Chapter 5

(a) Which ions in this list are likely to be found in ionic compounds: \(\mathrm{K}^{2+}, \mathrm{Cs}^{+}, \mathrm{Al}^{4+}, \mathrm{F}^{2-}, \mathrm{Se}^{2-} ?\) (b) Which, if any, of these ions have a noble-gas configuration?

Write the atomic orbital diagram for the \(4 s\) and \(3 d\) electrons in a (a) vanadium atom. (b) \(\mathrm{V}^{2+}\) ion. (c) \(\mathrm{V}^{4+}\) ion.

Write the atomic orbital diagram for the \(4 s\) and \(3 d\) electrons in a (a) manganese atom. (b) \(\mathrm{Mn}^{+}\) ion. (c) \(\mathrm{Mn}^{3+}\) ion.

Give the electron configurations of \(\mathrm{Mn}, \mathrm{Mn}^{2+},\) and \(\mathrm{Mn}^{3+}\). Use atomic orbital box diagrams to determine the number of unpaired electrons for each species.

Write the electron configurations of chromium: \(\mathrm{Cr}\), \(\mathrm{Cr}^{2+}\), and \(\mathrm{Cr}^{3+}\). Use atomic orbital box diagrams to determine the number of unpaired electrons for each species.

Write electron configurations for these elements. (a) Zirconium (Zr). This metal is exceptionally resistant to corrosion and so has important industrial applications. Moon rocks show a surprisingly high rirconium content compared with rocks on Earth. (b) Rhodium (Rh), which is used in jewelry and in industrial catalysts.

The lanthanides, or rare earths, are only "medium rare," because all can be purchased for a reasonable pricc. Give clectron configurations for atoms of thesc lanthanides. (a) Europium (Eu), the most expensive of the rare earth elements; \(1 \mathrm{~g}\) ean be purchased for ahout $$\$ 1.40$$. (b) Ytterbium (Yb). Less expensive than Eu, Yb costs only about $$\$ 0.35$$ per gram. It was named for the village of Yuerby in Sweden, where a mineral source of the element was found.

Locate these elements in the periodic table, and then draw a Lewis dot symbol that represents the number of valcnce electrons for an atom of each. (a) \(\mathrm{Sr}\) (b) \(B r\) (c) \(\mathrm{Ga}\) (d) \(\mathrm{Sb}\)

Locate these elements in the periodic table, and then draw a Lewis dot symbol that represents the number of valence clectrons for an atom of each. (a) \(\mathrm{F}\) (b) In (c) Te (d) \(\mathrm{Cs}\)

Give the electron configurations of these ions, and indicate which ones are isoelectronic. (a) \(\mathrm{Ca}^{2+}\) (b) \(\mathrm{K}^{+}\) (c) \(\mathrm{O}^{2-}\)

Give the electron configurations of these ions, and indicate which ones are isoelectronic. (a) \(\mathrm{Na}^{+}\) (b) \(\mathrm{Al}^{3+}\) (c) \(\mathrm{Cl}^{-}\)

(a) What is the electron configuration for an atom of tin? (b) What are the electron configurations for \(\mathrm{Sn}^{2+}\) and \(\mathrm{Sn}^{4+}\) ions?

What is the electron configuration for (a) a bromine atom? (b) a bromide ion?

(a) In the first transition series (in row four of the periodic table), which elements would you predict to be diamagnetic? (b) Which element in this series has the greatest number of unpaired electrons?

Nearly all first-row transition elements form \(2+\) ions. (a) For which of these elements are the \(2+\) ions paramagnetic? (b) For which element do compounds containing \(2+\) ions and chloride ions have the greatest paramagnetism? (Chloride ions have no unpaired electrons.)

How do the spins of unpaired clectrons from paramagnetic and ferromagnctic matcrials differ in their behavior in a magnctic ficld?

Consider titanium metal and its two oxides, TiO and \(\mathrm{TiO}_{2}\). The oxide ion has no unpaired clectrons. (a) Which of these titanium species is diamagnetic? Fxplain your answer. (b) Which titanium species will he most attracted to a magmetic field? Fxplain your answer.

The acctylacctonate ion, (acac) . has no unpaircd clectrons. It forms compounds with Fe \(^{21}\) and with \(\mathrm{Fc}^{3+}\) ions whosc formulas are \(\mathrm{Fc}(\mathrm{acac})_{2}\) and Fe(acac) \(_{3}\), respectively. Explain which compound has the grcater attraction to a magnctic ficld.

