Electronic Configuration of First 30 Elements - Meaning, Definition, Elements, Functions, FAQs

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Electronic configuration is an organized alignment of electrons in atomic orbitals according to the energy levels. This arrangement of electrons is directed by concepts like Hund's rule, Pauli exclusion, and the Aufbau rule. The location and the spin of the electron are expressed by its quantum number (n, l, mₗ, mₛ). Electronic configuration is helpful in determining the chemical properties, reactivity, and position in the periodic table. It also influences what type of bond the elements will form, the valency, and the ability to donate or accept electrons during a chemical reaction. The trend in the periodic table, like ionization enthalpy, electronegativity, and atomic size, can also be explained through this arrangement of electrons. The following electronic configurations of the first 30 elements are mentioned, which will help in understanding the trends it obeys.

This Story also Contains

Let's Understand Electronic Configuration
How to Write the Electronic Configuration of Elements
What are the three rules for writing electronic configurations of elements?
What Function Does Electronic Configuration of Elements Play?
Some Solved Examples

Electronic Configuration of First 30 Elements - Meaning, Definition, Elements, Functions, FAQs

Let's Understand Electronic Configuration

An electronic Configuration, also known as an electronic structure, is the arrangement of electrons at different energy levels surrounding an atomic nucleus. The electronic configuration of a molecule is the distribution of electrons in distinct molecular orbitals. The importance of the molecule cannot be overstated. It is possible to determine the number of electrons in bonding and antibonding molecular orbitals from a molecule's or molecular ion's electronic configuration.

The electrical configuration of an element is used to figure out where electrons are located in that element.
From the lowest to the highest energy level, electrons are arranged in ascending order.
The electrical configuration of an element is largely determined by its atomic number.
The electrical configuration of an atom is helpful in determining an element's valency, which aids in determining the element's reactivity.
The atomic spectra can also be interpreted using the electrical configuration.
Noble gases with totally filled outermost electrons, such as Neon, Argon, and Helium, are the most stable. Noble gases have filled valence shells, which give them their inertness.
Copper and chromium have a peculiar electrical structure in which the 3d- orbitals are filled first, rather than the 4s orbitals.
In chromium([Ar] 3d⁵ 4s¹) the d-orbital, which is filled with single electrons, boosts the atom's stability. Similarly, the d-orbital of Copper [Ar] 3d¹⁰ 4s¹ is totally filled with paired electrons, ensuring the atomic structure's stability.

How to Write the Electronic Configuration of Elements

The rules mentioned below are to be followed while writing electronic configurations. To get different quantum numbers, we first have to extract various information, such as the number of electrons, possible number of various shells and orbitals, energy levels, etc. of elements by using the periodic table.

The electronic configuration is as follows: 1s² 2s² 2p⁶ 3s²3p⁶4s¹. Its 19 electrons can be divided into shells in the following ways:

K shell (n=1) equals 2, L shell (n=2) equals 8, M shell (n=3) equals 8, and N shell (n=4) equals 1.

Electronic Configuration of Elements

Let us take an example of the electronic configuration of iron, which is mentioned below for a better understanding of the topic

Significance of the electronic configuration of Iron

Iron is a one-of-a-kind element that exists both outside and inside us. The atomic number of Iron is 26.Iron has 8 valence electrons and an electron configuration of 1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d⁶ 4s², which means it has

K shell – 2 electrons,
L shell – 8 electrons,
M shell – 14 electrons, and
N shell – 2 electrons.

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Various properties of Iron can be explained using the electronic configuration. Like, Iron is a silvery white metal that is ductile and malleable under normal conditions. Iron is a medium-activity metal that extracts hydrogen from water solutions of strong acids like HCl and sulfuric acid, resulting in the formation of iron salts. These can be explained by the number of valence shell electrons in Iron.

What are the three rules for writing electronic configurations of elements?

The Aufbau principle states that before occupying an orbital associated with a higher energy level, electrons must entirely fill the atomic orbitals of the previous energy level. In the sequence of increasing orbital energy levels, electrons occupy orbitals of lower energy first.
According to Pauli's exclusion principle, no two electrons may have the same values for all four quantum numbers. As a result, each subshell of an orbital may only hold a maximum of two electrons, both of which must have opposite spins for the spin quantum number to be different.
Hund's rule of maximum multiplicity states that all subshells in an orbital must be occupied singly before any subshell can be twice occupied. Furthermore, all electrons in singly occupied subshells must have the same spin (in order to maximize the overall spin).

