Group Trends in Metals and Nonmetals

As discussed in the text, we can draw an analogy between the attraction of an electron to a nucleus and the act of perceiving light from a lightbulbin essence, the more nuclear charge the electron sees, the greater the attraction. (a) Using this analogy, discuss how screening by core electrons is analogous to putting a frosted-glass lampshade between the lightbulb and your eyes, as shown in the illustration. Lightbulb Frosted glass Observer (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 change the frosted glass to mimic the effect of moving down a column of the periodic table? [Section 7.2]

Which of these spheres represents F, which represents Br, and which represents Br -? [Section 7.3]

Consider the Mg2+, Cl -, K+, and Se 2- ions. The four spheres below represent these four ions, scaled according to ionic size. (a) Without referring to Figure 7.8, match each ion to its appropriate sphere. (b) In terms of size, between which of the spheres would you find the (i) Ca 2 + and (ii) S2 - ions? [Section 7.3]

In the following reaction Reactants Products which sphere represents a metal and which represents a nonmetal? Explain your answer. [Section 7.3]

Consider the A2X4 molecule depicted here, where A and X are elements. The AA bond length in this molecule is d1, and the four AX bond lengths are each d2. (a) In terms of d1 and d2, how could you define the bonding atomic radii of atoms A and X? (b) In terms of d1 and d2, what would you predict for the XX bond length of an X2 molecule? [Section 7.3] A A X X X X d1 d2

Shown below is a qualitative diagram of the atomic orbital energies for an Na atom. The number of orbitals in each subshell is not shown. 2p 3p 3s 2s 1s E (a) Are all of the subshells for n = 1, n = 2, and n = 3 shown? If not, what is missing? (b) The 2s and 2p energy levels are shown as different. Which of the following is the best explanation for why this is the case? (i) The 2s and 2p energy levels have different energies in the hydrogen atom, so of course they will have different energies in the sodium atom, (ii) The energy of the 2p orbital is higher than that of the 2s in all many-electron atoms. (iii) The 2s level in Na has electrons in it, whereas the 2p does not. (c) Which of the energy levels holds the highest-energy electron in a sodium atom? (d) A sodium vapor lamp (Figure 7.23)operates by using electricity to excite the highest-energy electron to the next highest-energy level. Light is produced when the excited electron drops back to the lower level. Which two energy levels are involved in this process for the Na atom? [Section 7.7]

(a) Which of the following charts below shows the general periodic trends for each of the following properties of the main-group elements (you can neglect small deviations going either across a row or down a column of the periodic table)? (1) Bonding atomic radius, (2) first ionization energy, and(3) metallic character. (b) Do any of the charts show the general periodic trends in the electron affinities of the main-group elements? [Sections 7.2-7.6]

(b) Do any of the charts show the general periodic trends in the electron affinities of the maingroup elements? [Sections 7.27.6]

An element X reacts with F21g2 to form the molecular product shown here. (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] F X F F F

( a ) Evaluate the expressions 2 * 1, 2 * 11 + 32, 2 * 11 + 3 + 52, and 2 * 11 + 3 + 5 + 72

(b) How do the atomic numbers of the noble gases relate to the numbers from part (a)? (c) What topic discussed in Chapter 6 is the source of the number 2 in the expressions in part (a)?

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 eka-manganese, 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 7.1, what do we now call the element Mendeleev called eka-manganese

You might have expected that the elements would have been discovered in order of their relative abundance in the Earths crust (Figure 1.6), but this is not the case. Suggest a general reason.

(a) Moseleys experiments on X rays emitted from atoms led to the concept of atomic numbers. Where exactly do these X rays come from? Draw an energy-level diagram to explain

(b) Why are chemical and physical properties of the elements more closely related to atomic numbers than they are to atomic weights?

(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?

Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? (i) The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom. (ii) Effective nuclear charge increases going left to right across a row of the periodic table. (iii) Valence electrons screen the nuclear charge more effectively than do core electrons. (iv) The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. (v) The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table.

Detailed calculations show that the value of Zeff for the outermost electrons in 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 Slaters 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? (e) Predict Zeff for the outermost electrons in the Rb atom based on the calculations for Na and K

Detailed calculations show that the value of Zeff for the outermost electrons in 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 Slaters 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? (e) Predict Zeff for a valence electron in P, phosphorus, based on the calculations for Si and Cl.

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?

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, Ti.

(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?

(d) If a free atom reacts to become part of a molecule, would you say that the atom gets smaller or larger?

