Atomic Spectra

The ground state electronic configuration of carbon (atomic number 6) is: 1s² 2s² 2p² In the ground state, the 2p electrons occupy different orbitals according to Hund's rule: 1s² 2s² 2pₓ¹ 2pᵧ¹ 2p𝓏⁰
In the excited state, one of the electrons from the 2s orbital is promoted to the 2p orbital to allow bonding (especially in covalent molecules like CH₄), resulting in: 1s² 2s¹ 2pₓ¹ 2pᵧ¹ 2p𝓏¹
This configuration reflects the hybridization readiness of carbon (sp³) seen in many organic compounds.

The correct option is (D):

To determine the correct l values for the subshells, it is essential to understand the concept of azimuthal quantum numbers. The azimuthal quantum number, denoted as l, determines the shape of the electron's orbital and corresponds to the different subshells in an atom. The l values are defined as follows:

• s-subshell: l=0
• p-subshell: l=1
• d-subshell: l=2
• f-subshell: l=3

By matching each subshell with its corresponding l value, we have the order:

• d: 2
• s: 0
• f: 3
• p: 1

Given the options, the sequence 2, 0, 3, 1 matches the l values of the subshells d, s, f, and p, respectively.
Thus, the correct answer is: 2, 0, 3, 1

In the context of visible light, red color is known for having a high wavelength and a low frequency. This relationship between wavelength and frequency is derived from the equation that defines the speed of light:

c = λν

where:

• c is the speed of light in a vacuum, approximately 3.00 x 10⁸ meters per second.
• λ (lambda) represents the wavelength.
• ν (nu) represents the frequency.

According to this equation, wavelength and frequency are inversely proportional. This means as the wavelength increases, the frequency decreases. Red light is on the longer wavelength end of the visible spectrum (approximately 620-750 nm), which results in a lower frequency compared to colors with shorter wavelengths such as blue or violet.

Thus, in visible light, the red color possesses:

• High Wavelength (Longer in comparison to other colors)
• Low Frequency (Lower than colors with shorter wavelengths)

The correct choice that reflects this is: high wavelength and low frequency.

The term "degenerated orbitals" refers to orbitals that have the same energy level. In an atomic subshell, orbitals like , , and are examples of degenerated orbitals as they are energetically equivalent when not in the presence of an external field. Let us evaluate the given options:

: These orbitals are within the same 2p subshell and hence have the same energy in an isolated atom situation, making them degenerate.

: These are different energy levels; as s orbitals in different principal quantum numbers (n) are not degenerate.

: Similar to the 2p case, these are within the same 3p subshell, thus they are degenerate.

Both (1) and (3): Both sets of orbital triplets are degenerate under normal atom conditions without external fields or inter-electronic interactions involving differences in symmetry.

Hence, the correct answer is "Both (1) and (3)" because both the and subshell orbitals are degenerate.

The electronic configuration provided is . To understand which ions this configuration corresponds to, we need to evaluate each element's ground state and how gaining electrons affects their configurations.

1. Identify the ground state configuration of each element:

• Phosphorus ( ) has a ground state configuration of .
• Chlorine ( ) has a ground state configuration of .
• Sulfur ( ) has a ground state configuration of .

2. Account for gained electrons in ionic forms:

• has gained 3 electrons, resulting in , matching the configuration given ( ).
• has gained 1 electron, resulting also in , matching the provided configuration.
• has gained 2 electrons, resulting in , similarly fitting the configuration.

3. Conclusion:

Since all these ions can result in the provided electronic configuration, the correct answer is "All of these".

In the context of electron configurations and quantum chemistry, each atomic orbital can hold a specific number of electrons, and certain rules determine how electrons populate these orbitals. Let's explore the legitimacy of the given options based on their possible electron configurations:

: The 'p' orbital can hold a maximum of 6 electrons. Therefore, this configuration is correct.

: The 's' orbital can hold a maximum of 2 electrons, and having 1 electron is acceptable. Thus, this configuration is also correct.

: The 'f' orbital can hold a maximum of 14 electrons. With 12 electrons, this is a valid configuration.

: The 'd' orbital cannot exist in the second energy level (n=2) as the 'd' orbitals start from the third energy level (n=3). Therefore, this configuration is not possible, making it incorrect.

The correct answer is as it represents an impossible electron configuration according to quantum mechanical principles.

To determine the subshell based on the given quantum numbers, we need to understand what each quantum number represents in quantum chemistry:

Principal Quantum Number (n): This number indicates the energy level of the electron orbital. For n = 2, the electrons are in the second energy level.

Azimuthal Quantum Number (l): This number determines the shape of the orbital and is also referred to as the angular momentum quantum number. The values of l range from 0 to n-1.

For l = 1, the subshell is designated as p. Therefore, with n = 2 and l = 1, the subshell we are looking for is 2p.

Thus, the subshell corresponding to the quantum numbers n = 2 and l = 1 is 2p.