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Chapter 8: Problem 14

What is the Lewis symbol for each of the following atoms or ions? $(\mathbf{a}) \mathrm{Be},(\mathbf{b}) \mathrm{Rb},(\mathbf{c}) \mathrm{I}^{-},(\mathbf{d}) \mathrm{Se}^{2-} .$

Short Answer

Expert verified
The Lewis symbols for the given atoms and ions are: (a) Be (Beryllium): Be •• (b) Rb (Rubidium): Rb • (c) I⁻ (Iodide Ion): \[ \mathrm{[I \cdot \cdot \cdot \cdot \cdot \cdot \cdot ]^{-}} \] (d) Se²⁻ (Selenide Ion): \[ \mathrm{[Se \cdot \cdot \cdot \cdot \cdot \cdot \cdot \cdot ]^{2-}} \]

Step by step solution

01

Determine the Valence Electrons for Each Atom or Ion

For each atom or ion, we will identify its atomic number (which determines its place in the periodic table), which will guide us in determining the number of valence electrons. In case of ions, we will also account for any gained or lost electrons.
02

Create the Lewis Symbols for Each Atom or Ion

Using the chemical symbol for the element and the number of valence electrons, we will place dots around the symbol to represent the valence electrons. Keep in mind to distribute the electrons according to the octet rule, and add brackets and charge, if it is an ion. Now let's find the Lewis symbols for each given atom or ion:
03

(a) Be (Beryllium)

Atomic Number = 4 Valence Electrons = 2 Lewis Symbol for Be: Be ••
04

(b) Rb (Rubidium)

Atomic Number = 37 Valence Electrons = 1 Lewis Symbol for Rb: Rb •
05

(c) I- (Iodide Ion)

Atomic Number = 53 Valence Electrons = 7, adding 1 electron for the negative charge Lewis Symbol for I-: \[ \mathrm{[I \cdot \cdot \cdot \cdot \cdot \cdot \cdot ]^{-}} \]
06

(d) Se2- (Selenide Ion)

Atomic Number = 34 Valence Electrons = 6, adding 2 electrons for the double negative charge (-2) Lewis Symbol for Se2-: \[ \mathrm{[Se \cdot \cdot \cdot \cdot \cdot \cdot \cdot \cdot ]^{2-}} \]

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Most popular questions from this chapter

Write Lewis structures for the following: (a) \(\mathrm{H}_{2} \mathrm{CO}\) (both \(\mathrm{H}\) atoms are bonded to $\mathrm{C} \mathrm{)},(\mathbf{b}) \mathrm{H}_{2} \mathrm{O}_{2},(\mathbf{c}) \mathrm{C}_{2} \mathrm{~F}_{6}($ contains a \(\mathrm{C}-\mathrm{C}\) bond $),(\mathbf{d}) \mathrm{AsO}_{3}^{3-},(\mathbf{e}) \mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}$ is bonded to \(\mathrm{O})\) (f) \(\mathrm{NH}_{2} \mathrm{Cl}\).

A new compound is made that has a \(\mathrm{C}-\mathrm{O}\) bond length of $120 \mathrm{pm}$. Is this bond likely to be a single, double, or triple C-O bond?

Although \(\mathrm{I}_{3}^{-}\) is a known ion, \(\mathrm{F}_{3}^{-}\) is not. \((\mathbf{a})\) Draw the Lewis structure for \(I_{3}^{-}\) (it is linear, not a triangle). (b) One of your classmates says that \(\mathrm{F}_{3}^{-}\) does not exist because \(\mathrm{F}\) is too electronegative to make bonds with another atom. Give an example that proves your classmate is wrong. (c) Another classmate says \(\mathrm{F}_{3}^{-}\) does not exist because it would violate the octet rule. Is this classmate possibly correct? (d) Yet another classmate says \(\mathrm{F}_{3}^{-}\) does not exist because \(\mathrm{F}\) is too small to make bonds to more than one atom. Is this classmate possibly correct?

Under special conditions, sulfur reacts with anhydrous liquid ammonia to form a binary compound of sulfur and nitrogen. The compound is found to consist of \(69.6 \% \mathrm{~S}\) and \(30.4 \% \mathrm{~N}\). Measurements of its molecular mass yield a value of \(184.3 \mathrm{~g} / \mathrm{mol}\). The compound occasionally detonates on being struck or when heated rapidly. The sulfur and nitrogen atoms of the molecule are joined in a ring. All the bonds in the ring are of the same length. (a) Calculate the empirical and molecular formulas for the substance. (b) Write Lewis structures for the molecule, based on the information you are given. (Hint: You should find a relatively small number of dominant Lewis structures.) (c) Predict the bond distances between the atoms in the ring. (Note: The \(\mathrm{S}-\mathrm{S}\) distance in the \(\mathrm{S}_{8}\) ring is \(205 \mathrm{pm} .\) ) \((\mathbf{d})\) The enthalpy of formation of the compound is estimated to be $480 \mathrm{~kJ} / \mathrm{mol}^{-1} . \Delta H_{f}^{\circ}\( of \)\mathrm{S}(g)\( is \)222.8 \mathrm{~kJ} / \mathrm{mol}$. Estimate the average bond enthalpy in the compound.

Draw the Lewis structures for each of the following molecules or ions. Identify instances where the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state how many electrons surround these atoms: $(\mathbf{a}) \mathrm{PF}_{6}^{-},(\mathbf{b}) \mathrm{BeCl}_{2},(\mathbf{c}) \mathrm{NH}_{3},(\mathbf{d}) \mathrm{XeF}_{2} \mathrm{O}\( (the Xe is the central atom), (e) \)\mathrm{SO}_{4}^{2-}$.
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