Chapter 9

The following melting points are in degrees Celsius. Show that melting point is a periodic property of these elements: \(\mathrm{Al}, 660 ; \mathrm{Ar},-189 ; \mathrm{Be}, 1278 ; \mathrm{B}, 2300 ; \mathrm{C}\) \(3350 ; \mathrm{Cl},-101 ; \mathrm{F},-220 ; \mathrm{Li}, 179 ; \mathrm{Mg}, 651 ; \mathrm{Ne},-249 ; \mathrm{N}\) \(-210 ; \mathrm{O},-218 ; \mathrm{P}, 590 ; \mathrm{Si}, 1410 ; \mathrm{Na}, 98 ; \mathrm{S}, 119.\)

Mendeleev's periodic table did not preclude the possibility of a new group of elements that would fit within the existing table, as was the case with the noble gases. Moseley's work did preclude this possibility. Explain this difference.

Explain why the several periods in the periodic table do not all have the same number of members.

Assuming that the seventh period is 32 members long, what should be the atomic number of the noble gas following radon (Rn)? Of the alkali metal following francium (Fr)? What would you expect their approximate atomic masses to be?

Concerning the incomplete seventh period of the periodic table, what should be the atomic number of the element (a) for which the filling of the \(6 d\) subshell is completed; (b) that should most closely resemble bismuth; (c) that should be a noble gas?

For each of the following pairs, indicate the atom that has the larger size: (a) Te or Br; (b) \(\mathrm{K}\) or \(\mathrm{Ca} ;\) (c) Ca or Cs; (d) \(\mathrm{N}\) or \(\mathrm{O} ;\) (e) \(\mathrm{O}\) or \(\mathrm{P} ;\) (f) \(\mathrm{Al}\) or \(\mathrm{Au}.\)

Indicate the smallest and the largest species (atom or ion) in the following group: Al atom, F atom, As atom, \(\mathrm{Cs}^{+}\) ion, \(\mathrm{I}^{-}\) ion, \(\mathrm{N}\) atom.

Explain why the radii of atoms do not simply increase uniformly with increasing atomic number.

The masses of individual atoms can be determined with great precision, yet there is considerable uncertainty about the exact size of an atom. Explain why this is the case.

Which is (a) the smallest atom in group \(13 ;\) (b) the smallest of the following atoms: Te, In, Sr, Po, Sb? Why?

How would you expect the sizes of the hydrogen ion, \(\mathrm{H}^{+},\) and the hydride ion, \(\mathrm{H}^{-},\) to compare with that of the H atom and the He atom? Explain.

Arrange the following in expected order of increasing radius: \(\mathrm{Br}, \mathrm{Li}^{+}, \mathrm{Se}, \mathrm{I}^{-} .\) Explain your answer.

Among the following ions, several pairs are isoelectronic. Identify these pairs. \(\mathrm{Fe}^{2+}, \mathrm{Sc}^{3+}, \mathrm{Ca}^{2+}, \mathrm{F}^{-}\) \(\mathrm{Co}^{2+}, \mathrm{Co}^{3+}, \mathrm{Sr}^{2+}, \mathrm{Cu}^{+}, \mathrm{Zn}^{2+}, \mathrm{Al}^{3+}.\)

The following species are isoelectronic with the noble gas krypton. Arrange them in order of increasing radius and comment on the principles involved in doing so: \(\mathrm{Rb}^{+}, \mathrm{Y}^{3+}, \mathrm{Br}^{-}, \mathrm{Sr}^{2+}, \mathrm{Se}^{2-}.\)

All the isoelectronic species illustrated in the text had the electron configurations of noble gases. Can two ions be isoelectronic without having noble-gas electron configurations? Explain.

Is it possible for two different atoms to be isoelectronic? two different cations? two different anions? a cation and an anion? Explain.

Use principles established in this chapter to arrange the following atoms in order of increasing value of the first ionization energy: \(\mathrm{Sr}, \mathrm{Cs}, \mathrm{S}, \mathrm{F}, \mathrm{As}.\)

Are there any atoms for which the second ionization energy \(\left(I_{2}\right)\) is smaller than the first \(\left(I_{1}\right) ?\) Explain.

Some electron affinities are negative quantities, and some are zero or positive. Why is this not also the case with ionization energies?

How much energy, in joules, must be absorbed to convert to \(\mathrm{Na}^{+}\) all the atoms present in \(1.00 \mathrm{mg}\) of gaseous Na? The first ionization energy of Na is \(495.8 \mathrm{kJ} / \mathrm{mol}.\)

How much energy, in kilojoules, is required to remove all the third-shell electrons in a mole of gaseous silicon atoms?

