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Chapter 2: Problem 41

Consider the Lewis structures for the following molecules: $$\begin{equation} \mathrm{CO}_{2}, \mathrm{CO}_{3}^{2-}, \mathrm{NO}_{2}^{-}, \text {and } \mathrm{NO}_{3}^{-} \end{equation}$$ Which molecule would have the smallest bond angle between terminal atoms? (A) \(\mathrm{CO}_{2}\) (B) \(\mathrm{CO}_{3}^{2-}\) (C) \(\mathrm{NO}_{2}^{-}\) (D) \(\mathrm{NO}_{3}^{-}\)

Short Answer

Expert verified
So the molecule with the smallest bond angle among the options is NO2-.

Step by step solution

01

Understand Lewis Structures

Content for Step 1: You need to draw the Lewis structure of each molecule. Lewis structures show all valence electrons and therefore help to visualize molecular geometry.
02

Determine the Geometry

Content for Step 2: Using the Lewis structures, you can determine the geometry of each molecule. CO2 is linear, CO32- is trigonal planar, NO2- is bent, and NO3- is trigonal planar.
03

Evaluate Bond Angles

Content for Step 3: Then you need to denote the bond angles within each molecule. Linear molecules have bond angles of 180 degrees. Trigonal planar molecules have bond angles of 120 degrees. A bent molecule, on the other hand, due to the presence of lone pair(s) of electrons has bond angles less than120 degrees
04

Compare Angles

Content for Step 4: Lastly, compare the bond angles. The molecule with the smallest angle is the one having the bond angle less than 120 degrees that is present in the bent molecule -- NO2-, due to the existence of a lone pair of electrons.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Lewis Structures
Lewis structures are a way to represent molecules that show each atom and its valence electrons. This helps you understand how atoms connect to form molecules. For each atom, dots represent valence electrons, which are the electrons available for bonding. By using Lewis structures, you can visualize how atoms form different bonds.
  • Draw each atom and its valence electrons as dots around the element symbols.
  • Connect these atoms by sharing pairs of electrons to form bonds.
  • Double and triple bonds might be used if needed to satisfy each atom's full valence shell, typically eight electrons.
Each molecule in this exercise has a unique Lewis structure that gives clues to the overall structure and bonding nature. For instance, in \(\mathrm{NO}_{2}^{-}\), one lone pair and two bonded pairs are present on the nitrogen atom, resulting in a bent shape.
Molecular Geometry
Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule. It determines many of the molecule's properties, such as polarity and reactivity. This arrangement is mainly influenced by the number of bonds and lone pairs around a central atom.Key points about molecular geometry include:
  • Linear geometry results when there are no lone pairs, like in CO鈧�, where carbon forms two double bonds with oxygen atoms.
  • Trigonal planar geometry occurs when there are three regions of electron density around a central atom, such as in \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{NO}_{3}^{-}\).
  • Bent geometry appears due to lone pairs causing repulsion, leading to a smaller bond angle as seen in \(\mathrm{NO}_{2}^{-}\).
Understanding the molecular geometry helps you predict how molecules will interact in chemical reactions.
Bond Angles
Bond angles are the angles formed between three atoms across at least two bonds. Theoretical bond angles are altered by the presence of lone pairs of electrons, which exert more repulsion than bonding pairs.
  • Linear molecules, such as \(\mathrm{CO}_{2}\), typically have bond angles of 180掳.
  • For trigonal planar molecules, such as \(\mathrm{CO}_{3}^{2-}\) and \(\mathrm{NO}_{3}^{-}\), bond angles are typically around 120掳.
  • Bent molecules, such as \(\mathrm{NO}_{2}^{-}\), have bond angles less than 120掳 due to the repulsion by the lone pair, which decreases the angle.
In this exercise, \(\mathrm{NO}_{2}^{-}\) has the smallest bond angles because lone pairs push the bonded atoms closer together.
Valence Electrons
Valence electrons are the outermost electrons of an atom available for bonding. They play an essential role in determining the chemical properties and reactivity of the element. In the Lewis structure, each dot represents a valence electron. Here's how valence electrons influence bonding:
  • Atoms share or transfer valence electrons to achieve a stable electron configuration similar to noble gases.
  • In covalent bonding, atoms share valence electrons to fill their outermost shell, as demonstrated in the Lewis structures of the molecules discussed.
  • The number of valence electrons dictates how many bonds an atom can form. For example, oxygen forms two bonds because it has six valence electrons.
By understanding the number of valence electrons, you can deduce the structure and bonds within the molecule, crucial for predicting molecular properties and behaviors.

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Most popular questions from this chapter

Consider the Lewis structures for the following molecules: $$\begin{equation} \mathrm{CO}_{2}, \mathrm{CO}_{3}^{2-}, \mathrm{NO}_{2}^{-}, \text {and } \mathrm{NO}_{3}^{-} \end{equation}$$ Which molecule would have the shortest bonds? (A) \(\mathrm{CO}_{2}\) (B) \(\mathrm{CO}_{3}^{2-}\) (C) \(\mathrm{NO}_{2}^{-}\) (D) \(\mathrm{NO}_{3}^{-}\)

A sample of \(\mathrm{H}_{2} \mathrm{S}\) gas is placed in an evacuated, sealed container and heated until the following decomposition reaction occurs at \(1000 \mathrm{K} :\) $$2 \mathrm{H}_{2} \mathrm{S}(g) \rightarrow 2 \mathrm{H}_{2}(g)+\mathrm{S}_{2}(g) \qquad K_{\mathrm{c}}=1.0 \times 10^{-6}$$ As the reaction progresses at a constant temperature of 1000 K, how does the value for the Gibbs free energy constant for the reaction change? (A) It stays constant. (B) It increases exponentially. (C) It increases linearly. (D) It decreases exponentially.

A sample of oxygen gas at \(50^{\circ} \mathrm{C}\) is heated, reaching a final temperature of \(100^{\circ} \mathrm{C} .\) Which statement best describes the behavior of the gas molecules? (A) Their velocity increases by a factor of two. (B) Their velocity increases by a factor of four. (C) Their kinetic energy increases by a factor of 2. (D) Their kinetic energy increases by a factor of less than 2.

Which of the following is true for all bases? (A) All bases donate \(\mathrm{OH}^{-}\) ions into solution. (B) Only strong bases create solutions in which \(\mathrm{OH}^{-}\) ions are present. (C) Only strong bases are good conductors when dissolved in solution. (D) For weak bases, the concentration of the \(\mathrm{OH}^{-}\) ions exceeds the concentration of the base in the solution.

A solution of sulfurous acid, \(\mathrm{H}_{2} \mathrm{SO}_{3}\) , is present in an aqueous solution. Which of the following represents the concentrations of three different ions in solution? (A) \(\left[\mathrm{SO}_{3}^{2-}\right]>\left[\mathrm{HSO}_{3}^{-}\right]>\left[\mathrm{H}_{2} \mathrm{SO}_{3}\right]\) (B) \(\left[\mathrm{H}_{2} \mathrm{SO}_{3}\right]>\left[\mathrm{HSO}_{3}^{-}\right]>\left[\mathrm{SO}_{3}^{2-}\right]\) (C) \(\left[\mathrm{H}_{2} \mathrm{SO}_{3}\right]>\left[\mathrm{HSO}_{3}^{-}\right]=\left[\mathrm{SO}_{3}^{2-}\right]\) (D) \(\left[\mathrm{SO}_{3}^{2-}\right]=\left[\mathrm{HSO}_{3}^{-}\right]>\left[\mathrm{H}_{2} \mathrm{SO}_{3}\right]\)
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