Tag: electronic configuration

Introduction

• The chemical bond formation was first interpreted by the Lewis approach. However, the origin of covalent bonds and the nature of the attraction force between neighbouring atoms in molecules remain unexplained.
• As a solution, German scientists Fritz Wolfgang London and Walter Heinrich Heitler developed the valence bond hypothesis. Schrodinger’s wave equation was used to characterise the formation of a covalent bond between two hydrogen atoms by electron sharing.
• The principles of electronic configuration, electrical structure and atomic orbitals (and their overlap), and the hybridisation of most atomic orbitals are the main topics of this theory. Atomic orbitals that overlap and the electrons that are concentrated in the matching bond area create chemical bonds. 

What is Valence Bond Theory?

The valence bond theory describes the electrical structure of atoms and molecules. The concept explains how electrons are distributed among a molecule’s atomic orbitals. To better understand chemical bonding, the VBT incorporates quantum mechanical principles and explains the electrical makeup of molecules.

According to this theory, all the potential energy levels of an atom in a molecule are occupied, and all bonds are localized, atomic-scale bonds between two atoms that involve sharing an electron by both atoms. Because each atom has only one unpaired electron, their orbitals are weakly connected.

The two atomic orbitals don’t need to be the same. There can be interactions between, say, s and p orbitals. A sigma bond is formed when the orbitals of the two shared electrons intersect. However, Pi bonds are formed when the orbitals cross but remain perpendicular to one another. Since these atomic orbitals will overlap, the possibility of an electron existing in the bond position is highest. Because of the overlapping, electrons are most likely in the bond region.

Postulates of Valence Bond Theory

The main postulates of the Valence Bond Theory are outlined below:

• When two or more atoms come together, their unfilled outer electron shells overlap, creating a covalent bond.
• The orbitals of atoms must be sufficiently near together and properly oriented for overlap to occur.
• The only way for atoms to bind is through sharing electrons, which occurs when their valence orbitals overlap.
• There is an electron pair in overlapping orbitals, and their spins are antiparallel to one another.
• How tightly the orbitals overlap directly affects the quality of the established bond. Stronger covalent bonds are created when more atomic orbitals overlap. 
• When atoms combine to form a molecule, the resulting structure is identical to the original atoms.

Need for Valence Bond Theory

Combining the ideas of Lewis’s pair bonding with the Heitler-London theory, Linus Pauling proposed the valence bond theory in 1928. The concept of the valence bond was developed to account for resonance and orbital hybridisation.

Written by Pauling and published in 1931, “On the Nature of the Chemical Bond” is an in-depth analysis of valence bonds.

Lewis provides the systems to explain how molecules are arranged. However, it omitted to discuss the formation of chemical bonds. Similarly to that, the VSEPR hypothesis defines how essential molecules seem. However, its scope of use was relatively constrained. Additionally, the geometry of a complex atom was not described. To address and resolve these restrictions, scientists had to introduce the notion of valence bonding.

Modern Valence Bond Theory

• As an alternative to the valence bond idea, in which electron pairs are thought to be situated between a molecule’s two distinct atoms, covalent bonds can exist between atoms.
• Molecular orbital theory is the backbone of contemporary valence bond theory, which proposes that electron pairs are distributed around the molecule in sets of molecular orbitals.
• Predictions of magnetic and ionisation properties are simplified by molecular orbital theory, but valence bond theory provides greater nuance.
• The aromatic properties of molecules are attributed to the spin coupling between the pi orbitals, as proposed by contemporary valence bond theory.
• There is not much difference between the Dewar and Kekule arrangements from the perspective of the concept of resonance.
• According to molecular orbital theory, aromaticity is conceptualised as the delocalisation of the pi-electrons. According to molecular orbital theory, the ions participating in a chemical reaction are incorrectly predicted by the function representing the hydrogen molecule. In turn, this causes atoms to link together chemically.

Applications Of Valence Bond Theory

• VBT’s most significant use is in predicting molecular structure. VBT can be used to predict the shape and size of molecules by taking into account the distribution of electrons among atoms.
• The formation of chemical bonds can also be understood using VBT. VBT can explain the type of bond formed between two atoms by understanding the electron-sharing between the atoms. Predicting the strength of a bond and the reactivity of a molecule are two areas where this is particularly helpful.
• The physical and chemical properties of a molecule, like its melting and boiling points, can be predicted with this method.

