Until the early 1960s the group 8A elements were called the inert gases; before that

We can draw an analogy between the attraction of an electron to a nucleus and seeing a lightbulb-in essence, the more nuclear charge the electron "sees," the greater the attraction. (a) Within this analogy, discuss how the shielding by core electrons is analogous to putting a frosted-glass lampshade between the Lightbulb and your eyes, as shown in the illustration. (b) Explain how we could mimic moving to the right in a row of the periodic table by changing the wattage of the lightbulb. (c) How would you change the wattage of the bulb and/ or the frosted glass to mimic the effect of moving down a column of the periodic table? [Section 7.2]

Fluorine has atomic number 9. If we represent the radius of a fluorine atom with the billiard ball illustrated here, would the analogy be more appropriate for the bonding or nonbonding atomic radius? If we used the same billiard ball to illustrate the concept of fluorine's bonding atomic radius, would we overestimate or underestimate the bonding atomic radius? Explain. [Section 7.3]

onsider the A2X4 molecule depicted below, where A and X are elements. The A -A bond length in this molecule is d1, and the four A -X bond lengths are each dz. (a) In terms of d1 and dz, how could you define the bonding atomic radii of atoms A and X? (b) In terms of d1 and dz, what would you predict for the X -X bond length of an X2 molecule? [Section 7.3]

Make a simple sketch of the shape of the main part of the periodic table, as shown. (a) Ignoring H and He, write a single straight arrow from the element with the smallest bonding atomic radius to the element with the largest. (b) Ignoring H and He, write a single straight arrow from the element with the smallest first ionization energy to the element with the largest. (c) What significant observation can you make from the arrows you drew in parts (a) and (b)? [Sections 7.3 and 7.4]

In the chemical process called electron transfer, an electron is transferred from one atom or molecule to another (We will talk about electron transfer extensively in Chapter 20.) A simple electron transfer reaction is A(g) + A(g) ----+ A+(g) + A-(g) In terms of the ionization energy and electron affinity of atom A, what is the energy change for this reaction? For a representative nonmetal such as chlorine, is this process exothermic? For a representative metal such as sodium, is this process exothermic? [Sections 7.4 and 7.5]

An element X reacts with F2(g) to form the molecular product shown below. (a) Write a balanced equation for this reaction (do not worry about the phases for X and the product). (b) Do you think that X is a metal or nonmetal? Explain. [Section 7.6]

Why did Mendeleev leave blanks in his early version of the periodic table? How did he predict the properties of the elements that belonged in those blanks?

The prefix eka- comes from the Sanskrit word for one. Mendeleev used this prefix to indicate that the unknown element was one place away from the known element that followed the prefix. For example, eka-silicon, which we now call germanium, is one element below silicon. Mendeleev also predicted the existence of ekamanganese, which was not experimentally confirmed until 1937 because this element is radioactive and does not occur in nature. Based on the periodic table shown in Figure 7.2, what do we now call the element Mendeleev called eka-manganese?

In Chapter 1 we learned that silicon is the second most abundant element in Earth's crust, accounting for more than one-fourth of the mass of the crust (Figure 1.6). Yet we see that silicon is not among the elements that have been known since ancient times (Figure 7.2), whereas iron, which accounts for less than 5% of Earth's crust, has been known since prehistoric times. Given silicon's abundance how do you account for its relatively late discovery?

(a) What is meant by the term effective nuclear charge? (b) How does the effective nuclear charge experienced by the valence electrons of an atom vary going from left to right across a period of the periodic table?

(a) How is the concept of effective nuclear charge used to simplify the numerous electron-electron repulsions in a many-electron atom? (b) Which experiences a greater effective nuclear charge in a Be atom, the 1s electrons or the 2s electrons? Explain

Detailed calculations show that the value of Zeff for Na and K atoms is 2.51 + and 3.49+, respectively. (a) What value do you estimate for Zeff experienced by the outermost electron in both Na and K by assuming core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant? (b) What values do you estimate for Zeff using Slater's rules? (c) Which approach gives a more accurate estimate of Zeff? (d) Does either method of approximation account for the gradual increase in Zeff that occurs upon moving down a group?

Detailed calculations show that the value of Zeff for Si and Cl atoms is 4.29+ and 6.12+, respectively. (a) What value do you estimate for Zeff experienced by the outermost electron in both Si and Cl by assuming core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant? (b) What values do you estimate for Zeff using Slater's rules? (c) Which approach gives a more accurate estimate of Zeff? (d) Which method of approximation more accurately accounts for the steady increase in Zeff that occurs upon moving left to right across a period?