(a) Why does the quantum mechanical description of manyelectron 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?

Tungsten has the highest melting point of any metal in the periodic table: 3422 C. 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.) (b) If you put tungsten metal under high pressure, predict what would happen to the distance between W atoms.

Which of the following statements about the bonding atomic radii in Figure 7.7 is incorrect? (i) For a given period, the radii of the representative elements generally decrease from left to right across period. (ii) The radii of the representative elements for the n = 3 period are all larger than those of the corresponding elements in the n = 2 period. (iii) For most of the representative elements, the change in radius from the n = 2 to the n = 3 period is greater than the change in radius from n = 3 to n = 4. (iv) The radii of the transition elements generally increase moving from left to right within a period. (v) The large radii of the Group 1A elements are due to their relatively small effective nuclear charges.

Estimate the AsI bond length from the data in Figure 7.7 and compare your value to the experimental AsI bond length in arsenic triiodide, AsI3, 2.55 .

The experimental BiI bond length in bismuth triiodide, BiI3, is 2.81 . Based on this value and data in Figure 7.7, predict the atomic radius of Bi.

Using only the periodic table, arrange each set of atoms in order from largest to smallest: (a) K, Li, Cs; (b) Pb, Sn, Si; (c) F, O, N.

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

Identify each statement as true or false: (a) Cations are larger than their corresponding neutral atoms. (b) Li + is smaller than Li. (c) Cl - is bigger than I -.

Explain the following variations in atomic or ionic radii: (a) I - 7 I 7 I + (b) Ca 2+ 7 Mg2+ 7 Be 2+ (c) Fe 7 Fe 2+ 7 Fe 3+

Which neutral atom is isoelectronic with each of the following ions? Ga 3+, Zr 4+, Mn7+, I -, Pb2+.

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) Cl -, (b) Sc3 + , (c) Fe 2 + , (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 Slaters rules to estimate the screening constant, S. (d) For isoelectronic ions, how are effective nuclear charge and ionic radius related?

Consider the isoelectronic ions Cl - and K+. (a) Which ion is smaller? (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 Slaters 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.

Arrange each of the following sets of atoms and ions, in order of increasing size: (a) Se 2-, Te 2-, Se; (b) Co3+, Fe 2+, Fe 3+; (c) Ca, Ti 4+, Sc3+; (d) Be 2+, Na +, Ne.

Provide a brief explanation for each of the following: (a) O2- is larger than O. (b) S2- is larger than O2- . (c) S2- is larger than K+. (d) K+ is larger than Ca 2+

In the ionic compounds LiF, NaCl, KBr, and RbI, the measured cationanion distances are 2.01 (LiF), 2.82 (NaCl), 3.30 (KBr), and 3.67 (RbI), respectively. (a) Predict the cationanion distance using the values of ionic radii given in Figure 7.8. (b) Calculate the difference between the experimentally measured ionion distances and the ones predicted from Figure 7.8. Assuming we have an accuracy of 0.04 in the measurement, would you say that the two sets of ionion distances are the same or not? (c) What estimates of the cationanion 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 an aluminum atom. Which process would require the least amount of energy?

Write equations that show the process for (a) the first two ionization energies of lead and (b) the fourth ionization energy of zirconium.

(a) Why does Li have a larger first ionization energy than Na?

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

Identify each statement as true or false: (a) Ionization energies are always negative quantities. (b) Oxygen has a larger first ionization energy than fluorine. (c) The second ionization energy of an atom is always greater than its first ionization energy. (d) The third ionization energy is the energy needed to ionize three electrons from a neutral atom.

(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 7A 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 smaller 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 smaller first ionization energy: (a) Ti, Ba; (b) Ag, Cu; (c) Ge, Cl; (d) Pb, Sb.

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) Co2 + , (b) Sn2 + , (c) Zr 4 + , (d) Ag+, (e) S2 - .

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) Ru3 + , (b) As 3 - , (c) Y3 + , (d) Pd2 + , (e) Pb2 + , (f) Au3 + .

Find three examples of ions in the periodic table that have an electron configuration of nd8 1n = 3, 4, 5, c2.

Find three atoms in the periodic table whose ions have an electron configuration of nd6 1n = 3, 4, 5, c2.

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

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.

Although 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 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?