Summary

The valence bond theory is a chemical bonding concept that describes the nature of the chemical bond between two atoms. Like molecular orbital (MO) theory, it employs quantum mechanical principles to describe bonding. According to the valence bond theory, bonds are formed when unoccupied atomic orbitals overlap. To form a bond, the two atoms use the same unpaired electron to occupy an orbital, resulting in a hybrid orbital. Both sigma and pi bonds are accounted for in the valence bond theory.

Frequently Asked Questions

1. What are the limitations of the Valence Bond Theory?

Ans: Valence Bond Theory is limited in its ability to explain the behaviour of electrons in molecules that contain more than two atoms. It is also limited in its ability to explain the behaviour of electrons in molecules that contain multiple bonds.

2. How does the Molecular Orbital Theory vary from the Valence Bond Theory?

Ans. The Valence Bond Theory explains how atoms in a molecule share electrons to form chemical bonds. Molecular Orbital Theory is based on the idea that electrons in a molecule occupy discrete energy levels, or “orbitals,” around the nucleus.

3. To what extent does geometry factor into the theory of valence bonds? 

Ans. Quite simply said, valence bond theory relies heavily on geometric     considerations. How electron pairs are arranged around a molecule’s core atom determines the structure’s three-dimensional shape. Linear, trigonal planar, tetrahedral, and octahedral arrangements are all possible for the electron pairs. Chemical bond strength and reactivity are both affected by the shape of a molecule, which is determined by the arrangement of electron pairs.

Introduction

The valency of an element is indicated by its maximum capacity of making covalent bonds with any other element or the same element. Carbon has tetra valency since it can form a maximum of four bonds with another C atom or with other atoms like S, H, O, Cl, N, etc. Based on this, C atoms can form several organic compounds like Methane (\(C{H_4}\)), ethane(\({C_2}{H_6}\)), etc. Moreover, carbon can create compounds with both double and triple bonds among its atoms. Carbon chains can sometimes be ring-shaped, branched, or linear.

Define tetravalency of Carbon:

• Tetra signifies “four,” but “Valency” refers to “combining ability.” When carbon has a valency of four and can establish four covalent links with another atom, it is said to be tetravalent.
• The capability of carbon to form covalent linkages with the other carbon molecules tends to be referred to as catenation
• Due to the very small size of carbon, it can undergo catenation.

Reasons behind the tetravalency of Carbon:

The atomic number of carbon is 6 and its electronic configuration is \([He]2{s^2}2{p^2}\). The four electrons in the valence cells mean that the carbon atom can no longer lose or gain four electrons since doing so would require a significant amount of energy. Thus, the four electrons of carbon are shared with other elements. Since its electron has also been shared, there appear to be four shared electrons.

As a consequence, carbon’s valency becomes four, and as an outcome, “Carbon is Tetravalent.” Electrons can neither be generated nor taken up by carbon; it could only transfer them. Its tetravalent character influences the majority of the organic compounds.

Explanation of tetravalency by ground state and excited state configuration of carbon:

tetravalency at Ground state:

The lowest energy level is called the ground state. Most of the electrons in carbon’s ground state are at their lowest available energies. Despite having four electrons, carbon can only create two bonds because its ground state only has two unpaired electrons.

Excited state: 

Carbon is in an excited singlet state whenever the whole energy of the electrons may be lowered by first transferring one electron from the 2s orbital to the 2p orbital. In its excited state, carbon possesses four unpaired electrons, which allows it to form bonds with four other atoms.

For example, one 2s and three 2p electrons combine to form \(s{p^3}\) hybridization. Methane that is \(C{H_4}\) possesses \(s{p^3}\)  hybridization of the C atom. The structure of methane is tetrahedral where C is linked with four H atoms and forms covalent bonds.

