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Chapter 6: Problem 71

What is the maximum number of electrons that can occupy each of the following subshells? (a) \(3 s,(\) b) \(2 p,(\mathbf{c}) 4 d\), (d) \(5 s\).

Short Answer

Expert verified
The maximum number of electrons that can occupy each subshell are: (a) 3s: 2 electrons, (b) 2p: 6 electrons, (c) 4d: 10 electrons, and (d) 5s: 2 electrons.

Step by step solution

01

Identify the quantum numbers for each subshell

To identify the quantum numbers, we need to know the principal quantum number (n) and the letters corresponding to the orbital angular momentum quantum number (l). Here, s, p, d, and f represent l = 0, 1, 2, and 3, respectively. (a) 3s: n = 3 and l = 0 (b) 2p: n = 2 and l = 1 (c) 4d: n = 4 and l = 2 (d) 5s: n = 5 and l = 0
02

Calculate the number of electrons for each subshell

Now, use the formula: Number of electrons = 2(2l + 1) to find the maximum number of electrons that can occupy each subshell. (a) 3s: Number of electrons = 2(2(0) + 1) = 2 (b) 2p: Number of electrons = 2(2(1) + 1) = 6 (c) 4d: Number of electrons = 2(2(2) + 1) = 10 (d) 5s: Number of electrons = 2(2(0) + 1) = 2
03

State the maximum number of electrons for each subshell

(a) The maximum number of electrons that can occupy the 3s subshell is 2. (b) The maximum number of electrons that can occupy the 2p subshell is 6. (c) The maximum number of electrons that can occupy the 4d subshell is 10. (d) The maximum number of electrons that can occupy the 5s subshell is 2.

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Most popular questions from this chapter

Consider a transition in which the electron of a hydrogen atom is excited from \(n=1\) to \(n=\infty\). (a) What is the end result of this transition? (b) What is the wavelength of light that must be absorbed to accomplish this process? (c) What will occur if light with a shorter wavelength than that in part (b) is used to excite the hydrogen atom? (d) How are the results of parts \((\mathrm{b})\) and \((\mathrm{c})\) related to the plot shown in Exercise \(6.88 ?\)

(a) A green laser pointer emits light with a wavelength of \(532 \mathrm{nm}\). What is the frequency of this light? (b) What is the energy of one of these photons? (c) The laser pointer emits light because electrons in the material are excited (by a battery) from their ground state to an upper excited state. When the electrons return to the ground state, they lose the excess energy in the form of \(532-\mathrm{nm}\) photons. What is the energy gap between the ground state and excited state in the laser material?

Sketch the shape and orientation of the following types of orbitals: \((\mathbf{a}) s,(\mathbf{b}) p_{z},(\mathbf{c}) d_{x y}\).

Write the condensed electron configurations for the following atoms and indicate how many unpaired electrons each has: $(\mathbf{a}) \mathrm{Mg},(\mathbf{b}) \mathrm{Ge},(\mathbf{c}) \mathrm{Br},(\mathbf{d}) \mathrm{V},(\mathbf{e}) \mathrm{Y},(\mathbf{f}) \mathrm{Lu} .$

(a) Account for formation of the following series of oxides in terms of the electron configurations of the elements and the discussion of ionic compounds in Section 2.7: $\mathrm{K}_{2} \mathrm{O}, \mathrm{CaO}, \mathrm{Sc}_{2} \mathrm{O}_{3}, \mathrm{TiO}_{2}, \mathrm{~V}_{2} \mathrm{O}_{5}, \mathrm{CrO}_{3} .(\mathbf{b})$ Name these oxides. (c) Consider the metal oxides whose enthalpies of formation (in $\mathrm{kJ} \mathrm{mol}^{-1}$ ) are listed here. Calculate the enthalpy changes in the following general reaction for each case: $$\mathrm{M}_{n} \mathrm{O}_{m}(s)+\mathrm{H}_{2}(g) \longrightarrow n \mathrm{M}(s)+m \mathrm{H}_{2} \mathrm{O}(g)$$ (You will need to write the balanced equation for each case and then compute \(\left.\Delta H^{\circ} .\right)\) (d) Based on the data given, estimate a value of \(\Delta H_{f}^{\circ}\) for \(\mathrm{Sc}_{2} \mathrm{O}_{3}(s)\)
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