ATOMIC NUMBER 1 to 30	ELEMENT	ELECTRONIC CONFIGURATION
1	HYDROGEN (H)	1s¹
2	HELIUM (He)	1s²
3	LITHIUM (Li)	1s² 2s¹
4	BERYLLIUM (Be)	1s² 2s²
5	BORON (B)	1s² 2s¹ 2p¹
6	CARBON (C)	1s² 2s¹ 2p²
7	NITROGEN (N)	1s² 2s¹ 2p³
8	OXYGEN (O)	1s² 2s¹ 2p⁴
9	FLUORINE (F)	1s² 2s¹ 2p⁵
10	NEON (Ne)	1s² 2s¹ 2p⁶
11	SODIUM (Na)	1s² 2s¹ 2p⁶ 3s¹
12	MAGNESIUM (Mg)	1s² 2s¹ 2p⁶ 3s²
13	ALUMINIUM (Al)	1s² 2s¹ 2p⁶ 3s² 3p¹
14	SILICON (Si)	1s² 2s¹ 2p⁶ 3s² 3p²
15	PHOSPHORUS (P)	1s² 2s¹2p⁶ 3s² 3p³
16	SULPHUR (S)	1s² 2s¹ 2p⁶ 3s² 3p⁴
17	CHLORINE (Cl)	1s² 2s¹ 2p⁶ 3s² 3p⁵
18	ARGON (Ar)	1s² 2s¹ 2p⁶ 3s² 3p⁶
19	POTASSIUM (K)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 4s¹
20	CALCIUM (Ca)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 4s²
21	SCANDIUM (Sc)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d¹4s²
22	TITANIUM (Ti)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d² 4s²
23	VANADIUM (V)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d³ 4s²
24	CHROMIUM (Cr)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹
25	MANGANESE (Mn)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d⁵ 4s²
26	IRON (Fe)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d⁶4s²
27	COBALT (Co)	1s² 2s¹ 2p⁶ 3s² 3p⁶3d⁷ 4s²
28	NICKEL (Ni)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d⁸ 4s²
29	COPPER (Cu)	1s² 2s¹2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹
30	ZINC (Zn)	1s² 2s¹ 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s²

The above table contains atomic numbers 1 to 30 elements with symbols and electronic configuration. The electronic configuration of elements can also be written in the form of the electronic configuration of nearest noble gases.

For example- In this table, we can represent the electronic configuration of elements from 21 to 30 in the form of Ar as Argon (atomic number - 18) is the nearest noble gas for the first series elements of the d-block.

What Function Does Electronic Configuration of Elements Play?

The chemical characteristics of elements are largely determined by their electronic arrangement. Despite their small size, electrons are responsible for determining the nature of the elements. They determine the valency, ionization potential, ionization enthalpy, chemical bonding, and practically all other chemical properties of the element. When an element lacks an electron, it is classified as an electron acceptor, and when it has an excess electron, it is classified as an electron giver. As a result, electrical configuration, like the electron, is a deciding factor.

Some Solved Examples

Example 1. Which law indicates the pairing of electrons in the same orbital?

1) Newton’s first law

2) (correct) Hund’s rule

3) Aufbau principle

4) Pauli exclusion principle

Solution

Hund’s rule states that “pairing of electrons in the orbitals belonging to the same subshell (p, d or f) does not take place until each orbital belonging to that subshell has got one electron each. It is singly occupied”.

Hence, the answer is the option (2).

Example 2. The number of electrons that Vanadium (Z = 23) has in p-orbitals is equal to ______

1) (correct) 12

2) 11

3) 10

4) 9

Solution

The electronic configuration of V(Z= 23) is given as

1 s ² 2 s ² 2 p ⁶ 3 s ² 3 p ⁶ 4 s ² 3 d ³

Thus, there are 12 electrons in the p-orbitals of Vanadium.

Hence, the answer is the option (1).

Example 3. Identify the element for which electronic configuration in +3 oxidation state is [ Ar ] 3 d ⁵ : :

1) Ru

2) Mn

3) Co

4) Fe

Solution

As we have learned,

Fe has an electronic configuration of [ Ar ] 4 s ² 3 d⁶

So, Fe 3 + has an electronic configuration [ Ar ] 3 d ⁵.

Hence, the answer is the option (4).

Example 4.