The hybridization of carbon in different compounds:

• \(s{p^3}\) hybridization: C atom forms four covalent bonds with four H atoms in\(C{H_4}\). The hybridization of C in \(C{H_4}\) is \(s{p^3}\). \(C{H_4}\) has a tetrahedral structure.
• \(s{p^2}\) hybridization: In ethylene molecules, two C atoms are joined together by one sigma bond and one pi bond. The hybridization of C in \({C_2}{H_4}\) is \(s{p^2}\).
• sp hybridization: Two sigma bonds and two pi bonds are observed in the structure of \({C_2}{H_2}\) (acetylene). Since carbon typically forms two sigma bonds, two of its valence orbitals can combine to generate two orbitals that seem to be equivalent to one another.

Summary 

We may simply conclude that carbon tends to have the closest noble gas configuration since it couples its four valence electrons with several other elements and forms four separate covalent bonds. Tetravalency is the terminology used for this. In a later experiment, carbon proves that it has tetravalency across all hydrocarbons. The following features of carbon can make it the most flexible element in the periodic table: catenation, tetravalency, and isomerism. The production of such a large number of combinations from carbon compounds can be credited to each of these packages. Because they do not form bonds, inorganic composites have a lower number than organic composites.

 

Frequently Asked Questions

1. What other element other than carbon has tetravalency?

Ans: Si also can form four covalent bonds with other elements. That is Si is also tetravalent. But due to the larger size of Si as compared to C, Si doesn’t participate in catenation. Also, Si can’t form various compounds as C atoms can.

2. Is tetravalency possible for all semiconductors?

Ans: There is only one kind of element in it. The most abundant intrinsic semiconductor elements are silicon (Si) and germanium (Ge). Four valence electrons make them tetravalent. At the temperature of absolute zero, a covalent bond connects them to the atom.

3. Why is carbon considered to be a weak conductor of electricity?

Ans: Carbon compounds are found to contain covalent bonds. Covalent molecules do not disintegrate into ions in the aqueous phase;  hence they do not possess any free electrons. Even though there would be no charge transfer, this becomes a poor conductor of electricity.

Introduction

Aufbau is a German word that means “building up.” Like a construction build-up from the ground up. Atoms are also filled with electrons in this manner. An atom has orbitals that are arranged in increasing energy level order. According to the Aufbau principle, electrons are filled in the order of increasing energy of the atomic orbitals. That is from the bottom to the top. This principle aids in the electronic configuration of atoms as well as the placement of electrons in orbitals. In all atoms, the orbital is always the first orbital to be filled with electrons. After filling this orbital, electrons are filled in orbitals further away.

Explain Aufbau Principle

Niels Bohr, a Danish physicist developed the principle. According to this principle, the increasing order of energy levels of atoms causes the filling of electrons in an atom. They are entering a perfect order that corresponds to the energy level of orbitals.  We can predict the electron configurations of atoms or ions by using this rule.

The Madelung rule or rule is also related to this rule. According to this rule, the filling of electrons in an atom occurs as the value of n+l increases. That is, the electrons are filled to a lower-valued orbital. Where n represents the principal quantum number value, and l represents the angular momentum quantum number value. This is known as the Madelung rule or the diagonal rule.

Some features of the Aufbau Principle

1. Electrons are assigned to the subshell with the lowest energetically available energy.
2. An orbital can only hold two electrons.
3. If two or more energetically equivalent orbitals (e.g., p, d, etc.) are available, electrons should be spread out before being paired up (Hund’s rule).

Some Exceptions 

Some elements exhibit exceptional behaviour in terms of the Aufbau principle. They are chromium and copper, respectively. According to the Aufbau principle, the electronic configuration of chromium is \(\left[ {Ar} \right]3{d^4}4{s^2}\). However, chromium’s electronic configuration is \(\left[ {Ar} \right]3{d^5}4{s^1}\). And this is because chromium achieves stability by having a half-filled orbital. Elements require a filled state at all times. A fully-filled orbital is always more stable. Even though a half-filled orbital has partial stability.

Copper’s electronic configuration is \(\left[ {Ar} \right]{\rm{ }}3{d^{10}}4{s^1}\) rather than \(\left[ {Ar} \right]{\rm{ }}3{d^9}4{s^2}\). This is due to the presence of a fully-filled d-orbital configuration, which provides additional stability.

Summary

An electronic configuration is present for all elements to locate electrons in orbitals. As a result, the chemical properties of elements can be explained. When combined with other rules, this can result in a proper electronic configuration. According to Aufbau’s principle, the filling of electrons in an atomic orbital occurs in the order of increasing energy of atomic orbitals. The elements chromium and copper are exceptions to this rule. Because they achieve a half-filled and fully-filled atomic orbital, these elements can be more stable.