Consider the following equation: Ca +1g2 + e- Ca1g2 Which of the following statements are true? (i) The energy change for this process is the electron affinity of the Ca + ion. (ii) The energy change for this process is the negative of the first ionization energy of the Ca atom. (iii) The energy change for this process is the negative of the electron affinity of the Ca atom.

(a) Does metallic character increase, decrease, or remain unchanged as one goes from left to right across a row of the periodic table?

(b) Does metallic character increase, decrease, or remain unchanged as one goes down a column of the periodic table?

(c) Are the periodic trends in (a) and (b) the same as or different from those for first ionization energy?

You read the following statement about two elements X and Y: One of the elements is a good conductor of electricity, and the other is a semiconductor. Experiments show that the first ionization energy of X is twice as great as that of Y. Which element has the greater metallic character?

Discussing this chapter, a classmate says, An element that commonly forms a cation is a metal. Do you agree or disagree? Explain your answer.

Discussing this chapter, a classmate says, Since elements that form cations are metals and elements that form anions are nonmetals, elements that do not form ions are metalloids. Do you agree or disagree? Explain your answer.

Predict whether each of the following oxides is ionic or molecular: SnO2, Al2O3, CO2, Li2O, Fe2O3, H2O.

Some metal oxides, such as Sc2O3, do not react with pure water, but they do react when the solution becomes either acidic or basic. Do you expect Sc2O3 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: CO2, CaO, Al2O3, SO3, SiO2, P2O5

Chlorine reacts with oxygen to form Cl2O7. (a) What is the name of this product (see Table 2.6)? (b) Write a balanced equation for the formation of Cl2O71l2 from the elements. (c) Under usual conditions, Cl2O7 is a colorless liquid with a boiling point of 81 C. Is this boiling point expected or surprising? (d) Would you expect Cl2O7 to be more reactive toward H+1aq2 or OH-1aq2? (e) If the oxygen in Cl2O7 is considered to have the -2 oxidation state, what is the oxidation state of the Cl? What is the electron configuration of Cl in this oxidation state?

An element X reacts with oxygen to form XO2 and with chlorine to form XCl4. XO2 is a white solid that melts at high temperatures (above 1000 C). Under usual conditions, XCl4 is a colorless liquid with a boiling point of 58 C. (a) XCl4 reacts with water to form XO2 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? (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.

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

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

Silver and rubidium both form +1 ions, but silver is far less reactive. Suggest an explanation, taking into account the ground-state electron configurations of these elements and their atomic radii.

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

(a) As described in Section 7.7, the alkali metals react with hydrogen to form hydrides and react with halogens to form halides. Compare the roles of hydrogen and halogens in these reactions. How are the forms of hydrogen and halogens 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?

Potassium and hydrogen react to form the ionic compound potassium hydride. (a) Write a balanced equation for this reaction. (b) Use data in Figures 7.10 and 7.12 to determine the energy change in kJ/mol for the following two reactions: K1g2 + H1g2 K+1g2 + H-1g2 K1g2 + H1g2 K-1g2 + H+1g2 (c) Based on your calculated energy changes in (b), which of these reactions is energetically more favorable (or less unfavorable)? (d) Is your answer to (c) consistent with the description of potassium hydride as containing hydride ions?

Compare the elements bromine 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) Would you expect At to be a metal, nonmetal, or metalloid? Explain. (c) 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. (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?

(a) Why does xenon react with fluorine, whereas neon does not?

(b) Using appropriate reference sources, look up the bond lengths of XeF bonds in several molecules. How do these numbers compare to the bond lengths calculated from the atomic radii of the elements?

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 1Z = 822. In how many instances are the atomic weights of the elements out of order relative to the atomic numbers of the elements?

Figure 7.4 shows the radial probability distribution functions for the 2s orbitals and 2p orbitals. (a) Which orbital, 2s or 2p, has more electron density close to the nucleus? (b) How would you modify Slaters rules to adjust for the difference in electronic penetration of the nucleus for the 2s and 2p orbitals?

(a) If the core electrons were totally effective at screening the valence electrons and the valence electrons provided no screening 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 Slaters 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?

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 5A hydrides, of general formula MH3, the measured bond distances are PH, 1.419 ; AsH, 1.519 ; SbH, 1.707 . (a) Compare these values with those estimated by use of the atomic radii in Figure 7.7. (b) Explain the steady increase in MH bond distance in this series in terms of the electron configurations of the M atoms

In Table 7.8, the bonding atomic radius of neon is listed as 0.58 , whereas that for xenon is listed as 1.40 . A classmate of yours states that the value for Xe is more realistic than the one for Ne. Is she correct? If so, what is the basis for her statement?