Use electron configurations to explain why (a) sulfur has a lower electron affinity than chlorine. (b) boron has a lower first ionization energy than beryllium. (c) chlorine has a lower first ionization energy than fluorine. (d) oxygen has a lower tirst ionization energy than nitrogen. (e) iodine has a lower electron affinity than bromine.

A fellow chemistry student states that periodic trends cover the main-group elements but do not apply to the transition metals. Is the student correet? Explain.

Arrange these elements in order of increasing effective nuclear charge: \(\mathrm{N}, \mathrm{F}, \mathrm{B}, \mathrm{O},\) and \(\mathrm{C}\)

Arrange these ions in order of decreasing size: \(\mathrm{Be}^{2+}\), \(\mathrm{Rb}^{+}, \mathrm{C}_{\mathrm{a}}^{2+}, \mathrm{Mg}^{2+}\)

Arrange these elements in order of increasing atomic size: \(\mathrm{Ca}, \mathrm{Rb}, \mathrm{P}, \mathrm{Ge}, \mathrm{Sr}\). (Try arranging these without looking at Figure 5.25 and then check yourself by looking up the necessary atomic radii.)

Arrange these elements in order of increasing atomic size: \(\mathrm{Al}, \mathrm{B}, \mathrm{C}, \mathrm{K}, \mathrm{Na}\). (Try arranging these without looking at Figure 5.25 and then check yourself by looking up the necessary atomic radii.)

Select the atom or ion in each pair that has the smaller radius. (a) Cs or Rb (b) \(\mathrm{O}^{2-}\) or \(\mathrm{O}\) (c) Br or As (d) \(\mathrm{Ba}\) or \(\mathrm{Ba}^{2+}\) (e) \(\mathrm{Cl}^{-}\) or \(\mathrm{Ca}^{2+}\)

Select the atom or ion in each pair that has the larger radius. (a) \(\mathrm{Cl}\) or \(\mathrm{Cl}^{-}\) (b) Ca or \(\mathrm{Ca}^{2+}\) (c) Al or \(N\) (d) \(\mathrm{Cl}^{-}\) or \(\mathrm{K}^{+}\) (e) \(\ln\) or Sn

Which of these groups of elements is arranged correctly in order of increasing ionization energy? (a) \(\mathrm{C}\). \(\mathrm{Si}, \mathrm{Li}, \mathrm{Ne}\) (b) Ne, Si, \(\mathrm{C}, \mathrm{Li}\) (c) Li, Si, \(\mathrm{C}, \mathrm{Ne}\) (d) \(\mathrm{Ne}, \mathrm{C}, \mathrm{Si}, \mathrm{L} \mathrm{i}\)

Rank these ionization energies (IE) from the smallest to the largest value. Briefly explain your answer. (a) First \(I E\) of \(B e\) (b) First IE of \(\mathrm{Li}\) (c) Second IE of \(\mathrm{Be}\) (d) Second IE of Na (e) First IE of \(\mathrm{K}\)

Predict which of these clements would have the greatest difference betwcen the first and sccond ionization cncrgics: Si, \(\mathrm{Na}, \mathrm{P}, \mathrm{Mg} .\) Bricfly explain your answer.

Compare the clemcnts \(\mathrm{B}, \mathrm{Al}, \mathrm{C}, \mathrm{Si}\). (a) Which has the most metallic character? (b) Which has the largest atomic radius? (c) Arrange the three elements \(\mathrm{B}, \mathrm{Al},\) and \(\mathrm{C}\) in order of increasing first ionization energy.

Compare the elements \(\mathrm{I} \mathrm{i}, \mathrm{K}, \mathrm{C}, \mathrm{N}\). (a) Which has the largest atomic radius? (b) Arrange the elements in order of increasing first ionization cnergy.

The first electron affinity of oxygen is negative, the second is positive. Fxplain why this change in sign occurs.

Explain why nitrogen has a higher first ionization energy than does carbon, the element that precedes it in the periodic table.

Which group of the periodic table has elements with high first ionization encrgics and very negative electron affinities? Explain this behavior.