Frequently Asked Questions

1. Define Hund’s rule of maximum multiplicity

For an orbital of the same sub-shell, the filling of electrons takes place in a way that all the electrons are singly occupied before pairing occurs. The pairing of electrons takes place only when all the subshells are singly occupied.

2. What do you understand by Pauli’s exclusion principle?

All the quantum number values are distinct for each electron present in an atom. This principle states that no two electrons in an atom can have an equal set of all the quantum number values. And thereby we can easily locate all the electrons in an atom.

3. What is the principal quantum number?

The number that deals with the energy and size of orbitals are a principal quantum number. It will explain how far an electron is from the nucleus. For example, the electronic configuration of Helium is \(1{s^2}\) so the principal quantum number is 1.

Introduction

The electronic configuration describes the distribution of electrons within an atomic subshell. An electron configuration is a summary of the prediction of the position of the electrons surrounding a nucleus. In every neutral atom, the electron number is the same as the proton number. Now we’ll arrange those electrons so that they form a ring around the nucleus, displaying their energy and the orbital type in which they are located. Electrons occupy orbitals in a specific order based on their energy.

What do you understand by Electron Configuration?

• The electronic configuration describes the distribution of electrons within an atomic subshell.
• Atomic electronic configurations follow a standard format in which each atomic subshell containing an electron is listed in ascending order.
• For high atomic numbers, the standard representation of electronic configuration can be quite lengthy. In some cases, an abbreviated/condensed symbol may be used instead of the standard representation.
• The electron configuration of Na, for example, is \(1{s^2}2{s^2}2{p^6}3{s^1}\).

How Subshells are important for Electron Configuration?

• The azimuthal quantum no., represented by the letter “l,” determines the distribution of electrons into subshells.
• The magnitude of the principal quantum no., n, dictates the magnitude of this quantum number. As a result, when n equals 4, four distinct subshells can exist.
• For n = 4, the s, p, d, and f subshells correspond to l=0, 1, 2, 3 quantities.
• Equation 2(2l+1) gives the maximum number of electrons that a subshell can hold.
• The s, p, d, and f subshells can hold a maximum of 2, 6, 10, and 14 electrons, respectively.

Atomic Electronic Configuration Representation

This section provides examples of a few elements’ electronic configurations.

• The electron configuration of hydrogen has an atomic number of one. As a result, an H atom has one electron, which will be assigned to the subshell of the first shell/s orbit. \(1{s^1}\) is the electronic configuration of H.
• The electron configuration of chlorine

Cl has the atomic number 17. As a result, its 17 electrons are distributed as follows:

The K has two electrons.

The L has 8 electrons and the M has 7 electrons.

The electron configuration of Cl is depicted below. It is written as \(1{s^2}2{s^2}2{p^6}3{s^2}3{p^5}\).

Filling Atomic Orbitals

The following concepts govern how electrons are occupied in atomic orbitals.

Aufbau Principle

“The energy of an atomic orbital is calculated by adding the principal and azimuthal quantum numbers, and according to the Aufbau principle, electrons begin in relatively low energy orbitals and progress to higher energy orbitals.”

Pauli Exclusion Principle

“Only electron pairs with opposite spins can be carried in an atomic orbital, and no two electrons in the same atom have the same values for all four quantum numbers. If two electrons have the same principle, azimuthal, and magnetic numbers, they should have opposing spins.”

Hund’s Law

“Before a second electron is placed in an orbital, each orbital in a specific subshell is said to be entirely filled by electrons.”

Summary

It can be concluded that Electron configuration is the depiction of electron distribution inside an element’s atomic shells. Because the electrons are mathematically positioned in these subshells, the configuration aids in determining their position. The periodic table categorises elements based on their electron configurations. These make up the s, p, d, and f-block elements. The maximum number of electrons that can fit in a shell is determined by the principal quantum number (n). The azimuthal quantum number, represented by the letter “l,” governs the distribution of electrons into subshells.

