The substances NaF and CaO are isoelectronic (have the same number of valence electrons). (b) What are the charges of each of the anions in each compound?

The substances NaF and CaO are isoelectronic (have the same number of valence electrons). (d) Using the lattice energies in Table 8.1, predict the lattice energy of ScN.

(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?

Consider the ionic compounds KF, NaCl, NaBr, and LiCl. (a) Use ionic radii (Figure 7.8) to estimate the cation–anion distance for each compound.

Which of the following trends in lattice energy is due to differences in ionic radii: a. NaCl > RbBr > CsBr, b. BaO > KF, c. SrO > SrCl2?

Energy is required to remove two electrons from Ca to form Ca2+, and energy is required to add two electrons to O to form O2 - . Yet CaO is stable relative to the free elements. Which statement is the best explanation? (a) The lattice energy of CaO is large enough to overcome these processes. (b) CaO is a covalent compound, and these processes are irrelevant. (c) CaO has a higher molar mass than either Ca or O. (d) The enthalpy of formation of CaO is small. (e) CaO is stable to atmospheric conditions.

List the individual steps used in constructing a Born–Haber cycle for the formation of BaI₂ from the elements. Which of the steps would you expect to be exothermic?

Use data from Appendix C, Figure 7.11, and Figure 7.13 to calculate the lattice energy of KI.

(a) Based on the lattice energies of MgCl2 and SrCl2 given in Table 8.1, what is the range of values that you would expect for the lattice energy of CaCl2?

(b) Using data from Appendix C, Figure 7.11, Figure 7.13, and the value of the second ionization energy for Ca, 1145 kJ/mol, calculate the lattice energy of CaCl2.

(b) A substance, XY, formed from two different elements, melts at −33 °C. Is XY likely to be a covalent or an ionic substance?

Using Lewis symbols and Lewis structures, diagram the formation of SiCl4 from Si and Cl atoms, showing valence-shell electrons. e. How many bonding pairs of electrons are in the SiCl4 molecule?

Using Lewis symbols and Lewis structures, diagram the formation of PF₃ from P and F atoms, showing valence-shell electrons. (a) How many valence electrons does P have initially? (c) How many valence electrons surround the P in the PF₃ molecule? (d) How many valence electrons surround each P in the PF₃ molecule?

Use Lewis symbols and Lewis structures to diagram the formation of PF3 from P and F atoms, showing valence-shell electrons. e. How many bonding pairs of electrons are in the PF3 molecule?

(b) How many bonding electrons are in the structure?

(c) Do you expect the O—O bond in H₂O₂ to be longer or shorter than the O—O bond in O₂? Explain.

Which of the following statements about electronegativity is or are true? i. The alkali metals are the family with the largest electronegativity values. ii. The numerical values for electronegativity have no units. iii. Electronegativity is the ability of an atom in a molecule to attract electron density toward itself.

Place the following pairs of elements in order from smallest to largest difference in electronegativity: K and F, S and O, Br and I, Ca and Se, Li and Cl.

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Se, Te, Br, I; (b) Be, Mg, C, Si (c) Al, Si, P, S (d) Zn, Ge, Ga, As.

By referring only to the periodic table, select (d) the element in the group K, C, Zn, F that is most likely to form an ionic compound with Ba.

Which of the following bonds are polar: a. B —F, b. Cl — Cl, c. Se — O, d. H—I? Which is the more electronegative atom in each polar bond?

Arrange the bonds in each of the following sets in order of increasing polarity: (a) C—F, O—F, Be—F (b) O—Cl,S—Br, C—P

Arrange the bonds in each of the following sets in order of increasing polarity: (c) C—S, B— F, N — O.

(a) From the data in Table 8.2, calculate the effective charges on the H atom of the HBr molecule in units of the electronic charge, e.

(b) If you were to put HBr under very high pressure, so its bond length decreased significantly, would its dipole moment increase, decrease, or stay the same, if you assume that the effective charges on the atoms do not change?

The bromine monofluoride molecule, BrF, has a bond length of 1.76 Å and a dipole moment of 1.29 D. a. Which atom of the molecule is expected to have a negative charge?

In the following pairs of binary compounds, determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) SiF₄ and LaF₃ (for ionic or molecular substances) to assign a name to each compound: (b) FeCl₂ and ReCl₆ (c) PbCl₄ and RbCl.

In the following pairs of binary compounds, determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) TiCl₄ and CaF₂ (b) ClF₃ and VF₃

In the following pairs of binary compounds, determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (c) SbCl₅ and AlF₃.