Which of the following configurations represents a noble gas?

A. 1s² 2s² 2p⁶
B. 1s² 2s² 2p⁶ 3s² 3p⁵
C. 1s² 2s² 2p⁶ 3s² 3p⁶
D. 1s² 2s² 2p³

Solution-1s² 2s² 2p⁶ 3s² 3p⁶ totals 18 electrons = Argon (Ar), a noble gas.
Noble gases have full outer shells, making them stable.

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Frequently Asked Questions (FAQs)

1. What is the significance of the noble gas configuration?

Noble gas configurations are particularly stable due to a full valence shell of electrons. Elements tend to gain, lose, or share electrons to achieve a noble gas configuration, which is a driving force behind chemical reactions.

2. Can you provide an example of how Hund's Rule is applied?

In nitrogen (N), which has the electronic configuration 1s² 2s² 2p³, the three electrons in the 2p subshell will occupy separate p orbitals before any pairing occurs. This arrangement minimizes repulsion and stabilizes the atom.

3. What role do orbitals play in electronic configuration?

Orbitals are regions in an atom where electrons are likely to be found. They come in different shapes (s, p, d, f) and energy levels, and they dictate how electrons are arranged in an atom's electronic configuration.

4. How do transition metals differ in their electronic configurations?

Transition metals have partially filled d orbitals. Their electronic configurations can exhibit irregularities due to the stability associated with half-filled and fully filled subshells, leading to variations in expected configurations.

5. What is the periodic trend in the electronic configurations?

As you move across a period in the periodic table, the atomic number increases, leading to an increase in the number of electrons. This results in a gradual filling of orbitals, following the rules of electron configuration.

6. How is atomic mass calculated?

Atomic mass is the weighted average mass of an element's isotopes, measured in atomic mass units (amu). It accounts for both the mass of protons, neutrons, and the relative abundance of each isotope.

7. What is the Aufbau principle?

The Aufbau principle states that electrons fill orbitals from lowest energy to highest energy. This means electrons occupy the innermost shells first before moving to outer shells, following the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on.

8. How do you write the electronic configuration for hydrogen?

The electronic configuration for hydrogen (atomic number 1) is 1s¹. This means it has one electron in the 1s orbital.

9. What is the electronic configuration of carbon?

Carbon (atomic number 6) has the electronic configuration 1s² 2s² 2p². This means it has two electrons in the 1s orbital, two in the 2s orbital, and two in the 2p orbitals.

10. How does Hund's rule apply to 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.

11. What is the Pauli exclusion principle?

The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers. In practice, this means that each orbital can hold a maximum of two electrons with opposite spins.

12. Why is the electronic configuration of elements important?

Electronic configuration is crucial because it determines an element's chemical properties, reactivity, and position in the periodic table. It helps predict how atoms will form bonds and interact with other elements.

13. What is the difference between core electrons and valence electrons?

Core electrons are those in completely filled inner shells, while valence electrons are in the outermost shell. Valence electrons participate in chemical bonding and reactions, while core electrons generally do not.

14. How does the electronic configuration of sodium explain its reactivity?

Sodium's electronic configuration is 1s² 2s² 2p⁶ 3s¹. The single electron in the 3s orbital is easily lost, making sodium highly reactive. This explains why sodium readily forms positive ions in chemical reactions.

15. What is the significance of noble gas configuration?

Noble gas configuration refers to a full outer shell of electrons, which is very stable. Elements tend to gain, lose, or share electrons to achieve this configuration, explaining many chemical reactions and bonding patterns.

16. Why does the 4s orbital fill before the 3d orbital?

The 4s orbital fills before the 3d orbital because it has slightly lower energy. This is due to the complex interplay of electron-electron repulsion and nuclear attraction. However, when electrons are removed, the 4s electrons are lost before 3d electrons.

17. 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.

18. What is an orbital?

An orbital is a region in space around an atom's nucleus where an electron is likely to be found. It's described by a set of quantum numbers and can be visualized as a three-dimensional probability distribution of an electron's location.

19. What is meant by "ground state" electronic configuration?

The ground state electronic configuration refers to the lowest energy arrangement of electrons in an atom. It's the most stable state and the one in which an atom is typically found under normal conditions.

20. How many electrons can the first shell hold?

The first shell (n=1) can hold a maximum of 2 electrons. This is because it only contains one s orbital, and each orbital can hold up to two electrons with opposite spins.