Which will experience the greater effective nuclear charge, the electrons in the n = 3 shell in Ar or the n = 3 shell in Kr? Which will be closer to the nucleus? Explain.

Arrange the following atoms in order of increasing effective nuclear charge experienced by the electrons in the n = 3 electron shell: K, Mg, P, Rh, and Ti. Explain the basis for your order.

(a) Because an exact outer boundary cannot be measured or even calculated for an atom, how are atomic radii determined? (b) What is the difference between a bonding radius and a nonbonding radius? (c) For a given element, which one is larger?

a) Why does the quantum mechanical description of many-electron atoms make it difficult to define a precise atomic radius? (b) When nonbonded atoms come up against one another, what determines how closely the nuclear centers can approach?

The distance between W atoms in tungsten metal is 2.74 A. What is the atomic radius of a tungsten atom in this environment? (This radius is called the metallic radius.)

Estimate the As-I bond length from the data in Figure 7.7, and compare your value to the experimental As-I bond length in arsenic triiodide, Asi3, 2.55 A.

The experimental Bi-I bond length in bismuth triiodide, Bil3, is 2.81 A. Based on this value and data in Figure 7.7, predict the atomic radius of Bi.

How do the sizes of atoms change as we move (a) from left to right across a row in the periodic table, (b) from top to bottom in a group in the periodic table? (c) Arrange the following atoms in order of increasing atomic radius: F, P, S, As.

(a) Among the nonmetallic elements, the change in atomic radius in moving one place left or right in a row is smaller than the change in moving one row up or down. Explain these observations. (b) Arrange the following atoms in order of increasing atomic radius: Si, Al, Ge, Ga.

Using only the periodic table, arrange each set of atoms in order of increasing radius: (a) Ca, Mg, Be; (b) Ga, Br, Ge; (c) AI, Tl, Si.

Using only the periodic table, arrange each set of atoms in order of increasing radius: (a) Ba, Ca, Na; (b) Sn, Sb, As; (c) AI, Be, Si.

(a) Why are monatomic cations smaller than their corresponding neutral atoms? (b) Why are monatomic anions larger than their corresponding neutral atoms? (c) Why does the size of ions increase as one proceeds down a column in the periodic table?

Explain the following variations in atomic or ionic radii: (a) l- > I> I+, (b)Ca2+ > Mg2+ > Be2+, (c) Fe > Fe2+ > Fe3+

Consider a reaction represented by the following spheres: Reactants Products [ij] Wltich sphere represents a metal and which a nonmetal? Explain.

(a) What is an isoelectronic series? (b) Which neutral atom is isoelectronic with each of the following ions: At 3+, Ti4+, Br-, Sn2+.

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. If such an atom does not exist explain why: (a) CC (b) Sc3+, (c) Fe2+, (d) zn2+ , (e) Sn4+

Consider the isoelectronic ions F- and Na+ (a) Which ion is smaller? (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. (c) Repeat this calculation using Slater's rules to estimate the screening constant, S. (d) For isoelectronic ions, how are effective nuclear charge and ionic radius related?

Consider the isoelectronic ions o- and K + (a) Which ion is smaller? (b) Use 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. (d) For isoelectronic ions how are effective nuclear charge and ionic radius related?

Consider S, Cl, and K and their most common ions. (a) List the atoms in order of increasing size. (b) List the ions in order of increasing size. (c) Explain any differences in the orders of the atomic and ionic sizes.

For each of the following sets of atoms and ions, arrange the members in order of increasing size: (a) Se2-, Te2-, Se; (b) Co3+ , Fe2+ , Fe3+; (c) Ca, Ti4+, Sc3+; (d) Be2+ , Na+, Ne.

For each of the following statements, provide an explanation: (a) 02 - is larger than 0; (b) s 2- is larger than 02-; (c) S 2 - is larger than K + ; (d) K+ is larger than Ca2+

1n the ionic compounds LiF, NaCl, KBr, and Rbi, the measured cation-anion distances are 2.01 A (Li-F), 2.82 A (Na-Cl), 3.30 A (K-Br), and 3.67 A (Rb-I), respectively. (a) Predict the cation-anion distance using the values of ionic radii given in Figure 7.8. (b) Is the agreement between the prediction and the experiment perfect? If not, why not? (c) What estimates of the cation-anion distance would you obtain for these four compounds using 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 boron atom.

