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14th Edition

Burdge

ISBN: 9781259327933

**(a)**

Interpretation Introduction

**Interpretation:**

The values of the quantum numbers associated with the given orbitals should be identified using the concept of quantum numbers.

**Concept Introduction:**

Each electron in an atom is described by four different quantum numbers. The first three (n, l, m_{l}) specify the particular orbital of interest, and the fourth (m_{s}) specifies how many electrons can occupy that orbital.

Principal Quantum Number (n)

The principal quantum number (n) assigns the **size of the orbital** and specifies the **energy** of an electron. If the value of n is larger, then the average distance of an electron in the orbital from the nucleus will be greater. Therefore the size of the orbital is large. The principal quantum numbers have the integral values of 1, 2, 3 and so forth and it corresponds to the quantum number in Bohr’s model of the hydrogen atom. If all orbitals have the same value of ‘n’, they are said to be in the same **shell** (**level**). The total number of orbitals for a given n value is n^{2}. As the value of ‘n’ increases, the energy of the electron also increases.

Angular Momentum Quantum Number (l)

The angular momentum quantum number (l) explains the **shape of the atomic orbital**. The values of l are integers which depend on the value of the principal quantum number, n. For a given value of n, the possible values of l range are from 0 to n − 1. If n = 1, there is only one possible value of l (l=0). If n = 2, there are two values of l: 0 and 1. If n = 3, there are three values of l: 0, 1, and 2. The value of l is selected by the letters s, p, d, and f. If l = 0, we have an s orbital; if l = 1, we have a p orbital; if l = 2, we have a d orbital and finally if l = 3, we have a f orbital. A collection of orbitals with the same value of n is called a shell. One or more orbitals with the same n and l values are referred to a **subshell (sublevel)**. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).

Magnetic Quantum Number (m_{l})

The magnetic quantum number (m_{l}) explains the **orientation of the orbital in space**. The value of m_{l} depends on the value of l in a subshell. This number divides the subshell into individual **orbitals** which hold the electrons. For a certain value of l, there are (2l + 1) integral values of m_{l} which is explained as follows:

m_{l }= ‒ l, ..., 0, ..., +l

If l = 0, there is only one possible value of m_{l}: 0.

If l = 1, then there are three values of m_{l}: −1, 0, and +1.

If l = 2, there are five values of m_{l}, namely, −2, −1, 0, +1, and +2.

If l = 3, there are seven values of m_{l}, namely, −3, −2, −1, 0, +1, +2, and +3, and so on.

The number of m_{l} values indicates the number of orbitals in a subshell with a particular l value. Therefore, each m_{l} value refers to a different orbital.

Electron Spin Quantum Number (m_{s})

It specifies the **orientation of the spin axis** of an electron. An electron can spin in only one of two directions. There are two possible ways to represent m_{s }values. They are +½ and ‒½. One electron spins in the clockwise direction. Another electron spins in the anticlockwise direction. But, no two electrons should have the same spin quantum number.

**To find:** Get the values of the quantum numbers (n, l, m_{l}, m_{s}) associated with the given orbital (a) 2p

Get the values of the quantum numbers ‘n’, ‘l’ in (a)

(b)

Interpretation Introduction

**Interpretation:**

The values of the quantum numbers associated with the given orbitals should be identified using the concept of quantum numbers.

**Concept Introduction:**

Each electron in an atom is described by four different quantum numbers. The first three (n, l, m_{l}) specify the particular orbital of interest, and the fourth (m_{s}) specifies how many electrons can occupy that orbital.

Principal Quantum Number (n)

The principal quantum number (n) assigns the **size of the orbital** and specifies the **energy** of an electron. If the value of n is larger, then the average distance of an electron in the orbital from the nucleus will be greater. Therefore the size of the orbital is large. The principal quantum numbers have the integral values of 1, 2, 3 and so forth and it corresponds to the quantum number in Bohr’s model of the hydrogen atom. If all orbitals have the same value of ‘n’, they are said to be in the same **shell** (**level**). The total number of orbitals for a given n value is n^{2}. As the value of ‘n’ increases, the energy of the electron also increases.

Angular Momentum Quantum Number (l)

The angular momentum quantum number (l) explains the **shape of the atomic orbital**. The values of l are integers which depend on the value of the principal quantum number, n. For a given value of n, the possible values of l range are from 0 to n − 1. If n = 1, there is only one possible value of l (l=0). If n = 2, there are two values of l: 0 and 1. If n = 3, there are three values of l: 0, 1, and 2. The value of l is selected by the letters s, p, d, and f. If l = 0, we have an s orbital; if l = 1, we have a p orbital; if l = 2, we have a d orbital and finally if l = 3, we have a f orbital. A collection of orbitals with the same value of n is called a shell. One or more orbitals with the same n and l values are referred to a **subshell (sublevel)**. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).

