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Chapter 8: Problem 11

Consider the element silicon, Si. (a) Write its electron configuration. (b) How many valence electrons does a silicon atom have? (c) Which subshells hold the valence electrons?

Short Answer

Expert verified
The electron configuration of Silicon (Si) is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^2}\). Silicon has 4 valence electrons, which are found in the 3s and 3p subshells.

Step by step solution

01

Writing the electron configuration of Silicon (Si)

Silicon is the element with atomic number 14, which means it has 14 electrons. To write its electron configuration, we can follow the Aufbau principle, which suggests filling the orbitals in the order of increasing energy levels. The order is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f, 5d, 6s, 5f, 6d, 7s, 6f, 7d, 7f, 8s, ... Filling in each orbital with the given 14 electrons, we get: 1s虏, 2s虏, 2p鈦�, 3s虏, 3p虏 So, the electron configuration for Silicon is \(\mathrm{1s^2 2s^2 2p^6 3s^2 3p^2}\).
02

Determining the number of valence electrons

Valence electrons are the electrons found in the outermost energy level of an atom. In the case of Silicon, the outermost energy level is n=3. We have 2 electrons in the 3s subshell and 2 electrons in the 3p subshell. Therefore, Silicon has a total of 2 + 2 = 4 valence electrons.
03

Identifying the subshells holding the valence electrons

As mentioned in Step 2, the valence electrons are in the outermost energy level, which is n=3 for Silicon. These valence electrons are found in two subshells: the 3s subshell and the 3p subshell. So, the subshells that hold the valence electrons of Silicon are 3s and 3p.

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Most popular questions from this chapter

What is the Lewis symbol for each of the following atoms or ions? (a) \(\mathrm{K}\), (b) \(\mathrm{As}\), (c) \(\mathrm{Sn}^{2+}\), (d) \(\mathrm{N}^{3-}\).

Illustrated are four ions - \(A, B, X\), and \(Y-\) showing their relative ionic radii. The ions shown in red carry positive charges: a \(2+\) charge for A and a \(1+\) charge for B. Ions shown in blue carry negative charges: a \(1-\) charge for \(X\) and a \(2-\) charge for \(Y\). (a) Which combinations of these ions produce ionic compounds where there is a \(1: 1\) ratio of cations and anions? (b) Among the combinations in part (a), which leads to the ionic compound having the largest lattice energy? [Section 8.2]

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Cd}^{2+}\), (b) \(\mathrm{P}^{3-}\), (c) \(\mathrm{Zr}^{4+}\), (d) \(\mathrm{Ru}^{3+}\), (e) \(\mathrm{As}^{3-}\), (f) \(\mathrm{Ag}^{+}\).

Which of the following trends in lattice energy is due to differences in ionic radii? (a) \(\mathrm{NaCl}>\mathrm{RbBr}>\mathrm{CsBr}\), (b) \(\mathrm{BaO}>\mathrm{KF}\), (c) \(\mathrm{SrO}>\mathrm{SrCl}_{2}\).

One scale for electronegativity is based on the concept that the electronegativity of any atom is proportional to the ionization energy of the atom minus its electron affinity: electronegativity \(=k(I-E A)\), where \(k\) is a proportionality constant. (a) How does this definition explain why the electronegativity of \(\mathrm{F}\) is greater than that of \(\mathrm{Cl}\) even though \(\mathrm{Cl}\) has the greater electron affinity? (b) Why are both ionization energy and electron affinity relevant to the notion of electronegativity? (c) By using data in Chapter 7 , determine the value of \(k\) that would lead to an electronegativity of \(4.0\) for \(F\) under this definition. (d) Use your result from part (c) to determine the electronegativities of \(\mathrm{Cl}\) and \(\mathrm{O}\) using this scale. (e) Another scale for electronegativity defines electronegativity as the average of an atom's first ionization energy and its electron affinity. Using this scale, calculate the electronegativities for the halogens, and scale them so fluorine has an electronegativity of 4.0. On this scale, what is Br's electronegativity?
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