(a) Why are ionization energies always positive quantities? (b) Why does F have a larger first ionization energy than 0? (c) Why is the second ionization energy of an atom always greater than its first ionization energy?

(a) Why does Li have a larger first ionization energy than N a? (b) The difference between the third and fourth ionization energies of scandium is much larger than the difference between the third and fourth ionization energies of titanium. Why? (c) Why does Li have a much larger second ionization energy than Be?

(a) What is the general relationship between the size of an atom and its first ionization energy? (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 7 A 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?

Based on their positions in the periodic table, predict which atom of the following pairs will have the larger first ionization energy: (a) Cl, Ar; (b) Be, Ca; (c) K, Co; (d) S, Ge; (e) Sn, Te.

For each of the following pairs, indicate which element has the larger first ionization energy: (a) Ti, Ba; (b) Ag, Cu; (c) Ge, Cl; (d) Pb, Sb. (In each case use electron configuration and effective nuclear charge to explain your answer.)

Write the electron configurations for the following ions: (a) !n 3 +, (b) Sb3 +, (c) Te2 -, (d) Te6 +, (e) Hg2 +, (f) Rh3 +

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) Cr3+ , (b) N3 -, (c) Sc3+ , (d) Cu2+, (e) Tl+, (f) Au+

Write the electron configuration for (a) the Ni2 + ion and (b) the Sn2+ ion. How many unpaired electrons does each contain?

The first ionization energy of Ar and the electron affinity of Ar are both positive values. What is the significance of the positive value in each case?

The electron affinity of lithium is a negative value, whereas the electron affinity of beryllium is a positive value. Use electron configurations to account for this observation.

While the electron affinity of bromine is a negative quantity, it is positive for Kr. Use the electron configurations of the two elements to explain the difference.

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?

Consider the first ionization energy of neon and the electron affinity of fluorine. (a) Write equations, including electron configurations, for each process. (b) These two quantities will 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? Explain your answer.

Write an equation for the process that corresponds to the electron affinity of the Mg + ion. Also write the electron configurations of the species involved. What is the magnitude of the energy change in the process? [Hint: The answer is in Table 7.2.]

How are metallic character and first ionization energy related?

Arrange the following pure solid elements in order of increasing electrical conductivity: Ge, Ca, S, and Si. Explain the reasoning you used.

If we look at groups 3A through SA, we see two metalloids for groups 4A (Si, Ge) and SA (As, Sb), but only one metalloid in group 3A (B). To maintain a regular geometric pattern one might expect that aluminum would also be a metalloid, giving group 3A two metalloids. What can you say about the metallic character of aluminum with respect to its neighbors based on its first ionization energy?

Predict whether each of the following oxides is ionic or molecular: SOz, MgO, LizO, PzOs, Y z03, NzO, and Xe03. Explain the reasons for your choices.

Some metal oxides, such as Sc20:y do not react with pure water, but they do react when the solution becomes either acidic or basic. Do you expect Sc203 to react when the solution becomes acidic or when it becomes basic? Write a balanced chemical equation to support your answer.

(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: COz, CaO, AlzQ3, S03, SiOz, and PzOs.

Chlorine reacts with oxygen to form Cl207. (a) What is the name of this product (see Table 2.6)? (b) Write a balanced equation for the formation of Cl207(1) from the elements. (c) Under usual conditions, Cl207 is a colorless liquid with a boiling point of 81 C. Is this boiling point expected or surprising? (d) Would you expect Cl207 to be more reactive toward H+(aq) or OH-(aq)? Explain.

An element X reacts with oxygen to form X02 and with chlorine to form XC4. X02 is a white solid that melts at high temperatures (above 1000 C). Under usual conditions, XC4 is a colorless liquid with a boiling point of 58 C. (a) XC4 reacts with water to form X02 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 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, H202. 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 1960s the group 8A 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?

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 4s or a 4p 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 3s and 3p 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 3s electrons and 4.9+ for the 3p electrons. Why are the values for the 3s and 3p electrons different? (d) If you remove a single electron from a 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 F2), 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 2p 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 lor the change in Zeff from C to N? (b) What change do you predict using Slater's rules? (c) The actual change in z.11 from C to N is 0.70+. Which approach to estimating Zeu is more accurate? (d) The change in Zeff from N to 0 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?

In the series of group SA hydrides, of general formula MH3, the measured bond distances are P-H, 1.419 A; As-H, 1.519 A; Sb-H, 1.707 A. (a) Compare these values with those estimated by use of the atomic radii in Figure 7.7. (b) Explain the steady increase in M- H bond distance in this series in terms of the electronic configurations of the M atoms.