21. What is the "octet rule" and how does it relate to electronic configuration?

The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell, similar to noble gases. This rule is based on the stability of a full outer shell in electronic configuration.

22. What is the relationship between electronic configuration and ionization energy?

Ionization energy is the energy required to remove an electron from an atom. It's directly related to the electronic configuration - generally, the more stable the configuration (like noble gases), the higher the ionization energy.

23. What is the significance of the "18-electron rule" in transition metal complexes?

The 18-electron rule states that stable transition metal complexes often have a total of 18 valence electrons (including those from ligands). This corresponds to a filled valence shell, similar to the octet rule for main group elements.

24. What is meant by "effective nuclear charge" and how does it relate to electronic configuration?

Effective nuclear charge is the net positive charge experienced by an electron in an atom, considering the shielding effect of inner electrons. It increases across a period as more protons are added but the shielding remains relatively constant.

25. What is the relationship between electronic configuration and atomic radius?

Atomic radius generally decreases across a period due to increasing effective nuclear charge, while it increases down a group as new electron shells are added. This trend is directly related to the electronic configuration of elements.

26. What is the "lanthanide contraction" and how is it related to electronic configuration?

The lanthanide contraction is the decrease in atomic and ionic radii across the lanthanide series. It's caused by poor shielding of the nuclear charge by 4f electrons, leading to increased effective nuclear charge and smaller atoms.

27. What is the significance of "spin-orbit coupling" in electronic configuration?

Spin-orbit coupling is the interaction between an electron's spin and its orbital angular momentum. It becomes significant for heavier elements and can affect their spectroscopic properties and magnetic behavior.

28. How does the electronic configuration of chlorine explain its tendency to form negative ions?

Chlorine's electronic configuration is 1s² 2s² 2p⁶ 3s² 3p⁵. It has 7 valence electrons and needs only one more to achieve a stable octet. This makes chlorine prone to gaining an electron and forming a negative ion (Cl⁻).

29. How does the electronic configuration change when an atom becomes an ion?

When an atom becomes an ion, it either loses or gains electrons. For cations (positive ions), electrons are removed from the highest energy level. For anions (negative ions), electrons are added to the lowest unfilled energy level.

30. Why do transition elements often have similar properties despite different electronic configurations?

Transition elements have similar properties because they typically have the same number of electrons in their outermost s orbital (usually 2 or 1), while the d orbitals are being filled. The partially filled d orbitals contribute to their unique properties.

31. How does the electronic configuration of potassium differ from that of argon?

Potassium (1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹) has one more electron than argon (1s² 2s² 2p⁶ 3s² 3p⁶). This extra electron goes into the 4s orbital, making potassium highly reactive compared to the noble gas argon.

32. What is meant by "electron affinity" and how is it related to electronic configuration?

Electron affinity is the energy change when an atom in the gas phase gains an electron. It's generally related to how close an atom is to achieving a stable electronic configuration, often an octet in its outer shell.

33. How does the electronic configuration of elements change across a period in the periodic table?

Across a period, the number of electrons in the outermost shell increases. The main energy level remains the same, but electrons are added to sublevels (s, p, d, f) in order of increasing energy.

34. What is a "half-filled" orbital and why is it relatively stable?

A half-filled orbital contains one electron. It's relatively stable because the electron doesn't have to share the orbital, reducing electron-electron repulsion. Elements with half-filled subshells (like chromium) often have unexpected configurations due to this stability.

35. How does the electronic configuration of scandium differ from calcium?

Scandium (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹) has one more electron than calcium (1s² 2s² 2p⁶ 3s² 3p⁶ 4s²). This extra electron goes into the 3d orbital, making scandium the first transition element.

36. What is the "aufbau diagram" and how is it used?

The aufbau diagram is a visual representation of the order in which orbitals are filled. It shows energy levels and sublevels, helping to determine the electronic configuration of elements by filling orbitals from bottom to top.

37. How does the electronic configuration explain the formation of covalent bonds?

Covalent bonds form when atoms share electrons to achieve stable electronic configurations (usually octets). The shared electrons occupy molecular orbitals formed from the atomic orbitals of the bonding atoms.

38. What is meant by "excited state" electronic configuration?

An excited state electronic configuration occurs when an electron is promoted to a higher energy orbital. This is unstable and temporary, often resulting from the absorption of energy (like light).

