Chapter 7

(a) What is meant by the terms acidic oxide and basic oxide? (b) How can we predict whether an oxide will be acidic or basic, based on its composition?

Arrange the following oxides in order of increasing acidity: \(\mathrm{CO}_{2}, \mathrm{CaO}, \mathrm{Al}_{2} \mathrm{O}_{3}, \mathrm{SO}_{3}, \mathrm{SiO}_{2}\), and \(\mathrm{P}_{2} \mathrm{O}_{5}\)

Chlorine reacts with oxygen to form \(\mathrm{Cl}_{2} \mathrm{O}_{7} .\) (a) What is the name of this product (see Table 2.6)? (b) Write a balanced equation for the formation of \(\mathrm{Cl}_{2} \mathrm{O}_{7}(l)\) from the elements. (c) Under usual conditions, \(\mathrm{Cl}_{2} \mathrm{O}_{7}\) is a colorless liquid with a boiling point of \(81^{\circ} \mathrm{C}\). Is this boiling point expected or surprising? (d) Would you expect \(\mathrm{Cl}_{2} \mathrm{O}_{7}\) to be more reactive toward \(\mathrm{H}^{+}(a q)\) or \(\mathrm{OH}^{-}(a q) ?\) Explain.

An element \(X\) reacts with oxygen to form \(\mathrm{XO}_{2}\) and with chlorine to form \(\mathrm{XCl}_{4} \mathrm{XO}_{2}\) is a white solid that melts at high temperatures (above \(\left.1000^{\circ} \mathrm{C}\right)\). Under usual conditions, \(\mathrm{XCl}_{4}\) is a colorless liquid with a boiling point of \(58^{\circ} \mathrm{C}\). (a) \(\mathrm{XCl}_{4}\) reacts with water to form \(\mathrm{XO}_{2}\) and another product. What is the likely identity of the other product? (b) Do you think that element \(X\) is a metal, nonmetal, or metalloid? Explain. (c) By using a sourcebook such as the CRC Handbook of Chemistry and Physics, try to determine the identity of element \(\mathrm{X}\).

Write balanced equations for the following reactions: (a) barium oxide with water, (b) iron(II) oxide with perchloric acid, (c) sulfur trioxide with water, (d) carbon dioxide with aqueous sodium hydroxide.

Write balanced equations for the following reactions: (a) potassium oxide with water, (b) diphosphorus trioxide with water, (c) chromium(III) oxide with dilute hydrochloric acid, (d) selenium dioxide with aqueous potassium hydroxide.

Compare the elements sodium and magnesium with respect to the following properties: (a) electron configuration, (b) most common ionic charge, (c) first ionization energy, (d) reactivity toward water, (e) atomic radius. Account for the differences between the two elements.

(a) Why is calcium generally more reactive than magnesium? (b) Why is calcium generally less reactive than potassium?

(a) Why is cesium more reactive toward water than is lithium? (b) One of the alkali metals reacts with oxygen to form a solid white substance. When this substance is dissolved in water, the solution gives a positive test for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\). When the solution is tested in a burner flame, a lilac-purple flame is produced. What is the likely identity of the metal? (c) Write a balanced chemical equation for reaction of the white substance with water.

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Potassium metal burns in an atmosphere of chlorine gas, (b) Strontium oxide is added to water. (c) A fresh surface of lithium metal is exposed to oxygen gas. (d) Sodium metal is reacted with molten sulfur.

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Cesium is added to water. (b) Stontium is added to water. (c) Sodium reacts with oxygen. (d) Calcium reacts with iodine.

(a) If we arrange the elements of the second period (Li-Ne) in order of increasing first ionization energy, where would hydrogen fit into this series? (b) If we now arrange the elements of the third period (Na-Ar) in order of increasing first ionization energy, where would lithium fit into this series? (c) Are these series consistent with the assignment of hydrogen as a nonmetal and lithium as a metal?

(a) As described in Section \(7.7\), the alkali metals react with hydrogen to form hydrides and react with halogens-for example, fluorine-to form halides. Compare the roles of hydrogen and the halogen in these reactions. In what sense are the forms of hydrogen and halogen in the products alike? (b) Write balanced equations for the reaction of fluorine with calcium and for the reaction of hydrogen with calcium. What are the similarities among the products of these reactions?

Compare the elements fluorine and chlorine with respect to the following properties: (a) electron configuration, (b) most common ionic charge, (c) first ionization energy, (d) reactivity toward water, (e) electron affinity, (f) atomic radius. Account for the differences between the two elements.

