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Chapter 9: Problem 45

Write Lewis structures for the following molecules: (a) ICl, (b) \(\mathrm{PH}_{3}\), (c) \(\mathrm{P}_{4}\) (each \(\mathrm{P}\) is bonded to three other \(\mathrm{P}\) atoms (d) \(\mathrm{H}_{2} \mathrm{~S}\), (e) \(\mathrm{N}_{2} \mathrm{H}_{4}\), (f) \(\mathrm{HClO}_{3}\), (g) \(\mathrm{COBr}_{2}\) (C is bonded to \(\mathrm{O}\) and \(\mathrm{Br}\) atoms).

Short Answer

Expert verified
The Lewis structures are: ICl: \[I—Cl\] ; PH3: \[H—P—H \ \ \ \ \ H\]; P4: \[P—P \ \ \ \ \ P—P\]; H2S: \[H—S—H\]; N2H4: \[H—N—N—H \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ H \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ H\]; HClO3: \[ O=Cl—O—H \ \ \ \ \ \ \ \ \ O\]; COBr2: \[B—C=O—B\]

Step by step solution

01

ICl

Both iodine (I) and chlorine (Cl) have seven valence electrons cause they belong to Group 7A. After the formation of a Lewis structure, each atom shares one electron with the other, forming a single covalent bond, and both achieve a full octet state. The Lewis structure for ICl is \[ICl: I—Cl\] with three lone pair of electrons on each atom.
02

\(\mathrm{PH}_{3}\)

Phosphorus (P) has five and Hydrogen (H) has one valence electron. Phosphorus forms a single bond with each hydrogen, sharing one electron each, and retains two valence electrons. So, in \(\mathrm{PH}_{3}\), P is in the center surrounded by three hydrogens. The Lewis structure for \(\mathrm{PH}_{3}\) is \[PH_3: H—P—H\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ H\]
03

\(\mathrm{P}_{4}\)

Each phosphorus (P) atom has five valence electrons. In \(\mathrm{P}_{4}\), each phosphorus atom is bonded to three other phosphorus atoms, forming single bonds and retaining two valence electrons. The Lewis structure for \(\mathrm{P}_{4}\) is \[ P_4 : P—P \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ P—P \] with two lone pair of electrons in each P atom
04

\(\mathrm{H}_{2} \mathrm{~S}\)

Sulfur (S) has six and Hydrogen (H) has one valence electron. Sulfur forms a single bond with each Hydrogen atom, sharing one electron each, and retains four as two lone pairs. In \(\mathrm{H}_{2} \mathrm{~S}\), S is in the center surrounded by two hydrogen atoms. The Lewis structure for \(\mathrm{H}_{2} \mathrm{~S}\) is \[H_2S: H—S—H\]
05

\(\mathrm{N}_{2} \mathrm{H}_{4}\)

Each nitrogen (N) has 5 and each hydrogen (H) has one valence electron. Each nitrogen atom forms three bonds - one with another nitrogen and two with individual Hydrogen atoms. The structure resembles a 'H' shape. The Lewis structure for \(\mathrm{N}_{2} \mathrm{H}_{4}\) is \[N_2H_4 : H—N—N—H \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ H \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ H\]
06

\(\mathrm{HClO}_{3}\)

Chlorine (Cl) has seven, each oxygen (O) has six and hydrogen (H) has one valence electron. Chlorine forms 4 bonds – one each with three oxygens and a hydrogen. The bonding oxygen (O) atoms each form a double bond with Cl. The Lewis structure for \(\mathrm{HClO}_{3}\) is \[HClO_3 : O=Cl—O—H \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \O\]
07

\(\mathrm{COBr}_{2}\)

Each Bromine (Br) has seven, Oxygen (O) has six, and carbon (C) has four valence electrons. Carbon forms 4 bonds – a double bond with oxygen and a single bond with each bromine. The Lewis structure for \(\mathrm{COBr}_{2}\) is \[ COBr_2 : Br—C=O—Br\]

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Most popular questions from this chapter

The isolated \(\mathrm{O}^{2-}\) ion is unstable so it is not possible to measure the electron affinity of the \(\mathrm{O}^{-}\) ion directly. Show how you can calculate its value by using the lattice energy of \(\mathrm{MgO}\) and the Born-Haber cycle. [Useful information: \(\mathrm{Mg}(s) \rightarrow \mathrm{Mg}(g) \Delta H^{\circ}=\) \(148 \mathrm{~kJ} / \mathrm{mol} .]\)

A rule for drawing plausible Lewis structures is that the central atom is invariably less electronegative than the surrounding atoms. Explain why this is so. Why does this rule not apply to compounds like \(\mathrm{H}_{2} \mathrm{O}\) and \(\mathrm{NH}_{3} ?\)

For each of the following pairs of elements, state whether the binary compound they form is likely to be ionic or covalent. Write the empirical formula and name of the compound: (a) \(\mathrm{B}\) and \(\mathrm{F},\) (b) \(\mathrm{K}\) and \(\mathrm{Br}\)

For each of the following organic molecules draw a Lewis structure in which the carbon atoms are bonded to each other by single bonds: (a) \(\mathrm{C}_{2} \mathrm{H}_{6}\) (b) \(\mathrm{C}_{4} \mathrm{H}_{10},\) (c) \(\mathrm{C}_{5} \mathrm{H}_{12}\). For (b) and (c), show only structures in which each \(\mathrm{C}\) atom is bonded to no more than two other \(\mathrm{C}\) atoms.

Draw three resonance structures for the molecule \(\mathrm{N}_{2} \mathrm{O}_{3}\) (atomic arrangement is \(\mathrm{ONNO}_{2}\) ). Show formal charges.
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