39. How does the electronic configuration of copper deviate from the expected pattern?

Copper's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰) deviates from the expected [Ar] 4s² 3d⁹. One 4s electron is moved to the 3d orbital, resulting in a fully filled 3d sublevel, which is more stable.

40. How does the electronic configuration of iron explain its magnetic properties?

Iron's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶) includes unpaired electrons in the 3d orbitals. These unpaired electrons with parallel spins give iron its ferromagnetic properties.

41. How does the electronic configuration of zinc differ from other transition metals?

Zinc (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰) has a fully filled 3d sublevel, unlike most transition metals. This gives zinc properties more similar to main group elements than typical transition metals.

42. How does the electronic configuration of sulfur explain its ability to form expanded octets?

Sulfur's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁴) allows it to use its empty 3d orbitals to accommodate more than 8 electrons in its valence shell, forming expanded octets in some compounds.

43. What is the "diagonal rule" in electronic configuration?

The diagonal rule states that elements diagonally related in the periodic table often have similar properties. This is because they have similar ratios of valence electrons to electron shells, despite different total electron counts.

44. How does the electronic configuration of chromium demonstrate the stability of half-filled orbitals?

Chromium's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵) shows one 4s electron moving to the 3d orbital, resulting in a half-filled 4s and half-filled 3d sublevel, which is more stable than the expected [Ar] 4s² 3d⁴.

45. How does the electronic configuration of nitrogen explain its triple bond in N₂?

Nitrogen's electronic configuration (1s² 2s² 2p³) has three unpaired electrons in its 2p orbitals. These three electrons can form three covalent bonds with another nitrogen atom, resulting in the very stable triple bond in N₂.

46. What is meant by "diamagnetic" and "paramagnetic" in terms of electronic configuration?

Diamagnetic substances have all paired electrons in their electronic configuration, while paramagnetic substances have one or more unpaired electrons. This affects how the substance interacts with magnetic fields.

47. How does the electronic configuration of aluminum explain its ability to form both covalent and ionic bonds?

Aluminum's electronic configuration (1s² 2s² 2p⁶ 3s² 3p¹) has three valence electrons. It can lose all three to form ionic bonds (Al³⁺) or share them to form covalent bonds, giving it versatile bonding capabilities.

48. What is the significance of "electron correlation" in electronic configuration?

Electron correlation refers to the interaction between electrons in an atom. It's not fully accounted for in simple electronic configuration models but is important for understanding the detailed behavior of multi-electron atoms.

49. How does the electronic configuration of neon explain its chemical inertness?

Neon's electronic configuration (1s² 2s² 2p⁶) shows a completely filled outer shell. This stable configuration makes neon chemically inert, as it has no tendency to gain, lose, or share electrons.

50. How does the electronic configuration of manganese explain its variable oxidation states?

Manganese's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵) allows it to lose different numbers of electrons, resulting in various oxidation states (+2 to +7). This is due to the availability of both 4s and 3d electrons.

51. What is meant by "exchange energy" in electronic configuration?

Exchange energy is a quantum mechanical effect that lowers the energy of electrons with parallel spins. It contributes to Hund's rule and explains why some elements, like chromium, have unexpected electronic configurations.

52. How does the electronic configuration of phosphorus explain its ability to form multiple bonds?

Phosphorus's electronic configuration (1s² 2s² 2p⁶ 3s² 3p³) has three unpaired electrons in its 3p orbitals. This allows phosphorus to form single, double, or triple bonds, explaining its diverse chemistry.

53. What is the "spdf" notation in electronic configuration?

The "spdf" notation refers to the four types of orbitals: s (sharp), p (principal), d (diffuse), and f (fundamental). It's used to describe the arrangement of electrons in these orbitals, such as 1s² 2s² 2p⁶.

54. How does the electronic configuration of scandium explain the beginning of the transition series?

Scandium's electronic configuration (1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹) shows the first electron entering the 3d orbital. This marks the beginning of the transition series, where d orbitals are progressively filled.

55. What is the relationship between electronic configuration and electronegativity?

Electronegativity generally increases across a period and decreases down a group. This trend is related to the effective nuclear charge and the distance of valence electrons from the nucleus, both determined by electronic configuration.

56. How does the electronic configuration of boron explain its electron-deficient nature?

Boron's electronic configuration (1s² 2s² 2p¹) has only three valence electrons. This makes boron electron-deficient, as it needs five more electrons to achieve an octet, explaining its tendency to form electron-deficient compounds.