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Chapter 8: Problem 50

Write electron configurations for a. the cations \(\mathrm{Sr}^{2+}, \mathrm{Cs}^{+}, \mathrm{In}^{+},\) and \(\mathrm{Pb}^{2+} .\) b. the anions \(\mathrm{P}^{3-}, \mathrm{S}^{2-},\) and \(\mathrm{Br}^{-}\)

Short Answer

Expert verified
a. Cations: \(\mathrm{Sr}^{2+}\): [Kr] \(\mathrm{Cs}^{+}\): [Xe] \(\mathrm{In}^{+}\): [Kr] 5s^2 4d^10 \(\mathrm{Pb}^{2+}\): [Xe] 4f^14 5d^10 6s^2 b. Anions: \(\mathrm{P}^{3-}\): [Ne] 3s^2 3p^6 \(\mathrm{S}^{2-}\): [Ne] 3s^2 3p^6 \(\mathrm{Br}^{-}\): [Ar] 4s^2 3d^10 4p^6

Step by step solution

01

Determine the atomic number and electron configuration of the neutral atoms

First, we need to find the atomic numbers for each element using the periodic table: Strontium (Sr) - Atomic number: 38 Cesium (Cs) - Atomic number: 55 Indium (In) - Atomic number: 49 Lead (Pb) - Atomic number: 82 Phosphorus (P) - Atomic number: 15 Sulfur (S) - Atomic number: 16 Bromine (Br) - Atomic number: 35 Using the atomic numbers, we can write the electron configurations for the neutral atoms: Sr: [Kr] 5s^2 Cs: [Xe] 6s^1 In: [Kr] 5s^2 4d^10 5p^1 Pb: [Xe] 4f^14 5d^10 6s^2 6p^2 P: [Ne] 3s^2 3p^3 S: [Ne] 3s^2 3p^4 Br: [Ar] 4s^2 3d^10 4p^5
02

Determine electron configurations for the cations

Now, let's write electron configuration for each cation by removing the necessary electrons. a. \(\mathrm{Sr}^{2+}\): [Kr] 5s^2 → [Kr] b. \(\mathrm{Cs}^{+}\): [Xe] 6s^1 → [Xe] c. \(\mathrm{In}^{+}\): [Kr] 5s^2 4d^10 5p^1 → [Kr] 5s^2 4d^10 d. \(\mathrm{Pb}^{2+}\): [Xe] 4f^14 5d^10 6s^2 6p^2 → [Xe] 4f^14 5d^10 6s^2
03

Determine electron configurations for the anions

Finally, let's write electron configuration for each anion by adding the necessary electrons. a. \(\mathrm{P}^{3-}\): [Ne] 3s^2 3p^3 → [Ne] 3s^2 3p^6 b. \(\mathrm{S}^{2-}\): [Ne] 3s^2 3p^4 → [Ne] 3s^2 3p^6 c. \(\mathrm{Br}^{-}\): [Ar] 4s^2 3d^10 4p^5 → [Ar] 4s^2 3d^10 4p^6 So, the electron configurations for the given ions are: a. Cations: \(\mathrm{Sr}^{2+}\): [Kr] \(\mathrm{Cs}^{+}\): [Xe] \(\mathrm{In}^{+}\): [Kr] 5s^2 4d^10 \(\mathrm{Pb}^{2+}\): [Xe] 4f^14 5d^10 6s^2 b. Anions: \(\mathrm{P}^{3-}\): [Ne] 3s^2 3p^6 \(\mathrm{S}^{2-}\): [Ne] 3s^2 3p^6 \(\mathrm{Br}^{-}\): [Ar] 4s^2 3d^10 4p^6

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Most popular questions from this chapter

List all the possible bonds that can occur between the elements P, Cs, O, and H. Predict the type of bond (ionic, covalent, or polar covalent) one would expect to form for each bond.

Without using Fig. 8.3, predict the order of increasing electronegativity in each of the following groups of elements. a. \(\mathrm{C}, \mathrm{N}, \mathrm{O} \quad\) c. $\mathrm{Si}, \mathrm{Ge}, \mathrm{Sn}$ b. \(\mathrm{S}, \mathrm{Se}, \mathrm{Cl} \quad\) d. $\mathrm{TI}, \mathrm{S}, \mathrm{Ge}$

Give the formula of a negative ion that would have the same number of electrons as each of the following positive ions. a. \(\mathrm{Na}^{+} \quad\) c. \(\mathrm{Al}^{3+}\) b. \(\mathrm{Ca}^{2+} \quad\) d. \(\mathrm{Rb}^{+}\)

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. $\mathrm{C}-\mathrm{F}, \mathrm{Si}-\mathrm{F}, \mathrm{Ge}-\mathrm{F} \quad\( c. \)\mathrm{S}-\mathrm{F}, \mathrm{S}-\mathrm{Cl}, \mathrm{S}-\mathrm{Br}$ b. \(\mathrm{P}-\mathrm{Cl}\) or \(\mathrm{S}-\mathrm{Cl} \quad\) d. \(\mathrm{Ti}-\mathrm{Cl}, \mathrm{Si}-\mathrm{Cl}, \mathrm{Ge}-\mathrm{Cl}\)

Write Lewis structures for the following. Show all resonance structures where applicable. a. $\mathrm{NO}_{2}^{-}, \mathrm{NO}_{3}^{-}, \mathrm{N}_{2} \mathrm{O}_{4}\left(\mathrm{N}_{2} \mathrm{O}_{4} \text { exists as } \mathrm{O}_{2} \mathrm{N}-\mathrm{NO}_{2} .\right)$ b. \(\mathrm{OCN}^{-}, \mathrm{SCN}^{-}, \mathrm{N}_{3}^{-}\) (Carbon is the central atom in \(\mathrm{OCN}^{-}\) and \(\mathrm{SCN}^{-} . )\)
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