The AsAs bond length in elemental arsenic is 2.48 . The ClCl bond length in Cl2 is 1.99 . (a) Based on these data, what is the predicted AsCl bond length in arsenic trichloride, AsCl3, in which each of the three Cl atoms is bonded to the As atom? (b) What bond length is predicted for AsCl3, using the atomic radii in Figure 7.7?

The following observations are made about two hypothetical elements A and B: The AA and BB bond lengths in elemental A and B are 2.36 and 1.94 , respectively. A and B react to form the binary compound AB2, which has a linear structure (that is B - A - B = 180). Based on these statements, predict the separation between the two B nuclei in a molecule of AB2.

Elements in group 7A in the periodic table are called the halogens; elements in group 6A are called the chalcogens. (a) What is the most common oxidation state of the chalcogens compared to the halogens? (b) For each of the following periodic properties, state whether the halogens or the chalcogens have larger values: atomic radii, ionic radii of the most common oxidation state, first ionization energy, second ionization energy

Note from the following table that there is a significant increase in atomic radius upon moving from Y to La whereas the radii of Zr to Hf are the same. Suggest an explanation for this effect. Atomic Radii () Sc 1.70 Ti 1.60 Y 1.90 Zr 1.75 La 2.07 Hf 1.75

(a) Which ion is smaller, Co3+ or Co4+?

(b) In a lithium-ion battery that is discharging to power a device, for every Li + that inserts into the lithium cobalt oxide electrode, a Co4+ ion must be reduced to a Co3+ ion to balance charge. Using the CRC Handbook of Chemistry and Physics or other standard reference, find the ionic radii of Li +, Co3+, and Co4+. Order these ions from smallest to largest.

(c) Will the lithium cobalt electrode expand or contract as lithium ions are inserted?

(d) Lithium is not nearly as abundant as sodium. If sodium ion batteries were developed that function as lithium ion ones, do you think sodium cobalt oxide would still work as the electrode material? Explain.

e) If you dont think cobalt would work as the redox-active partner ion in the sodium version of the electrode, suggest an alternative metal ion and explain your reasoning.

The ionic substance strontium oxide, SrO, forms from the reaction of strontium metal with molecular oxygen. The arrangement of the ions in solid SrO is analogous to that in solid NaCl: (a) Write a balanced equation for the formation of SrO(s) from its 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 corner to the center of an atom at a neighboring corner). (c) The density of SrO is 5.10 g>cm3 . Given your answer to part (b), how many formula units of SrO are contained in the cube shown here?

Explain the variation in the ionization energies of carbon, as displayed in this graph: 1 2 3 4 5 6 10,000 20,000 30,000 40,000 50,000 Ionization number Ionization energies of carbon (kJ mol1 )

Group 4A elements have much more negative electron affinities than their neighbors in groups 3A and 5A (see Figure 7.12). Which of the following statements best explains this observation? (i) The group 4A elements have much higher first ionization energies than their neighbors in groups 3A and 5A. (ii) The addition of an electron to a group 4A element leads to a half-filled np3 outer electron configuration. (iii) The group 4A elements have unusually large atomic radii. (iv) The group 4A elements are easier to vaporize than are the group 3A and 5A elements.

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 A1g2 + A1g2 A+1g2 + A-1g2 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?

(a) Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons.

(b) Why does O3 - not exist?

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.

Identify two ions that have the following ground-state electron configurations: (a) [Ar], (b) 3Ar43d5 , (c) 3Kr45s 2 4d10 .

Which of the following chemical equations is connected to the definitions of (a) the first ionization energy of oxygen, (b) the second ionization energy of oxygen, and (c) the electron affinity of oxygen? (i) O1g2 + e- O-1g2 (ii) O1g2 O+1g2 + e- (iii) O1g2 + 2e- O2-1g2 (iv) O1g2 O2+1g2 + 2e- (v) O+1g2 O2+1g2 + e-

The electron affinities, in kJ>mol, for the group 1B and group 2B metals are as follows: Hg > 0 Au 223 Cd > 0 Ag 126 Zn > 0 Cu 119 (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 affinities 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 nonmetals. 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. (d) The hydride ion is H-. Write out the process corresponding to the first ionization energy of the hydride ion. (e) How does the process in part (d) compare to the process for the electron affinity of a neutral hydrogen atom?

