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Chapter 3: Problem 59

Which compound in each of the following pairs of ionic substances has the most negative lattice energy? Justify your answers. a. NaCl, KCl b. LiF, LiCl c. \(\mathrm{Mg}(\mathrm{OH})_{2}, \mathrm{MgO}\) d. \(\mathrm{Fe}(\mathrm{OH})_{2}, \mathrm{Fe}(\mathrm{OH})_{3}\) e. \(\mathrm{NaCl}, \mathrm{Na}_{2} \mathrm{O}\) f. \(\mathrm{MgO}, \mathrm{BaS}\)

Short Answer

Expert verified
a. NaCl has the most negative lattice energy because Na is smaller than K. b. LiF has the most negative lattice energy because F is smaller than Cl. c. MgO has the most negative lattice energy because O2- is smaller than OH-. d. Fe(OH)3 has the most negative lattice energy due to its higher cation charge e. Na2O has the most negative lattice energy because O2- has a higher charge than Cl-. f. MgO has the most negative lattice energy due to both its ion charges and ion sizes.

Step by step solution

01

a. NaCl vs. KCl

Between NaCl and KCl, the charges on the ions are the same (+1 for Na and K, -1 for Cl), so the deciding factor will be the size of the ions. Na is smaller than K, so the Na+ ion and the Cl- ion will be closer together, therefore NaCl will have a more negative lattice energy.
02

b. LiF vs. LiCl

In this case, the cations are the same (Li+), so the deciding factor will be the size of the anions. F is smaller than Cl, so the force of attraction will be stronger in LiF. Thus, LiF will have a more negative lattice energy than LiCl.
03

c. \(\mathrm{Mg}(\mathrm{OH})_{2}\) vs. \(\mathrm{MgO}\)

For these compounds, the cations are the same (Mg2+), as well as the magnitude of the anions charges(-2 in both cases). However, the size of the anions differs. O2- is smaller than the OH- anion, so MgO has a more negative lattice energy than \(\mathrm{Mg}(\mathrm{OH})_{2}\).
04

d. \(\mathrm{Fe}(\mathrm{OH})_{2}\) vs. \(\mathrm{Fe}(\mathrm{OH})_{3}\)

In this case, the cations are Fe(II) in \(\mathrm{Fe}(\mathrm{OH})_{2}\) and Fe(III) in \(\mathrm{Fe}(\mathrm{OH})_{3}\). So, there is a difference in charge between the two compounds. The cation with a higher charge (Fe(III)) will lead to a more negative lattice energy, so \(\mathrm{Fe}(\mathrm{OH})_{3}\) will have the most negative lattice energy.
05

e. \(\mathrm{NaCl}\) vs. \(\mathrm{Na}_{2} \mathrm{O}\)

These compounds have the same cation (Na+), so the deciding factor will be the charges of the anions. O2- has a higher anion charge than Cl-, and as a result, \(\mathrm{Na}_{2} \mathrm{O}\) will have the most negative lattice energy.
06

f. \(\mathrm{MgO}\) vs. \(\mathrm{BaS}\)

For this pair, we need to look at both the cation and anion properties. Mg2+ has a higher charge than Ba2+, and O2- has a higher charge than S2-. However, Mg2+ is smaller than Ba2+ and O2- is smaller than S2-. Due to both the charge and size of the ions, MgO has a more negative lattice energy than BaS.

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Most popular questions from this chapter

Which member of the following pairs would you expect to be more energetically stable? Justify each choice. a. NaBr or \(\mathrm{NaBr}_{2}\) b. \(\mathrm{ClO}_{4}\) or \(\mathrm{ClO}_{4}^{-}\) c. \(\mathrm{SO}_{4}\) or \(\mathrm{XeO}_{4}\) d. \(\mathrm{OF}_{4}\) or \(\mathrm{SeF}_{4}\)

Write Lewis structures that obey the octet rule for the following species. Assign the formal charge for each central atom. a. \(\mathrm{POCl}_{3}\) b. \(\mathrm{SO}_{4}^{2-}\) c. \(\mathrm{ClO}_{4}\) d. \(\mathrm{PO}_{4}^{3-}\) e. \(\mathrm{SO}_{2} \mathrm{Cl}_{2}\) f. \(\quad X \in O_{4}\) g. \(\mathrm{ClO}_{3}\) h. \(\mathrm{NO}_{4}^{3-}\)

Write Lewis structures that obey the octet rule for each of the following molecules and ions. (In each case the first atom listed is the central atom.) a. \(\mathrm{POCl}_{3}, \mathrm{SO}_{4}^{2-}, \mathrm{XeO}_{4}, \mathrm{PO}_{4}^{3-}, \mathrm{ClO}_{4}^{-}\) b. \(\mathrm{NF}_{3}, \mathrm{SO}_{3}^{2-}, \mathrm{PO}_{3}^{3-}, \mathrm{ClO}_{3}^{-}\) c. \(\mathrm{ClO}_{2}^{-}, \mathrm{SCl}_{2}, \mathrm{PCl}_{2}^{-}\) d. Considering your answers to parts a, b, and c, what conclusions can you draw concerning the structures of species containing the same number of atoms and the same number of valence electrons?

Write Lewis structures for \(\mathrm{CO}_{3}^{2-}, \mathrm{HCO}_{3}^{-},\) and \(\mathrm{H}_{2} \mathrm{CO}_{3}\). When acid is added to an aqueous solution containing carbonate or bicarbonate ions, carbon dioxide gas is formed. We generally say that carbonic acid \(\left(\mathrm{H}_{2} \mathrm{CO}_{3}\right)\) is unstable. Use bond energies to estimate \(\Delta E\) for the reaction (in the gas phase) $$\mathrm{H}_{2} \mathrm{CO}_{3} \longrightarrow \mathrm{CO}_{2}+\mathrm{H}_{2} \mathrm{O}$$Specify a possible cause for the instability of carbonic acid.

One type of exception to the octet rule are compounds with central atoms having fewer than eight electrons around them. \(\mathrm{BeH}_{2}\) and \(\mathrm{BH}_{3}\) are examples of this type of exception. Draw the Lewis structures for \(\mathrm{BeH}_{2}\) and \(\mathrm{BH}_{3}\)
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