Little is known about the properties of astatine, At, because of its rarity and high radioactivity. Nevertheless, it is possible for us to make many predictions about its properties. (a) Do you expect the element to be a gas, liquid, or solid at room temperature? Explain. (b) What is the chemical formula of the compound it forms with Na?

Until the early 1960 s the group 8 A elements were called the inert gases; before that they were called the rare gases. The term rare gases was dropped after it was discovered that argon accounts for roughly \(1 \%\) of Earth's atmosphere. (a) Why was the term inert gases dropped? (b) What discovery triggered this change in name? (c) What name is applied to the group now?

Why does xenon react with fluorine, whereas neon does not?

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Ozone decomposes to dioxygen. (b) Xenon reacts with fluorine. (Write three different equations.) (c) Sulfur reacts with hydrogen gas. (d) Fluorine reacts with water.

Write a balanced equation for the reaction that occurs in each of the following cases: (a) Chlorine reacts with water. (b) Barium metal is heated in an atmosphere of hydrogen gas. (c) Lithium reacts with sulfur. (d) Fluorine reacts with magnesium metal.

Consider the stable elements through lead \((Z=82)\). In how many instances are the atomic weights of the elements in the reverse order relative to the atomic numbers of the elements? What is the explanation for these cases?

(a) Which will have the lower energy, a 4 s or a \(4 p\) electron in an As atom? (b) How can we use the concept of effective nuclear charge to explain your answer to part (a)?

(a) If the core electrons were totally effective at shielding the valence electrons and the valence electrons provided no shielding for each other, what would be the effective nuclear charge acting on the 3 s and \(3 p\) valence electrons in \(P\) ? (b) Repeat these calculations using Slater's rules. (c) Detailed calculations indicate that the effective nuclear charge is \(5.6+\) for the 3 s electrons and \(4.9+\) for the \(3 p\) electrons. Why are the values for the 3 s and \(3 p\) electrons different? (d) If you remove a single electron from a \(\mathrm{P}\) atom, which orbital will it come from? Explain.

Nearly all the mass of an atom is in the nucleus, which has a very small radius. When atoms bond together (for example, two fluorine atoms in \(\mathrm{F}_{2}\) ), why is the distance separating the nuclei so much larger than the radii of the nuclei?

Consider the change in effective nuclear charge experienced by a \(2 p\) electron as we proceed from \(C\) to \(N\). (a) Based on a simple model in which core electrons screen the valence electrons completely and valence electrons do not screen other valence electrons, what do you predict for the change in \(Z_{\text {eff }}\) from \(C\) to \(N\) ? (b) What change do you predict using Slater's rules? (c) The actual change in \(Z_{\text {eff }}\) from \(C\) to \(N\) is \(0.70+.\) Which approach to estimating \(Z_{\text {eff }}\) is more accurate? (d) The change in \(Z_{\text {dif }}\) from \(N\) to \(O\) is smaller than that from \(C\) to \(N\). Can you provide an explanation for this observation?

As we move across a period of the periodic table, why do the sizes of the transition elements change more gradually than those of the representative elements?

It is possible to produce compounds of the form \(\mathrm{GeClH}_{3}, \mathrm{GeCl}_{2} \mathrm{H}_{2}\), and \(\mathrm{GeCl}_{3} \mathrm{H}\). What values do you predict for the Ge \(-\mathrm{H}\) and \(\mathrm{Ge}-\mathrm{Cl}\) bond lengths in these compounds?

Note from the following table that the increase in atomic radius in moving from \(Z r\) to \(H f\) is smaller than in moving from \(Y\) to La. Suggest an explanation for this effect. \begin{tabular}{llll} \hline \multicolumn{3}{l} { Atomic Radii \((\AA)\)} \\ \hline \(\mathrm{Sc}\) & \(1.44\) & \(\mathrm{Ti}\) & \(1.36\) \\ \(\mathrm{Y}\) & \(1.62\) & \(\mathrm{Zr}\) & \(1.48\) \\ \(\mathrm{La}\) & \(1.69\) & \(\mathrm{Hf}\) & \(1.50\) \\ \hline \end{tabular}

Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons. Why is it extremely difficult to add a third electron to the atom?

Use electron configurations to explain the following observations: (a) The first ionization energy of phosphorus is greater than that of sulfur. (b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen. (c) The second ionization energy of oxygen is greater than the first ionization energy of fluorine. (d) The third ionization energy of manganese is. greater than those of both chromium and iron.