Frequently Asked Questions

1. Why are specific electron configurations required for elements?
Ans. Electron configurations can shed light on an atom’s chemical behaviour by identifying its valence electrons. It also aids in the organisation of elements into different blocks such as s, p, d, and f blocks.

2. Describe the significance of electron configuration.
Ans. The significance is as follows:

They aid in determining the reactivity state of an atom.

It aids in the identification of both chemical and physical properties.

It foretells an atom’s magnetic properties.

3. For n=3, which subshells are present?
Ans. Each orbital can hold a maximum of two electrons, and there are four subshells present- s, p, d, and f for n=3. The maximum number of orbitals corresponding to the s, p, d, and f subshells is 1,3,5, and 7.

Introduction

An atom’s energy changes due to electron affinity. A neutral atom gains energy and a negative charge when electrons are added to its outer shell. To stabilise its octet, an element gains electrons. When an element accepts or loses an electron, energy is released. When an element accepts an electron to form a compound, it releases energy, which is referred to as an exothermic reaction. The energy is released in an exothermic reaction in order to attract the electron by a nucleus from another element. When an element loses an electron, it absorbs energy, a process known as endothermic. An atom gains energy when it loses electrons.

What do you mean by Electron Affinity?

When atoms accept electrons, they emit energy, which is referred to as an exothermic reaction. Atoms that lose an electron in a chemical reaction, on the other hand, absorb energy and are known as endothermic reactions. The ability to accept an electron is referred to as electron affinity. When a neutral gaseous atom accepts an electron, it gains a negative ion charge. The first electron affinity is always negative, while the second is always positive. It is difficult to measure the electron affinity of an atom. It is determined by the energy released by ionic compounds. The electron affinity is also measured by an atom’s tendency to act as an oxidising or reducing agent. It is measured in kilojoules/moles. Electron affinity is symbolised by EA.

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Factors Influencing Electron Affinity

The atomic size of the element, the nuclear charge on the molecules, and the electronic configuration of atoms are all factors that influence a molecule’s electron affinity.

1. Atomic size: Atoms with smaller sizes have greater electron affinity than atoms with larger sizes. The nucleus of smaller atoms is more attractive to electrons than the nucleus of larger atoms. As the atom’s size increases, the outer shell becomes further away from the nucleus, and the attraction for electrons in the outer shell decreases. 
2. Nuclear Charge: The nuclear charge influences electron affinity as well. As the charge on an atom increases, so does the attraction in electrons, and thus the electron affinity. When a molecule is already charged, electron repulsion increases, and the pull from the nucleus increases, resulting in increased electron affinity in charged ions.
3. Shielding Effect: As the screening effect on an atom’s inner shell is reduced, the electron affinity increases.
4. Electronic Configuration: The electronic configuration also affects electron affinity. Because elements with full octets have zero tendencies to accept electrons, electron affinity in inert gases is zero. The electronic configuration is crucial in electron affinity. Metals have a lower affinity for electrons than non-metals due to their electronic configuration.

Summary

The ability to accept electrons in gaseous form and form an anion is referred to as electron affinity. The process of accepting electrons generates energy, which is why it is referred to as an exothermic process. When we move from group to group, the electron affinity decreases and increases when we move from period to period. It is denoted by the symbol EA and measured in Kilojoules per Mole (KJ/Mol). Because of electron-electron repulsion, the first electron affinity is always less than the second electron affinity. The atomic size, electronic configuration, screening effect, and nuclear charge of elements all influence electron affinity.

Frequently Asked Questions

1. Why do noble gases have no electron affinity?

Ans. Noble gases have zero electron affinity because their octet is complete, and they do not have an affinity for electrons. As a result, noble gases have no electron affinity.

2. Why does group 17 have such a strong electron affinity?

Ans. Because the halogens are small and have more electrons in the outer shell, the elements of the halogens group have a high electron affinity. A halogen would rather accept an electron than lose seven electrons to complete its octet.

3. Why does fluorine have a lower electron affinity than chlorine?

Ans. Because the atomic size of fluorine molecules is smaller than that of chlorine molecules, the outer shell of fluorine is already filled with electrons, and the nucleus is much closer to the outer shell, the electron repulsion is greater than the force of attraction of the nucleus when an electron is placed in the outer shell of fluorine molecules compared to chlorine molecules.