Some elements can have multiple electronic configurations due to the close energy levels of certain orbitals. This allows for electron promotion between orbitals, resulting in different arrangements that may offer slightly different stabilities.
37.How does electronic configuration explain the color of transition metal compounds?
The color of transition metal compounds is explained by their electronic configuration, specifically the partially filled d-orbitals. These allow for electron transitions between d-orbitals when exposed to light, absorbing certain wavelengths and reflecting others, producing color.
38.What is the relationship between electronic configuration and an element's magnetic properties?
Electronic configuration determines an element's magnetic properties based on the presence of unpaired electrons. Elements with unpaired electrons in their d or f orbitals often exhibit paramagnetism, while those with paired electrons are typically diamagnetic.
39.How does electronic configuration influence the formation of complex ions?
Electronic configuration influences complex ion formation by determining the availability of empty orbitals and the number of valence electrons. Transition metals, with their partially filled d-orbitals, are particularly prone to forming complex ions.
40.Why do lanthanides and actinides have similar chemical properties despite different electronic configurations?
Lanthanides and actinides have similar properties because their differentiating electrons are added to inner f-orbitals, which are shielded by outer electrons. This results in minimal changes to their chemical behavior despite different electronic configurations.
41.How does electronic configuration explain the phenomenon of electron affinity?
Electronic configuration explains electron affinity by showing how easily an atom can accept an additional electron. Elements with nearly full valence shells typically have high electron affinities, as adding an electron completes their octet.
42.What is the significance of core electrons in electronic configuration?
Core electrons, the inner electrons not involved in chemical bonding, play a crucial role in electronic configuration by shielding outer electrons from the nuclear charge. This shielding effect influences the energy levels and behavior of valence electrons.
43.How does electronic configuration relate to the concept of effective nuclear charge?
Electronic configuration directly relates to effective nuclear charge by determining the degree of electron shielding. The arrangement of electrons in various orbitals affects how much nuclear charge the outer electrons experience, influencing atomic properties.
44.Why do elements in the same group have similar valence electron configurations but different total electron counts?
Elements in the same group have similar valence configurations because they have the same number of outer electrons. However, their total electron counts differ due to additional inner electrons in lower energy levels, which increase as you move down the group.
45.How does electronic configuration explain the trend in atomic size across the periodic table?
Electronic configuration explains atomic size trends by showing how electron shells are filled. Across a period, atomic size generally decreases due to increased nuclear charge pulling electrons closer. Down a group, size increases as new electron shells are added.
46.What is the relationship between electronic configuration and an element's first ionization energy?
Electronic configuration directly influences first ionization energy. Elements with stable configurations (like noble gases) have high ionization energies, while those with single electrons in their outer s or p orbitals (like alkali metals) have low ionization energies.
47.How does the concept of hybridization relate to electronic configuration?
Hybridization is the mixing of atomic orbitals to form new hybrid orbitals. It's based on the electronic configuration of an atom and explains molecular geometry and bonding in compounds that don't follow the octet rule, like in carbon's sp3 hybridization.
48.Why do transition metals often have variable oxidation states?
Transition metals often have variable oxidation states due to their electronic configuration, specifically their partially filled d-orbitals. These d-electrons can be easily lost or shared, allowing the atom to form compounds with different oxidation numbers.
49.How does electronic configuration explain the formation of coordinate covalent bonds?
Electronic configuration explains coordinate covalent bonds by showing how an atom with a lone pair of electrons (the donor) can share both electrons with another atom (the acceptor) that has an empty orbital. This is common in complex ion formation.
50.What is the connection between electronic configuration and an element's melting and boiling points?
Electronic configuration influences melting and boiling points by determining the strength of interatomic forces. Elements with more valence electrons or stronger metallic bonds (due to delocalized electrons) tend to have higher melting and boiling points.
51.How does electronic configuration relate to the concept of electronegativity?
Electronic configuration relates to electronegativity by determining an atom's ability to attract shared electrons. Elements with nearly full valence shells and small atomic radii tend to have high electronegativity due to their strong attraction for electrons.
52.Why do some elements violate the Aufbau principle in their ground state electronic configuration?
Some elements violate the Aufbau principle due to the extra stability gained from having half-filled or fully filled subshells, particularly in d and f orbitals. This can lead to unexpected electron arrangements that are more energetically favorable.
53.How does electronic configuration explain the formation of molecular orbitals?
Electronic configuration forms the basis for molecular orbital theory. When atoms combine, their atomic orbitals overlap to form molecular orbitals. The number and type of orbitals involved depend on the electronic configurations of the combining atoms.
54.What is the significance of the noble gas core in writing electronic configurations?
The noble gas core notation simplifies writing electronic configurations by using the symbol of the previous noble gas in brackets, followed by the remaining electron configuration. This shorthand emphasizes the valence electrons, which are most important for chemical behavior.
55.How does electronic configuration influence an element's reactivity?
Electronic configuration influences reactivity by determining how easily an atom can gain, lose, or share electrons. Elements with nearly empty or nearly full valence shells are often more reactive as they seek to achieve a stable octet configuration.
56.Why do some elements form multiple bonds, and how is this related to their electronic configuration?
Some elements form multiple bonds due to their electronic configuration allowing for the sharing of more than one pair of electrons. This is common in elements like carbon, nitrogen, and oxygen, which can form double or triple bonds to achieve stable octet configurations.
57.How does electronic configuration explain the diamagnetic or paramagnetic nature of elements?
Electronic configuration determines whether an element is diamagnetic or paramagnetic based on the presence of unpaired electrons. Atoms with all paired electrons are diamagnetic, while those with unpaired electrons are paramagnetic due to their magnetic moments.
58.What is the relationship between electronic configuration and the formation of ionic compounds?
Electronic configuration guides ionic compound formation by showing how atoms can gain or lose electrons to achieve stable noble gas configurations. Metals tend to lose electrons to form cations, while non-metals gain electrons to form anions, based on their valence electron arrangements.
59.How does electronic configuration explain the trend in electron affinity across the periodic table?
Electronic configuration explains electron affinity trends by showing how easily an atom can accept an additional electron. Generally, electron affinity increases across a period due to increasing nuclear charge, and decreases down a group as the atomic size increases.
60.Why do some elements have anomalous electronic configurations, and what are the implications?
Some elements have anomalous configurations due to the increased stability of half-filled or fully filled subshells. This can lead to unexpected chemical properties and reactivity, as seen in elements like copper and chromium, which have one electron in their outermost s-orbital instead of two.
61.How does electronic configuration relate to the concept of quantum numbers?
Electronic configuration is directly related to quantum numbers, which describe the energy, angular momentum, magnetic moment, and spin of electrons in an atom. The arrangement of electrons in various orbitals is determined by these quantum numbers, following specific rules and principles.