Magnetic Quantum Number (m_{l})

The magnetic quantum number (m_{l}) explains the **orientation of the orbital in space**. The value of m_{l} depends on the value of l in a subshell. This number divides the subshell into individual **orbitals** which hold the electrons. For a certain value of l, there are (2l + 1) integral values of m_{l} which is explained as follows:

m_{l }= ‒ l, ..., 0, ..., +l

If l = 0, there is only one possible value of m_{l}: 0.

If l = 1, then there are three values of m_{l}: −1, 0, and +1.

If l = 2, there are five values of m_{l}, namely, −2, −1, 0, +1, and +2.

If l = 3, there are seven values of m_{l}, namely, −3, −2, −1, 0, +1, +2, and +3, and so on.

The number of m_{l} values indicates the number of orbitals in a subshell with a particular l value. Therefore, each m_{l} value refers to a different orbital.

Electron Spin Quantum Number (m_{s})

It specifies the **orientation of the spin axis** of an electron. An electron can spin in only one of two directions. There are two possible ways to represent m_{s }values. They are +½ and ‒½. One electron spins in the clockwise direction. Another electron spins in the anticlockwise direction. But, no two electrons should have the same spin quantum number.

**To find:** Get the values of the quantum numbers (n, l, m_{l}, m_{s}) associated with the given orbital (b) 3s

Get the values of the quantum numbers ‘n’, ‘l’ in (b)

(c)

Interpretation Introduction

**Interpretation:**

The values of the quantum numbers associated with the given orbitals should be identified using the concept of quantum numbers.

**Concept Introduction:**

Each electron in an atom is described by four different quantum numbers. The first three (n, l, m_{l}) specify the particular orbital of interest, and the fourth (m_{s}) specifies how many electrons can occupy that orbital.

Principal Quantum Number (n)

The principal quantum number (n) assigns the **size of the orbital** and specifies the **energy** of an electron. If the value of n is larger, then the average distance of an electron in the orbital from the nucleus will be greater. Therefore the size of the orbital is large. The principal quantum numbers have the integral values of 1, 2, 3 and so forth and it corresponds to the quantum number in Bohr’s model of the hydrogen atom. If all orbitals have the same value of ‘n’, they are said to be in the same **shell** (**level**). The total number of orbitals for a given n value is n^{2}. As the value of ‘n’ increases, the energy of the electron also increases.

Angular Momentum Quantum Number (l)

The angular momentum quantum number (l) explains the **shape of the atomic orbital**. The values of l are integers which depend on the value of the principal quantum number, n. For a given value of n, the possible values of l range are from 0 to n − 1. If n = 1, there is only one possible value of l (l=0). If n = 2, there are two values of l: 0 and 1. If n = 3, there are three values of l: 0, 1, and 2. The value of l is selected by the letters s, p, d, and f. If l = 0, we have an s orbital; if l = 1, we have a p orbital; if l = 2, we have a d orbital and finally if l = 3, we have a f orbital. A collection of orbitals with the same value of n is called a shell. One or more orbitals with the same n and l values are referred to a **subshell (sublevel)**. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).

Magnetic Quantum Number (m_{l})

The magnetic quantum number (m_{l}) explains the **orientation of the orbital in space**. The value of m_{l} depends on the value of l in a subshell. This number divides the subshell into individual **orbitals** which hold the electrons. For a certain value of l, there are (2l + 1) integral values of m_{l} which is explained as follows:

m_{l }= ‒ l, ..., 0, ..., +l

If l = 0, there is only one possible value of m_{l}: 0.

If l = 1, then there are three values of m_{l}: −1, 0, and +1.

If l = 2, there are five values of m_{l}, namely, −2, −1, 0, +1, and +2.

If l = 3, there are seven values of m_{l}, namely, −3, −2, −1, 0, +1, +2, and +3, and so on.

The number of m_{l} values indicates the number of orbitals in a subshell with a particular l value. Therefore, each m_{l} value refers to a different orbital.

Electron Spin Quantum Number (m_{s})

It specifies the **orientation of the spin axis** of an electron. An electron can spin in only one of two directions. There are two possible ways to represent m_{s }values. They are +½ and ‒½. One electron spins in the clockwise direction. Another electron spins in the anticlockwise direction. But, no two electrons should have the same spin quantum number.

**To find:** Get the values of the quantum numbers (n, l, m_{l}, m_{s}) associated with the given orbital (c) 5d

Get the values of the quantum numbers ‘n’, ‘l’ in (c)