A metal oxide has a lattice energy in the range of 3300 kJ/mol. Is the charge on the metal ion M likely to be 1+, 2+, or 3+? Explain.

Verified step by step guidance

Step 1: Understand the concept of lattice energy. Lattice energy is the energy required to separate one mole of an ionic solid into its gaseous ions. It is influenced by the charges on the ions and the distance between them.

Step 2: Recall the relationship between lattice energy and ionic charge. Lattice energy increases with the magnitude of the charge on the ions. Higher charges result in stronger electrostatic attractions, leading to higher lattice energies.

Step 3: Consider the typical lattice energy values for different charges. Generally, compounds with 1+ and 1- charges have lower lattice energies, while those with 2+ and 2- charges have higher lattice energies. Compounds with 3+ and 3- charges have even higher lattice energies.

Step 4: Compare the given lattice energy with typical values. A lattice energy of 3300 kJ/mol is relatively high, suggesting that the metal ion likely has a higher charge, such as 2+ or 3+.

Step 5: Conclude based on the comparison. Given the high lattice energy, it is more likely that the metal ion M has a charge of 2+ or 3+, as these charges correspond to stronger ionic interactions and higher lattice energies.

Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Lattice Energy

Lattice energy is the amount of energy released when gaseous ions combine to form an ionic solid. It is a measure of the strength of the forces between the ions in an ionic compound. Higher lattice energy indicates stronger ionic bonds, which typically results from higher charges on the ions and smaller ionic radii.

Ionic Charge

The charge on an ion is determined by the number of electrons lost or gained by an atom. In metals, a higher positive charge (e.g., 2+ or 3+) usually corresponds to the loss of more electrons. The charge affects the lattice energy; ions with higher charges lead to stronger electrostatic attractions, resulting in higher lattice energies.

Trends in Lattice Energy

Lattice energy trends can be predicted based on the charges and sizes of the ions involved. Generally, as the charge on the ions increases, the lattice energy increases due to stronger attractions. Conversely, larger ionic radii decrease lattice energy because the distance between the ions increases, weakening the electrostatic forces.

Recommended video:

Guided course

00:49

Lattice Energy

Related Practice

Textbook Question

Consider the lattice energies of the following Group 2A compounds: BeH₂, 3205 kJ/mol; MgH₂, 2791 kJ/mol; CaH₂, 2410 kJ/mol; SrH₂, 2250 kJ/mol; BaH₂, 2121 kJ/mol. (a) What is the oxidation number of H in these compounds?

Textbook Question

Consider the lattice energies of the following Group 2A compounds: BeH₂, 3205 kJ/mol; MgH₂, 2791 kJ/mol; CaH₂, 2410 kJ/mol; SrH₂, 2250 kJ/mol; BaH₂, 2121 kJ/mol. (c) Consider BeH₂. Does it require 3205 kJ of energy to break one mole of the solid into its ions, or does breaking up one mole of solid into its ions release 3205 kJ of energy?

Textbook Question

Consider the lattice energies of the following Group 2A compounds: BeH₂, 3205 kJ/mol; MgH₂, 2791 kJ/mol; CaH₂, 2410 kJ/mol; SrH₂, 2250 kJ/mol; BaH₂, 2121 kJ/mol. (d) The lattice energy of ZnH₂ is 2870 kJ/mol. Considering the trend in lattice enthalpies in the Group 2 compounds, predict which Group 2 element is most similar in ionic radius to the Zn²⁺ ion.

Textbook Question

The ionic compound CaO crystallizes with the same structure as sodium chloride (Figure 8.3). (a) In this structure, how many O2- are in contact with each Ca2+ ion (Hint: Remember the pattern of ions shown in Figure 8.3 repeats over and over again in all three directions.)

Textbook Question

Construct a Born–Haber cycle for the formation of the hypothetical compound NaCl₂, where the sodium ion has a 2⁺ charge (the second ionization energy for sodium is given in Table 7.2). (a) How large would the lattice energy need to be for the formation of NaCl₂ to be exothermic?