Note from the following table that the increase in atomic radius in moving from Zr to Hf is smaller than in moving from Y to La. Suggest an explanation for this effect.

The "Chemistry and Life" box on ionic size in Section 7.3 compares the ionic radii of Zn2+ and Cd2+ (a) The 2+ ion of which other element seems the most obvious one to compare to Zn2+ and Cd2+? (b) With reference to Figure 2.24, is the element in part (a) essential for life? (c) Estimate the ionic radius of the 2+ ion of the element in part (a). Explain any assumptions you have made. (d) Would you expect the 2+ ion of the element in part (a) to be physiologically more similar to Zn2+ or to Cd2+? (e) Use a sourcebook or a Web search to determine whether the element in part (a) is toxic to humans.

The ionic substance strontium oxide, SrO, forms from the direct reaction of strontium metal with molecular oxygen. The arrangement of the ions in solid SrO is analogous to that in solid NaCl (see Figure 2.23) and is shown here. (a) Write a balanced equation for the formation of SrO(s) from the elements. (b) Based on the ionic radii in Figure 7.8, predict the length of the side of the cube in the figure (the distance from the center of an atom at one comer to the center of an atom at a neighboring comer). (c) The experimental density of SrO is 5.10 g/cm3 Given your answer to part (b), what is the number of formula units of SrO that are contained in the cube in the figure? (We will examine structures like those in the figure more closely in Chapter 11.)

Explain the variation in ionization energies of carbon, as displayed in the following graph:

Do you agree with the following statement? " A negative value for the electron affinity of an atom occurs when the outermost electrons incompletely shield one another from the nucleus." If not, change it to make it more nearly correct in your view. Apply either the statement as given or your revised statement to explain why the electron affinity of bromine is -325 kJ/mol and that for its neighbor Kr is > 0.

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 kJ/mol, for the group 1B and group 2B metals: (a) Why are the electron affinities of the group 2B elements greater than zero? (b) Why do the electron affinities of the group 1B 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.]

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.

Based on your reading of this chapter, arrange the following in order of increasing melting point: K, Br:u Mg, and 02. Explain the factors that determine the order.

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 Mg, Be resembles 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 oc Density: 7.3 g/ cm3 . Electrical conductivity: 20% that of copper. Hardness: About 1% as hard as iron. When 4.20 g of the unknown is heated in an excess of oxygen, 5.08 g of a white solid is formed. The solid could be sublimed by heating to over 800 oc. (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: Element Wavelength (A) Ne 14.610 Ca 3.358 Zn 1.435 Zr 0.786 Sn 0.491 (a) Calculate the frequency, v, of the X-rays emitted by each of the elements, in Hz. (b) Using graph paper (or suitable computer software), plot the square root of v versus the atomic number of the element. What do you observe about the plot? (c) Explain how the plot in part (b) allowed Moseley to predict the existence of undiscovered elements. (d) Use the result from part (b) to predict the X-ray wavelength emitted by iron. (e) A particular element emits X-rays with a wavelength of 0.980 A. What element do you think it is?

(a) Write the electron configuration for 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 (-2.18 X 10-l B J)(:) where 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. ooo (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 nm. (a) What is the energy of a photon of this light, in eV? (b) Write an equation that shows the process corresponding to the first ionization energy of Hg. (c) The kinetic energy of the emitted electrons is measured to be 10.75 eV. What is the first ionization energy of Hg, in kJ/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: Na{g) + Cl{g) ----> Na+{g) + 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 NaCl(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, MgO. The other is the product of the reaction of 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 MgO and magnesium nitride after burning is 0.470 g. Water is added to the crucible, further reaction occurs, and the crucible is heated to dryness until the final product is 0.486 g of 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-g Mg ribbon reacts with 2.57 g NH3{g) and the reaction goes to completion, which component is the limiting reactant? What mass of H2{g) is formed in the reaction? (e) The standard enthalpy of formation of solid magnesium nitride is -461.08 kJ/mol. Calculate the standard enthalpy change for the reaction between magnesium metal and ammonia gas.

Potassium superoxide, KOz, is often used in oxygen masks (such as those used by firefighters) because K02 reacts with C02 to release molecular oxygen. Experiments indicate that 2 mol of K02(s) react with each mole of C02{g). (a) The products of the reaction are K2C03(s) and 02{g). Write a balanced equation for the reaction between K02(s) and COz(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 K02(s) is needed to consume 18.0 g C02(g)? What mass of 02{g) is produced during this reaction?