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Chapter 8: Problem 88

A major challenge in implementing the "hydrogen economy" is finding a safe, lightweight, and compact way of storing hydrogen for use as a fuel. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. For example, \(\mathrm{NaAlH}_{4}\) can release \(5.6 \%\) of its mass as \(\mathrm{H}_{2}\) upon decomposing to \(\mathrm{NaH}(s), \mathrm{Al}(s),\) and \(\mathrm{H}_{2}(g) . \mathrm{NaAlH}_{4}\) pos- sesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (a) Write a balanced equation for the decomposition of \(\mathrm{NaAlH}_{4}\). (b) Which element in \(\mathrm{NaAlH}_{4}\) is the most electronegative? Which one is the least electronegative? (c) Based on electronegativity differences, predict the identity of the polyatomic anion. Draw a Lewis structure for this ion. (d) What is the formal charge on hydrogen in the polyatomic ion?

Short Answer

Expert verified
(a) Balanced equation: \(\mathrm{NaAlH}_4 (s) \rightarrow \mathrm{NaH} (s) + \mathrm{Al} (s) + \mathrm{\frac{3}{2} H_2} (g)\). (b) Most electronegative: H, least: Na. (c) Polyatomic anion: \([\mathrm{AlH}_4]^-\); draw Al at the center bonded to 4 H atoms. (d) Formal charge on H in \([\mathrm{AlH}_4]^-\) is 0.

Step by step solution

01

Write the Balanced Equation

The decomposition of sodium aluminum hydride (\(\mathrm{NaAlH}_4\)) releases hydrogen gas. It follows the reaction: \[\mathrm{NaAlH}_4 (s) \rightarrow \mathrm{NaH} (s) + \mathrm{Al} (s) + \mathrm{\frac{3}{2} H_2} (g)\]. This equation is balanced with respect to mass and charge.
02

Identify Most and Least Electronegative Elements

In \(\mathrm{NaAlH}_4\), hydrogen (H) is the most electronegative element, as it is typically more electronegative than sodium (Na) and aluminum (Al), while sodium (Na) is the least electronegative since it is a metal and located further to the left in the periodic table.
03

Determine and Draw the Polyatomic Anion

To determine the anion, consider electronegativity. The polyatomic anion is likely \([\mathrm{AlH}_4]^-\) since aluminum (Al) bonds with hydrogen to form a complex anion. The Lewis structure involves Al at the center with single bonds to four H atoms, surrounded by a negative charge due to gaining an electron, thereby balancing the compound's charge.
04

Calculate Formal Charge on Hydrogen

The formula for formal charge is: \(\text{Formal charge} = \text{valence electrons} - (\text{non-bonding electrons} + \frac{1}{2} \times \text{bonding electrons})\). Hydrogen typically has one valence electron and shares a pair entirely with Al, thus: \(1 - (0 + \frac{1}{2} \times 2) = 0\). Therefore, the formal charge on hydrogen in \([\mathrm{AlH}_4]^-\) is 0.

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Key Concepts

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

Chemical Bonds
Chemical bonds are the glue holding atoms together in molecules and complexes. In a compound like \( \mathrm{NaAlH_4} \), these bonds can be either ionic or covalent. Ionic bonds are formed when one atom donates an electron to another, creating ions. These oppositely charged ions are held together by electrostatic attraction. For instance, in \( \mathrm{NaAlH_4} \), sodium (Na) forms an ionic bond with the \([\mathrm{AlH}_4]^鈥揬) anion.
The polyatomic anion \([\mathrm{AlH}_4]^鈥揬) is held together by covalent bonds. This involves sharing electrons between atoms. In \([\mathrm{AlH}_4]^鈥揬), aluminum forms single covalent bonds with hydrogen atoms, creating a stable arrangement. The presence of both ionic and covalent bonds in \( \mathrm{NaAlH_4} \) allows it to efficiently store hydrogen.
Electronegativity
Electronegativity is a measure of how strongly an atom can attract electrons in a bond. In the context of \( \mathrm{NaAlH_4} \), understanding electronegativity helps identify bond types and predict molecular behavior.
Hydrogen, being more electronegative than sodium and aluminum, tends to attract electrons more strongly. Unlike sodium, which sits further left and lower on the periodic table, giving it a lower electronegativity. This means sodium easily loses electrons, forming positive ions.
  • Hydrogen: Most electronegative in \( \mathrm{NaAlH_4} \)
  • Sodium: Least electronegative
Recognizing these differences in electronegativity helps predict the formation of the \([\mathrm{AlH}_4]^鈥揬) anion, as aluminum's greater tendency to hold onto electrons over sodium aids in this bond formation.
Lewis Structure
The Lewis structure is a simplified representation of the arrangement of atoms in a molecule, showing how they bond and accommodate electrons. For the \([\mathrm{AlH}_4]^鈥揬) anion, constructing the Lewis structure begins with placing aluminum (Al) at the center since it's the most likely to form multiple bonds due to its size and valence.
Aluminum connects to the four hydrogen atoms, forming single covalent bonds with each. The overall negative charge on \([\mathrm{AlH}_4]^鈥揬) indicates an extra electron in the structure, balancing the compound's formal charge. It is often represented by placing this charge outside the brackets enclosing the structure.
When drawing, ensure Al is bonded to H with lines representing shared pairs of electrons. This simple depiction helps in understanding how electrons are distributed in the anion and verifies that the formal charge on each hydrogen remains \(0\). It shows a clear picture of how the hydride functions and maintains stability within the atomic structure.

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Most popular questions from this chapter

(a) Draw the best Lewis structure(s) for the nitrite ion, \(\mathrm{NO}_{2}^{-}\). (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in \(\mathrm{NO}_{2}^{-}\) relative to \(\mathrm{N}-\mathrm{O}\) single bonds and double bonds?

The hypochlorite ion, \(\mathrm{ClO}^{-},\) is the active ingredient in bleach. The perchlorate ion, \(\mathrm{ClO}_{4}^{-},\) is a main component of rocket propellants. Draw Lewis structures for both ions. (a) What is the formal charge of Cl in the hypochlorite ion? (b) What is the formal charge of Cl in the perchlorate ion, assuming the Cl-O bonds are all single bonds? (c) What is the oxidation number of Cl in the hypochlorite ion? (d) What is the oxidation number of \(\mathrm{Cl}\) in the perchlorate ion, assuming the Cl- O bonds are all single bonds? (e) In a redox reaction, which ion would you expect to be more easily reduced?

By referring only to the periodic table, select \((\mathbf{a})\) the most electronegative element in group \(13 ;(\mathbf{b})\) the least electronegative element in the group As, Se and Br; (c) the most electronegative element in the group K, Mg, Al and In; (d) the element in the group \(\mathrm{Na}\), Be, \(\mathrm{Si}\), Ar, that is most likely to form an ionic compound with \(B \mathrm{r}\).

(a) Describe the molecule xenon trioxide, \(\mathrm{XeO}_{3},\) using four possible Lewis structures, one each with zero, one, two, or three \(\mathrm{Xe}-\mathrm{O}\) double bonds. (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? (c) Do any of the four Lewis structures have multiple resonance structures? If so, how many resonance structures do you find? (d) Which of the Lewis structures in part (a) yields the most favorable formal charges for the molecule?

(a) Does the lattice energy of an ionic solid increase or decrease (i) as the charges of the ions increase, (ii) as the sizes of the ions increase? (b) Arrange the following substances not listed in Table 8.1 according to their expected lattice energies, listing them from lowest lattice energy to the highest: \(\mathrm{MgS},\) KI, GaN, LiBr.
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