The following table gives the electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group \(1 \mathrm{~B}\) and group \(2 \mathrm{~B}\) metals: (a) Why are the electron affinities of the group \(2 \mathrm{~B}\) elements greater than zero? (b) Why do the electron affinities of the group \(1 \mathrm{~B}\) elements become more negative as we move down the group? [Hint: Examine the trends in the electron affinity of other groups as we proceed down the periodic table.] \begin{tabular}{|c|l|} \hline \(\mathrm{Cu}\) & \multirow{2}{*} { \(\mathrm{Zn}\) \(-119\)} & \(>0\) \\ \hline \(\mathrm{Ag}\) & \(\mathrm{Cd}\) \\ \(-126\) & \(>0\) \\ \hline \(\mathrm{Au}\) & \(\mathrm{Hg}\) \\ \(-223\) & \(>0\) \\ \hline \end{tabular}

Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like a nonmetal. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity. (a) Explain why the electron affinity of hydrogen is much closer to the values for the alkali elements than for the halogens. (b) Is the following statement true? "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds." If not, correct it. If it is, explain in terms of electron configurations. (c) Explain why the ionization energy of hydrogen is closer to the values for the halogens than for the alkali metals.

The first ionization energy of the oxygen molecule is the energy required for the following process: $$ \mathrm{O}_{2}(g) \longrightarrow \mathrm{O}_{2}^{+}(g)+\mathrm{e} $$ The energy needed for this process is \(1175 \mathrm{~kJ} / \mathrm{mol}\), very similar to the first ionization energy of Xe. Would you expect \(\mathrm{O}_{2}\) to react with \(\mathrm{F}_{2}\) ? If so, suggest a product or products of this reaction.

The element strontium is used in a variety of industrial processes. It is not an extremely hazardous substance, but low levels of strontium ingestion could affect the health of children. Radioactive strontium is very hazardous, it was a by-product of nuclear weapons testing and was found widely distributed following nuclear tests. Calcium is quite common in the environment, including food products, and is frequently present in drinking water. Discuss the similarities and differences between calcium and strontium, and indicate how and why strontium might be expected to accompany calcium in water supplies, uptake by plants, and so on.

There are certain similarities in properties that exist between the first member of any periodic family and the element located below it and to the right in the periodic table. For example, in some ways Li resembles \(\mathrm{Mg}\), Be resembles \(\mathrm{Al}\), and so forth. This observation is called the diagonal relationship. Using what we have learned in this chapter, offer a possible explanation for this relationship.

A historian discovers a nineteenth-century notebook in which some observations, dated 1822 , on a substance thought to be a new element, were recorded. Here are some of the data recorded in the notebook: Ductile, silverwhite, metallic looking. Softer than lead Unaffected by water. Stable in air. Melting point: \(153^{\circ} \mathrm{C}\) Density: \(7.3 \mathrm{~g} / \mathrm{cm}^{3} .\) Electrical conductivity: \(20 \%\) that of copper. Hardness. About \(1 \%\) as hard as iron. When \(4.20 \mathrm{~g}\) of the unknown is heated in an excess of oxygen, \(5.08 \mathrm{~g}\) of a white solid is formed. The solid could be sublimed by heating to over \(800^{\circ} \mathrm{C}\). (a) Using information in the text and a handbook of chemistry, and making allowances for possible variations in numbers from current values, identify the element reported. (b) Write a balanced chemical equation for the reaction with oxygen. (c) Judging from Figure 7.2, might this nineteenth- century investigator have been the first to discover a new element?

Moseley established the concept of atomic number by studying X-rays emitted by the elements. The X-rays emitted by some of the elements have the following wavelengths: \begin{tabular}{ll} \hline Element & Wavelength (A) \\ \hline \(\mathrm{Ne}\) & \(14.610\) \\ \(\mathrm{Ca}\) & \(3.358\) \\ \(\mathrm{Zn}\) & \(1.435\) \\ \(\mathrm{Zr}\) & \(0.786\) \\ \(\mathrm{Sn}\) & \(0.491\) \\ \hline \end{tabular} (a) Calculate the frequency, \(\nu\), of the \(X\) -rays emitted by each of the elements, in Hz. (b) Using graph paper (or

(a) Write the electron configuration for \(\mathrm{Li}\), and estimate the effective nuclear charge experienced by the valence electron. (b) The energy of an electron in a one-electron atom or ion equals \(\left(-2.18 \times 10^{-18} \mathrm{~J}\right)\left(\frac{\mathrm{Z}^{2}}{n^{2}}\right)\) where \(\mathrm{Z}\) is the nuclear charge and \(n\) is the principal quantum number of the electron. Estimate the first ionization energy of Li. (c) Compare the result of your calculation with the value reported in table \(7.4\), and explain the difference. (d) What value of the effective nuclear charge gives the proper value for the ionization energy? Does this agree with your explanation in (c)?

