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Chapter 1: Problem 7

Draw Lewis structures for these ions, and show which atom in each bears the formal charge. (a) \(\mathrm{CH}_{3} \mathrm{NH}_{3}^{+}\) (b) \(\mathrm{CO}_{3}{ }^{2-}\) (c) \(\mathrm{OH}^{-}\)

Short Answer

Expert verified
(a) \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\) (b) \(\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{OH}^{-}\) Answer: (a) In \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\), the Nitrogen atom bears the formal charge. (b) In \(\mathrm{CO}_{3}^{2-}\), one of the Oxygen atoms bears the formal charge. (c) In \(\mathrm{OH}^{-}\), the Oxygen atom bears the formal charge.

Step by step solution

01

Count total electrons

Count the total number of valence electrons in the molecule, adding or subtracting one electron per positive or negative charge. Here, we have 3 Hydrogen atoms (3 x 1 e-), 1 Carbon atom (4 e-), 1 Nitrogen atom (5 e-) and a +1 charge (-1 e-). Thus, we have a total of 3 + 4 + 5 - 1 = 11 electrons.
02

Arrange atoms and place electrons

Start with the least electronegative atom (Carbon) in the center and arrange other atoms around it, resulting in the structure H - C - N - H, with a hydrogen atom attached to both C and N. Distribute the electrons around the atoms in pairs, starting with the atoms directly bonded together.
03

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, making sure not to exceed the octet rule. Calculate the formal charge of each atom in the ion using this formula: Formal Charge = Valence electrons - Non-bonding electrons - 1/2 Bonding electrons The \(\mathrm{CH}_{3}\mathrm{NH}_{3}^{+}\) Lewis structure becomes: H H H | | | C―N | H Since Nitrogen has a lower formal charge than the other atoms, it will bear the formal charge in the ion. (b) For \(\mathrm{CO}_{3}{ }^{2-}\)
04

Count total electrons

Count the total number of valence electrons in the ion, adding or subtracting one electron per positive or negative charge. Here, we have 1 Carbon atom (4 e-), 3 Oxygen atoms (3 x 6 e-) and a -2 charge (+2 e-). Thus, we have a total of 4 + 18 + 2 = 24 electrons.
05

Arrange atoms and place electrons

Start with the least electronegative atom (Carbon) in the center and arrange other atoms (Oxygen) around it. Connect Carbon to each Oxygen with a single bond, and distribute the electrons around the atoms in pairs, starting with the atoms directly bonded together.
06

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, deciding on single, double, or triple bonds as necessary to satisfy the octet rule. Calculate the formal charge of each atom in the ion. The \(\mathrm{CO}_{3}{ }^{2-}\) Lewis structure becomes: O // C \ O- \ O Each of the three Oxygen atoms has a formal charge of -1. One Oxygen atom bears a formal charge in the ion. (c) For \(\mathrm{OH}^{-}\)
07

Count total electrons

Count the total number of valence electrons in the ion, adding or subtracting one electron per positive or negative charge. Here, we have 1 Hydrogen atom (1 e-), 1 Oxygen atom (6 e-), and a -1 charge (+1 e-). Thus, we have a total of 1 + 6 + 1 = 8 electrons.
08

Arrange atoms and place electrons

Start with the least electronegative atom (Oxygen) and connect with bonding electrons. In this ion, there are only two atoms: Oxygen and Hydrogen, which will be connected by a single bond.
09

Complete octets and formal charges

Fill up the remaining electrons in the outer shell of each atom, making sure not to exceed the octet rule. Calculate the formal charge of each atom in the ion. The \(\mathrm{OH}^{-}\) Lewis structure becomes: O // H The Oxygen has a formal charge of -1 and bears the formal charge in the ion.

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Most popular questions from this chapter

What is the meaning of the term tertiary \(\left(3^{\circ}\right)\) when it is used to classify amines? Draw a structural formula for the one tertiary \(\left(3^{\circ}\right)\) amine with the molecular formula \(\mathrm{C}_{4} \mathrm{H}_{11} \mathrm{~N}\).

(a) Draw a Lewis structure for the ozone molecule, \(\mathrm{O}_{3}\). (The order of atom attachment is \(\mathrm{O}-\mathrm{O}-\mathrm{O}\), and they do not form a ring.) Chemists use ozone to cleave carboncarbon double bonds (Section 6.5C). (b) Draw four contributing resonance structures; include formal charges. (c) How does the resonance model account for the fact that the length of each \(\mathrm{O}-\mathrm{O}\) bond in ozone \((128 \mathrm{pm})\) is shorter than the \(\mathrm{O}-\mathrm{O}\) single bond in hydrogen peroxide (HOOH, \(147 \mathrm{pm}\) ) but longer than the \(\mathrm{O}-\mathrm{O}\) double bond in the oxygen molecule \((123 \mathrm{pm})\) ?

Write and compare the ground-state electron configurations for each pair of elements. (a) Carbon and silicon (b) Oxygen and sulfur (c) Nitrogen and phosphorus

Draw Lewis structures and condensed structural formulas for the four alcohols with the molecular formula \(\mathrm{C}_{4} \mathrm{H}_{10} \mathrm{O}\). Classify each alcohol as primary, secondary, or tertiary.

Problem \(1.3\) Judging from their relative positions in the Periodic Table, which element in each set is the more electronegative? (a) Lithium or potassium (b) Nitrogen or phosphorus (c) Carbon or silicon
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