The first ionization energy of the oxygen molecule is the energy required for the following process: O21g2 O2 +1g2 + eThe energy needed for this process is 1175 kJ>mol, very similar to the first ionization energy of Xe. Would you expect O2 to react with F2? If so, suggest a product or products of this reaction.

It is possible to define metallic character as we do in this book and base it on the reactivity of the element and the ease with which it loses electrons. Alternatively, one could measure how well electricity is conducted by each of the elements to determine how metallic the elements are. On the basis of conductivity, there is not much of a trend in the periodic table: Silver is the most conductive metal, and manganese the least. Look up the first ionization energies of silver and manganese; which of these two elements would you call more metallic based on the way we define it in this book?

Which of the following is the expected product of the reaction of K(s) and H21g2? (i) KH(s), (ii) K2H1s2, (iii) KH21s2, (iv) K2H21s2, or (v) K(s) and H21g2 will not react with one another.

Elemental cesium reacts more violently with water than does elemental sodium. Which of the following best explains this difference in reactivity? (i) Sodium has greater metallic character than does cesium. (ii) The first ionization energy of cesium is less than that of sodium. (iii) The electron affinity of sodium is smaller than that of cesium. (iv) The effective nuclear charge for cesium is less than that of sodium. (v) The atomic radius of cesium is smaller than that of sodium.

(a) 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, H2O2. When the solution is tested in a burner flame, a lilac-purple flame is produced. What is the likely identity of the metal?

(b) Write a balanced chemical equation for the reaction of the white substance with water.

Zinc in its 2+ oxidation state is an essential metal ion for life. Zn2+ is found bound to many proteins that are involved in biological processes, but unfortunately Zn2+ is hard to detect by common chemical methods. Therefore, scientists who are interested in studying Zn2+-containing proteins frequently substitute Cd2+ for Zn2+, since Cd2+ is easier to detect. (a) On the basis of the properties of the elements and ions discussed in this chapter and their positions in the periodic table, describe the pros and cons of using Cd2+ as a Zn2+ substitute. (b) Proteins that speed up (catalyze) chemical reactions are called enzymes. Many enzymes are required for proper metabolic reactions in the body. One problem with using Cd2+ to replace Zn2+ in enzymes is that Cd2+ substitution can decrease or even eliminate enzymatic activity. Can you suggest a different metal ion that might replace Zn2+ in enzymes instead of Cd2+? Justify your answer.

A historian discovers a nineteenth-century notebook in which some observations, dated 1822, were recorded on a substance thought to be a new element. Here are some of the data recorded in the notebook: Ductile, silver-white, metallic looking. Softer than lead. Unaffected by water. Stable in air. Melting point: 153 C. 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 C. (a) Using information in the text and the CRC Handbook of Chemistry and Physics, 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 7.1, might this nineteenth- century investigator have been the first to discover a new element?

In April 2010, a research team reported that it had made Element 117. This discovery was confirmed in 2012 by additional experiments. Write the ground-state electron configuration for Element 117 and estimate values for its first ionization energy, electron affinity, atomic size, and common oxidation state based on its position in the periodic table.

We will see in Chapter 12 that semiconductors are materials that conduct electricity better than nonmetals but not as well as metals. The only two elements in the periodic table that are technologically useful semiconductors are silicon and germanium. Integrated circuits in computer chips today are based on silicon. Compound semiconductors are also used in the electronics industry. Examples are gallium arsenide, GaAs; gallium phosphide, GaP; cadmium sulfide, CdS; and cadmium selenide, CdSe. (a) What is the relationship between the compound semiconductors compositions and the positions of their elements on the periodic table relative to Si and Ge? (b) Workers in the semiconductor industry refer to IIVI and IIIV materials, using Roman numerals. Can you identify which compound semiconductors are IIVI and which are IIIV? (c) Suggest other compositions of compound semiconductors based on the positions of their elements in the periodic table.

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 () Ne 14.610 Ca 3.358 Zn 1.435 Zr 0.786 Sn 0.491 (a) Calculate the frequency, n, of the X rays emitted by each of the elements, in Hz. (b) Plot the square root of n 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 . 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 1-2.18 * 10-18 J2 a Z2 n2 b, 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 part (c)?

One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect. (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 Design an Experiment 297 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) Using Figure 7.10, determine which of the halogen elements has a first ionization energy closest to that of mercury.

