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Problem 29
Use data from Appendix C, Figure 7.10, and Figure 7.12 to calculate the lattice energy of RbCl.
Problem 29a
Which neutral atom is isoelectronic with each of the following ions? Ga3+, Zr4+, Mn7+, I−, Pb2+.
Problem 30
Some ions do not have a corresponding neutral atom that has the same electron configuration. For each of the following ions, identify the neutral atom that has the same number of electrons and determine if this atom has the same electron configuration. (a) CI−, (b) Sc3+, (c) Fe2+, (d) Zn2+, (e) Sn4+.
Problem 31b
Consider the isoelectronic ions F- and Na+. (b) Using Equation 7.1 and assuming that core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant, S, calculate Zeff for the 2p electrons in both ions.
Problem 31c
Consider the isoelectronic ions F- and Na+. (c) Repeat this calculation using Slater’s rules to estimate the screening constant, S.
Problem 31d
Consider the isoelectronic ions F- and Na+. (d) For isoelectronic ions, how are effective nuclear charge and ionic radius related?
- Consider the isoelectronic ions Cl- and K+. (b) Using Equation 7.1 and assuming that core electrons contribute 1.00 and valence electrons contribute nothing to the screening constant, S, calculate Zeff for these two ions. (c) Repeat this calculation using Slater’s rules to estimate the screening constant, S.
Problem 32
Problem 33a
Consider S, Cl, and K and their most common ions. (a) List the atoms in order of increasing size.
Problem 33b,c
Consider S, Cl, and K and their most common ions. (b) List the ions in order of increasing size. (c) Explain any differences in the orders of the atomic and ionic sizes.
- Arrange each of the following sets of atoms and ions, in order of increasing size: (a) Pb, Pb2+, Pb4+
Problem 34
Problem 35
Provide a brief explanation for each of the following: (a) Cl- is larger than Ar. (b) P3- is larger than S2-. (c) K+ is larger than Na+. (d) F- is larger than F.
Problem 36b
In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cation–anion distances are 201 pm (Li–F), 282 pm (Na–Cl), 330 pm (K–Br), and 367 pm (Rb–I), respectively. (b) Calculate the difference between the experimentally measured ion–ion distances and the ones predicted from Figure 7.8.
Problem 36c
In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cation–anion distances are 201 pm (Li–F), 282 pm (Na–Cl), 330 pm (K–Br), and 367 pm (Rb–I), respectively. (c) What estimates of the cation– anion distance would you obtain for these four compounds using neutral atom bonding atomic radii? Are these estimates as accurate as the estimates using ionic radii?
- Write equations that show the processes that describe the first, second, and third ionization energies of a chlorine atom. Which process would require the least amount of energy?
Problem 37
- Write equations that show the process for (a) the first two ionization energies of zinc (b) the fourth ionization energy of calcium.
Problem 38
Problem 39
Which element has the highest second ionization energy: Li, K, or Be?
- Identify each statement as true or false: (a) Ionization energies are always endothermic. (b) Potassium has a larger first ionization energy than lithium. (c) The second ionization energy of the sodium atom is larger than the second ionization energy of the magnesium atom. (d) The third ionization energy is three times the first ionization energy of an atom.
Problem 40
Problem 41a
(a) What is the general relationship between the size of an atom and its first ionization energy?
Problem 41b
(b) Which element in the periodic table has the largest ionization energy? Which has the smallest?
- (a) What is the trend in first ionization energies as one proceeds down the group 17 elements? Explain how this trend relates to the variation in atomic radii. (b) What is the trend in first ionization energies as one moves across the fourth period from K to Kr? How does this trend compare with the trend in atomic radii?
Problem 42
Problem 43
Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) Br, Kr (b) C, Ca (c) Li, Rb (d) S, Ge (e) Al, B.
- For each of the following pairs, indicate which element has the smaller first ionization energy: (a) Cs, Cl (b) Fe, Zn (c) I, Cl (d) Se, Sn.
Problem 44
- Give three examples of +2 ions that have an electron configuration of nd10 (n = 3, 4, 5...).
Problem 47
Problem 48
Give examples of transition metal ions with +3 charge that have an electron configuration of nd5 (n = 3, 4, 5...).
- Write an equation for the first electron affinity of helium. Would you predict a positive or a negative energy value for this process? Is it possible to directly measure the first electron affinity of helium?
Problem 49
- If the electron affinity for an element is a negative number, does it mean that the anion of the element is more stable than the neutral atom? Explain.
Problem 50
- Which of the following, I or I-, will have a negative electron affinity?
Problem 51
- What is the relationship between the ionization energy of an anion with a 1- charge, such as F-, and the electron affinity of the neutral atom, F?
Problem 52
- Consider the first ionization energy of neon and the electron affinity of fluorine. (b) These two quantities have opposite signs. Which will be positive, and which will be negative? (c) Would you expect the magnitudes of these two quantities to be equal? If not, which one would you expect to be larger?
Problem 53
- Consider the following equation: Al3+1g2 + e-¡Al2+1g2 Which of the following statements are true? (i) The energy change for this process is the second electron affinity of Al atom since Al2+1g2 is formed. (ii) The energy change for this process is the negative of the third ionization energy of the Al atom. (iii) The energy change for this process is the electron affinity of the Al2+ ion.
Problem 54