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Chapter 9: Problem 92

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 \(\mathrm{NH}_{3} ?\)

Short Answer

Expert verified
In Lewis structures, the central atom is typically less electronegative to enhance molecular stability by allowing more even electron sharing. Exceptions exist, such as H2O and NH3, where Hydrogen atoms surround a more electronegative central atom due to Hydrogen's individual electronegativity and the molecular structure's requirements.

Step by step solution

01

Understanding Electronegativity

Electronegativity refers to an atom's ability to attract electrons towards itself in a chemical bond. The higher the electronegativity, the greater an atom's ability to attract bonding electrons.
02

Explaining Lewis Structures

In drawing plausible Lewis structures, it's often observed that the central atom is less electronegative than the surrounding atoms. This is because the central atom, being less electronegative, has a lesser ability to draw bonding electrons towards itself. This allows the electrons to be more evenly shared across the molecule, enhancing stability.
03

Discussing Exceptions

While this rule generally applies, exceptions exist such as in the case of water (H2O) and ammonia (NH3). Here, the central atoms (Oxygen in H2O and Nitrogen in NH3) are more electronegative than the Hydrogen atoms surrounding them. This is due to the fact that Hydrogen is less electronegative than most other elements, combined with the steric requirements of the molecule that prevent other, more electronegative elements from occupying the central position.

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Most popular questions from this chapter

Draw Lewis structures for the following organic molecules: (a) methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right) ;\) (b) ethanol \(\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right) ;\) (c) tetraethyllead \(\left[\mathrm{Pb}\left(\mathrm{CH}_{2} \mathrm{CH}_{3}\right)_{4}\right]\) which was used in "leaded gasoline"; (d) methylamine \(\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right),\) which is used in tanning; (e) mustard gas \(\left(\mathrm{ClCH}_{2} \mathrm{CH}_{2} \mathrm{SCH}_{2} \mathrm{CH}_{2} \mathrm{Cl}\right),\) a poisonous gas used in World War I; (f) urea [(NH \(_{2}\) ) \(_{2}\) CO], a fertilizer; and (g) glycine \(\left(\mathrm{NH}_{2} \mathrm{CH}_{2} \mathrm{COOH}\right)\), an amino acid.

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: \(\mathrm{Mg}(s) \rightarrow \mathrm{Mg}(g) \Delta H^{\circ}=\) \(148 \mathrm{~kJ} / \mathrm{mol} .]\)

Write Lewis structures for the following molecules and ions: (a) \(\mathrm{NCl}_{3},(\mathrm{~b}) \mathrm{OCS},(\mathrm{c}) \mathrm{H}_{2} \mathrm{O}_{2},(\mathrm{~d}) \mathrm{CH}_{3} \mathrm{COO}^{-}\) (e) \(\mathrm{CN}^{-}\) (f) \(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{NH}_{3}^{+}\)

Define bond length, resonance, and resonance structure. What are the rules for writing resonance structures?

Specify which compound in the following pairs of ionic compounds has the higher lattice energy: (a) \(\mathrm{KCl}\) or \(\mathrm{MgO},\) (b) LiF or \(\mathrm{LiBr}\), (c) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}\) or \(\mathrm{NaCl} .\) Explain your choice.
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