One way to measure ionization energies is photoelectron spectroscopy (PES), a technique based on the photoelectric effect. em (Section 6.2) In PES, monochromatic light is directed onto a sample, causing electrons to be emitted. The kinetic energy of the emitted electrons is measured. The difference between the energy of the photons and the kinetic energy of the electrons corresponds to the energy needed to remove the electrons (that is, the ionization energy). Suppose that a PES experiment is performed in which mercury vapor is irradiated with ultraviolet light of wavelength \(58.4 \mathrm{~nm}\). (a) What is the energy of a photon of this light, in \(\mathrm{eV}\) ? (b) Write an equation that shows the process corresponding to the first ionization energy of \(\mathrm{Hg}\). (c) The kinetic energy of the emitted electrons is measured to be \(10.75 \mathrm{eV}\). What is the first ionization energy of \(\mathrm{Hg}\), in \(\mathrm{kJ} / \mathrm{mol} ?\) (d) With reference to Figure \(7.11\), determine which of the halogen elements has a first ionization energy closest to that of mercury.

Consider the gas-phase transfer of an electron from a sodium atom to a chlorine atom: $$ \mathrm{Na}(\mathrm{g})+\mathrm{Cl}(\mathrm{g}) \longrightarrow \mathrm{Na}^{+}(\mathrm{g})+\mathrm{Cl}^{-}(g) $$ (a) Write this reaction as the sum of two reactions, one that relates to an ionization energy and one that relates to an electron affinity. (b) Use the result from part (a), data in this chapter, and Hess's law to calculate the enthalpy of the above reaction. Is the reaction exothermic or endothermic? (c) The reaction between sodium metal and chlorine gas is highly exothermic and produces \(\mathrm{NaCl}(\mathrm{s})\), whose structure was discussed in Section 2.7. Comment on this observation relative to the calculated enthalpy for the aforementioned gas-phase reaction.

When magnesium metal is burned in air (Figure 3.5), two products are produced. One is magnesium oxide, \(\mathrm{MgO}\). The other is the product of the reaction of \(\mathrm{Mg}\) with molecular nitrogen, magnesium nitride. When water is added to magnesium nitride, it reacts to form magnesium oxide and ammonia gas. (a) Based on the charge of the nitride ion (Table 2.5), predict the formula of magnesium nitride. (b) Write a balanced equation for the reaction of magnesium nitride with water. What is the driving force for this reaction? (c) In an experiment a piece of magnesium ribbon is burned in air in a crucible. The mass of the mixture of \(\mathrm{MgO}\) and magnesium nitride after burning is \(0.470 \mathrm{~g}\). Water is added to the crucible, further reaction occurs, and the crucible is heated to dryness until the final product is \(0.486 \mathrm{~g}\) of \(\mathrm{MgO}\). What was the mass percentage of magnesium nitride in the mixture obtained after the initial burning? (d) Magnesium nitride can also be formed by reaction of the metal with ammonia at high temperature. Write a balanced equation for this reaction. If a \(6.3-\mathrm{g} \mathrm{Mg}\) ribbon reacts with \(2.57 \mathrm{~g} \mathrm{NH}_{3}(\mathrm{~g})\) and the reaction goes to \(\mathrm{com}\) pletion, which component is the limiting reactant? What mass of \(\mathrm{H}_{2}(\mathrm{~g})\) is formed in the reaction? (e) The standard enthalpy of formation of solid magnesium nitride is \(-461.08 \mathrm{~kJ} / \mathrm{mol}\). Calculate the standard enthalpy change for the reaction between magnesium metal and ammonia gas.

Potassium superoxide, \(\mathrm{KO}_{2}\), is often used in oxygen masks (such as those used by firefighters) because \(\mathrm{KO}_{2}\) reacts with \(\mathrm{CO}_{2}\) to release molecular oxygen. Experiments indicate that \(2 \mathrm{~mol}\) of \(\mathrm{KO}_{2}(\mathrm{~s})\) react with each mole of \(\mathrm{CO}_{2}(g) .\) (a) The products of the reaction are \(\mathrm{K}_{2} \mathrm{CO}_{3}(s)\) and \(\mathrm{O}_{2}(g) .\) Write a balanced equation for the reaction between \(\mathrm{KO}_{2}(\mathrm{~s})\) and \(\mathrm{CO}_{2}(\mathrm{~g}) .\) (b) Indicate the oxidation number for each atom involved in the reaction in part (a). What elements are being oxidized and reduced? (c) What mass of \(\mathrm{KO}_{2}(s)\) is needed to consume \(18.0 \mathrm{~g}\) \(\mathrm{CO}_{2}(g) ?\) What mass of \(\mathrm{O}_{2}(g)\) is produced during this reaction?