Use ionization energies and electron affinities listed in the text to determine whether the following reaction is endothermic or exothermic. $$ \mathrm{Mg}(\mathrm{g})+2 \mathrm{F}(\mathrm{g}) \longrightarrow \mathrm{Mg}^{2+}(\mathrm{g})+2 \mathrm{F}^{-}(\mathrm{g}) $$

The \(\mathrm{Na}^{+}\) ion and the Ne atom are isoelectronic. The ease of loss of an electron by a gaseous Ne atom, \(I_{1}\) has a value of \(2081 \mathrm{kJ} / \mathrm{mol} .\) The ease of loss of an electron from a gaseous \(\mathrm{Na}^{+}\) ion, \(I_{2}\), has a value of \(4562 \mathrm{kJ} / \mathrm{mol} .\) Why are these values not the same?

Compare the elements \(\mathrm{Al}, \mathrm{Si}, \mathrm{S},\) and \(\mathrm{Cl}.\) (a) Place the elements in order of increasing ionization energy. (b) Place the elements in order of increasing electron affinity.

Compare the elements \(\mathrm{Na}, \mathrm{Mg}, \mathrm{O},\) and \(\mathrm{P}.\) (a) Place the elements in order of increasing ionization energy. (b) Place the elements in order of increasing electron affinity.

Unpaired electrons are found in only one of the following species. Indicate which one, and explain why: \(\mathrm{F}^{-}, \mathrm{Ca}^{2+}, \mathrm{Fe}^{2+}, \mathrm{S}^{2-}\)

Which of the following species has the greatest number of unpaired electrons (a) \(\mathrm{Ge} ;\) (b) \(\mathrm{Cl} ;\) (c) \(\mathrm{Cr}^{3+}\) (d) Br -?

Which of the following species would you expect tobe diamagnetic and which paramagnetic? (a) \(\mathrm{K}^{+}=\) (b) \(\mathrm{Cr}^{3+} ;\) (c) \(\mathrm{Zn}^{2+} ;\) (d) \(\mathrm{Cd} ;\) (e) \(\mathrm{Co}^{3+} ;\) (f) \(\mathrm{Sn}^{2+} ;\) (g) Br.

Write electron configurations consistent with the following data on numbers of unpaired electrons: \(\mathrm{Ni}^{2+}, 2 ; \mathrm{Cu}^{2+}, 1 ; \mathrm{Cr}^{3+}, 3.\)

Must all atoms with an odd atomic number be paramagnetic? Must all atoms with an even atomic number be diamagnetic? Explain.

Neither \(\mathrm{Co}^{2+}\) nor \(\mathrm{Co}^{3+}\) has \(4 \mathrm{s}\) electrons in its electron configuration. How many unpaired electrons would you expect to find in each of these ions? Explain.

For the following groups of elements, select the one that has the property noted: (a) the largest atom: \(\mathrm{Mg}, \mathrm{Mn}, \mathrm{Mo}, \mathrm{Ba}, \mathrm{Bi}, \mathrm{Br}.\) (b) the lowest first ionization energy: \(\mathrm{B}, \mathrm{Sr}, \mathrm{Al}, \mathrm{Br}\) \(\mathrm{Mg}_{\ell} \mathrm{Pb}.\) (c) the most negative electron affinity: \(\mathrm{As}, \mathrm{B}, \mathrm{Cl}\) \(\mathrm{K}, \mathrm{Mg}, \mathrm{S}.\) (d) the largest number of unpaired electrons: \(\mathrm{F}, \mathrm{N}, \mathrm{S}^{2-}, \mathrm{Mg}^{2+}, \mathrm{Sc}^{3+}, \mathrm{Ti}^{3+}.\)

Which of the following ions are unlikely to be found in chemical compounds: \(\mathrm{K}^{+}, \mathrm{Ga}^{4+}, \mathrm{Fe}^{6+} \mathrm{S}^{2-}, \mathrm{Ge}^{5+},\) or \(\mathrm{Br}^{-} ?\) Explain briefly.

Which of the following ions are likely to be found in chemical compounds: \(\mathrm{Na}^{2+}, \mathrm{Li}^{+}, \mathrm{Al}^{4+}, \mathrm{F}^{2-},\) or \(\mathrm{Te}^{2-} ?\) Explain briefly.

Complete and balance the following equations. If no reaction occurs, so state. (a) \(\operatorname{Rb}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\) (b) \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Na}^{+}(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq}) \longrightarrow\) (c) \(\operatorname{SrO}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\) (d) \(\mathrm{SO}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\)

Write balanced equations to represent (a) the displacement of a halide anion from aqueous solution by liquid bromine (b) the reaction with water of an alkali metal with \(Z>50\) (c) the reaction of tetraphosphorus decoxide with water (d) the reaction of aluminum oxide with aqueous sulfuric acid

Sketch a periodic table that would include all the elements in the main body of the table. How many "numbers" wide would the table be?