Determine the lattice energy for L.iCl(s) given these data: Sublimation enthalpy of \(\mathrm{Li}, 161 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{IF}_{1}\) for \(\mathrm{Li}\), \(520 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{BE}\) of \(\mathrm{Cl}_{2}(\mathrm{~g}), 242 \mathrm{~kJ} / \mathrm{mol} ;\) electron affinity of \(\mathrm{Cl},-349 \mathrm{~kJ} / \mathrm{mol} ;\) formation enthalpy of \(\mathrm{LiCl}(\mathrm{s})\) \(-408.7 \mathrm{~kJ} / \mathrm{mol}\)

Which ionic compound has the lowest melting point? Explain your choice. \(\begin{array}{lll}\mathrm{LiCl} & \mathrm{NaBr} & \mathrm{KCl}\end{array}\)

Which ionic compound has the largest lattice energy? Explain your choice. \(\begin{array}{lll}\mathrm{MgS} & \mathrm{RbI} & \mathrm{Li}_{2} \mathrm{~S}\end{array}\)

Give the symbol of the ground-state atom that (a) is in Group 8 A but has no \(p\) electrons. (b) has a single electron in the \(3 d\) subshell. (c) forms a \(1+\) ion with a \(1 s^{2} 2 s^{2} 2 p^{6}\) electron configuration.

Give the symbol of all the ground-state atoms that have (a) no \(p\) electrons. (b) from two to four \(d\) electrons. (c) from two to four \(s\) electrons.

Answer these questions about the elements \(X\) and \(Z\), which have the electron configurations shown. $$ \mathrm{X}=[\mathrm{Kr}] 4 d^{10} 5 s^{1} \quad Z=[\mathrm{Ar}] 3 d^{10} 4 s^{2} 4 p^{4} $$ (a) Is element \(X\) a metal or a nonmetal? (b) Which element has the larger atomic radius? (c) Which element would have the greater first ionization energy?

(a) Rank these elements in order of increasing atomic radius: \(\mathrm{O}, \mathrm{S}, \mathrm{F}\). Briefly explain your reasoning. (b) Which element has the largest first ionization energy: P, Si, S, Se? Briefly explain your reasoning.

(a) Rank these in order of increasing radius: \(\mathrm{Ne}, \mathrm{O}^{2-}, \mathrm{N}^{3-}\), \(\mathrm{F}^{-}\). Briefly explain your reasoning- (b) Place these elements in order of increasing first ionization energy: Cs, Sr, Ba. Briefly explain your reasoning.

Name the clement corrcsponding to cach of these charactcristics. (a) The element whose atoms have the electron configura- $$ \text { tion } 1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{4} $$ (b) The element in the alkaline-earth group that has the largest atomic radius (c) The element in Group \(5 \Lambda\) whose atoms have the largest first ionization energy (d) The element whose \(2+\) ion has the configuration \([\mathrm{Kr}] 4 d^{6}\) (e) The clement whose neutral atoms have the electron configuration \(\lfloor\mathrm{\Lambdar}] 3 d^{10} 4 s^{1}\)

The ionization cnergies for the removal of the first electron from atoms of \(\mathrm{Si}, \mathrm{P}, \mathrm{S},\) and \(\mathrm{Cl}\) are listed below. Briefly rationalize this trend.

Answer these questions about the elements with the electron configurations shown. $$ \mathrm{X}=\left[\begin{array}{ll} \mathrm{Ar}] & 3 d^{\mathrm{s}} 4 s^{2} \end{array} \quad \mathrm{Z}=[\mathrm{Ar}] 3 d^{10} 4 s^{2} 4 p^{5}\right. $$ (a) An atom of which element is expected to have the larger first ionization energy? (b) An atom of which element would be the smaller of the two?

Place these atoms and ions in order of decreasing size: \(\mathrm{Ar}, \mathrm{K}^{+}, \mathrm{Cl}^{-}, \mathrm{S}^{2-}, \mathrm{Ca}^{2+} .\) Briefly explain your reasoning.

Which of these ions are unlikely, and why: \(\mathrm{Cs}^{+}, \mathrm{In}^{4+},\) \(\mathrm{V}^{6+}, \mathrm{Te}^{2-}, \mathrm{Sn}^{5+}, \mathrm{I}^{-}\) ? Briefly explain your reasoning.

Rank these in order of increasing first ionization energy: \(\mathrm{Zn}, \mathrm{Ca}, \mathrm{Ca}^{2+}, \mathrm{Cl}^{-} .\) Briefly explain your answer.

Criticize these statements. If a statement is incorrect, rewrite it so that it is correct. (a) The energy of a photon is inversely proportional to its frequency. (b) The energy of the hydrogen electron is inversely proportional to its principal quantum number \(n\). (c) Electrons start to enter the fourth energy level as soon as the third level is full. (d) Light emitted by an \(n=4\) to \(n=2\) transition has a lower frequency than that from an \(n=5\) to \(n=2\) transition.

A general chemistry student tells a chemistry classmate that when an electron goes from a \(2 d\) atomic orbital to a \(1 s\) atomic orbital, it emits more energy than that for a \(2 p\) to \(1 s\) transition. The other student is skeptical and says that such an energy change is not possible and explains why. What explanation was given?