Listed below are two atomic properties of the element germanium. Refer only to the periodic table on the inside front cover and indicate probable values for each of the following elements, expressed as greater than, about equal to, or less than the value for Ge. $$\begin{array}{lcc} \hline \text { Element } & \text { Atomic Radius } & \begin{array}{c} \text { First lonization } \\ \text { Energy } \end{array} \\ \hline \mathrm{Ge} & 122 \mathrm{pm} & 762 \mathrm{kJ} / \mathrm{mol} \\ \mathrm{Al} & ? & ? \\ \mathrm{In} & ? & ? \\ \mathrm{Se} & ? & ? \\ \hline \end{array}$$

In estimating the boiling point and melting point of bromine in Example \(9-5,\) could we have used Celsius or Fahrenheit instead of Kelvin temperature? Explain.

Two elements, \(A\) and \(B\), have the electron configurations shown. $$ \mathrm{A}=\left[\begin{array}{ll} \mathrm{Ar} & 4 s^{1} \end{array} \quad \mathrm{B}=\left[\begin{array}{l} \mathrm{Ar} \end{array}\right] 3 d^{10} 4 \mathrm{s}^{2} 4 p^{3}\right. $$ (a) Which element is a metal? (b) Which element has the greater ionization energy? (c) Which element has the larger atomic radius? (d) Which element has the greater electron affinity?

Two elements, \(A\) and \(B\), have the electron configurations shown. $$ \mathrm{A}=[\mathrm{Kr}] 4 s^{2} \quad \mathrm{B}=[\mathrm{Ar}] 3 d^{10} 4 s^{2} 4 p^{5} $$ (a) Which element is a metal? (b) Which element has the greater ionization energy? (c) Which element has the larger atomic radius? (d) Which element has the greater electron affinity?

Refer only to the periodic table on the inside front cover, and arrange the following ionization energies in order of increasing value: \(I_{1}\) for \(\mathrm{F} ; I_{2}\) for \(\mathrm{Ba}\) \(I_{3}\) for \(\mathrm{Sc} ; I_{2}\) for \(\mathrm{Na} ; I_{3}\) for \(\mathrm{Mg}\). Explain the basis of any uncertainties.

Plot a graph of the square roots of the ionization energies versus the nuclear charge for the two series \(\mathrm{Li}, \mathrm{Be}^{+}, \mathrm{B}^{2+}, \mathrm{C}^{3+},\) and \(\mathrm{Na}, \mathrm{Mg}^{+}, \mathrm{Al}^{2+}, \mathrm{Si}^{3+} .\) Explain the observed relationship with the aid of Bohr's expression for the binding energy of an electron in a one electron atom.

A method for estimating electron affinities is to extrapolate \(Z_{\text {eff }}\) values for atoms and ions that contain the same number of electrons as the negative ion of interest. Use the data in the table on the next page to answer the questions that follow. $$\begin{array}{lll} \hline \begin{array}{l} \text { Atom or lon: } \\ \text { I(kJmol }^{-1} \text {) } \end{array} & \begin{array}{l} \text { Atom or lon: } \\ \text { I(kJmol }^{-1} \text {) } \end{array} & \begin{array}{l} \text { Atom or lon: } \\ \text { I(kJmol }^{-1} \text {) } \end{array} \\ \hline \text { Ne: 2080 } & \text { F: 1681 } & \text { O: } 1314 \\ \text { Na }^{+}: 4565 & \text { Ne }^{+}: 3963 & \text { F }^{+}: 3375 \\ \text { Mg }^{2+} \text { : 7732 } & \text { Na }^{2+}: 6912 & \text { Ne }^{2+}: 6276 \\ \text { A1 }^{\text {3 }^{+}: 11,577} & \text { Mg }^{3+}: 10,548 & \text { Na }^{3+}: 9540 \\ \hline \end{array}$$ (a) Estimate the electron affinity of \(F\), and compare it with the experimental value. (b) Estimate the electron affinities of \(\mathrm{O}\) and \(\mathrm{N}\) (c) Examine your results in terms of penetration and screening.

In your own words, define the following terms: (a) isoelectronic; (b) valence- shell electrons; (c) metal; (d) nonmetal; (e) metalloid.

Briefly describe each of the following ideas or phenomena: (a) the periodic law; (b) ionization energy; (c) electron affinity; (d) paramagnetism.

Explain the important distinctions between each pair of terms: (a) actinide and lanthanide element; (b) covalent and metallic radius; (c) atomic number and effective nuclear charge; (d) ionization energy and electron affinity; (e) paramagnetic and diamagnetic.

The element whose atoms have the electron configuration \([\mathrm{Kr}] 4 d^{10} 5 \mathrm{s}^{2} 5 p^{3}(\mathrm{a})\) is in group 13 of the periodic table; (b) bears a similarity to the element Bi; (c) is similar to the element \(\mathrm{Te} ;\) (d) is a transition element.

The fourth-period element with the largest atom is (a) \(\mathrm{K} ;\) (b) \(\mathrm{Br} ;\) (c) \(\mathrm{Pb} ;\) (d) \(\mathrm{Kr}\).

The largest of the following is (a) an Ar atom; (b) a \(\mathrm{K}^{+}\) ion; \((c) a C a^{2+}\) ion; \((d) a C l^{-}\) ion.