Mercury in the environment can exist in oxidation states 0, +1, and +2. One major question in environmental chemistry research is how to best measure the oxidation state of mercury in natural systems; this is made more complicated by the fact that mercury can be reduced or oxidized on surfaces differently than it would be if it were free in solution. XPS, X-ray photoelectron spectroscopy, is a technique related to PES (see Exercise 7.111), but instead of using ultraviolet light to eject valence electrons, X rays are used to eject core electrons. The energies of the core electrons are different for different oxidation states of the element. In one set of experiments, researchers examined mercury contamination of minerals in water. They measured the XPS signals that corresponded to electrons ejected from mercurys 4f orbitals at 105 eV, from an X-ray source that provided 1253.6 eV of energy. The oxygen on the mineral surface gave emitted electron energies at 531 eV, corresponding to the 1s orbital of oxygen. Overall the researchers concluded that oxidation states were +2 for Hg and -2 for O. (a) Calculate the wavelength of the X rays used in this experiment. (b) Compare the energies of the 4f electrons in mercury and the 1s electrons in oxygen from these data to the first ionization energies of mercury and oxygen from the data in this chapter. (c) Write out the ground-state electron configurations for Hg2+ and O2-; which electrons are the valence electrons in each case? (d) Use Slaters rules to estimate Zeff for the 4f and valence electrons of Hg2+ and O2-; assume for this purpose that all the inner electrons with 1n - 32 or less screen a full +1.

When magnesium metal is burned in air (Figure 3.6), 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 NH31g2 and the reaction goes to completion, which component is the limiting reactant? What mass of H21g2 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

(a) The measured BiBr bond length in bismuth tribromide, BiBr3, is 2.63 . Based on this value and the data in Figure 7.8, predict the atomic radius of Bi.

(b) Bismuth tribromide is soluble in acidic solution. It is formed by treating solid bismuth(III) oxide with aqueous hydrobromic acid. Write a balanced chemical equation for this reaction.

(c) While bismuth(III) oxide is soluble in acidic solutions, it is insoluble in basic solutions such as NaOH(aq). Based on these properties, is bismuth characterized as a metallic, metalloid, or nonmetallic element?

(d) Treating bismuth with fluorine gas forms BiF5. Use the electron configuration of Bi to explain the formation of a compound with this formulation.

(e) While it is possible to form BiF5 in the manner just described, pentahalides of bismuth are not known for the other halogens. Explain why the pentahalide might form with fluorine but not with the other halogens. How does the behavior of bismuth relate to the fact that xenon reacts with fluorine to form compounds but not with the other halogens?

Potassium superoxide, KO2, is often used in oxygen masks (such as those used by firefighters) because KO2 reacts with CO2 to release molecular oxygen. Experiments indicate that 2 mol of KO21s2 react with each mole of CO21g2. (a) The products of the reaction are K2CO31s2 and O21g2. Write a balanced equation for the reaction between KO21s2 and CO21g2. (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 KO21s2 is needed to consume 18.0 g CO21g2? What mass of O21g2 is produced during this reaction?

Problem 1DE In this chapter we have seen that the reaction of potassium metal with oxygen leads to a product that we might not expect, namely, potassium superoxide, KO2(s). Let’s design some experiments to learn more about this unusual product. (a) One of your team members proposes that the capacity to form a superoxide such as KO2 is related to a low value for the first ionization energy. How would you go about testing this hypothesis for the metals of group 1A? What other periodic property of the alkali metals might be considered as a factor favoring superoxide formation? ________________ (b) KO2(s) is the active ingredient in many breathing masks used by firefighters because it can be used as a source of O2(g). In principle, KO2(s) can react with both major components of human breath, H2O(g) and CO2(g), to produce O2(g) and other products (all of which follow the expected patterns of reactivity we have seen). Predict the other products in these reactions and design experiments to determine whether KO2(s) does actually react with both H2O(g) and CO2(g). ________________ (c) Propose an experiment to determine whether either of the reactions in part (b) is more important in the operation of a firefighter’s breathing mask. ________________ (d) The reaction of K(s) and O2(g) leads to a mixture of KO2(s) and K2O(s). Use ideas presented in this exercise to design an experiment to determine the percentages of KO2(s) and K2O(s) in the product mixture that results from the reaction of K(s) with excess O2(g).

As discussed in the text, we can draw an analogy between the attraction of an electron to a nucleus and the act of perceiving light from a lightbulb—in essence, the more nuclear charge the electron “sees,” the greater the attraction. (a) Using this analogy, discuss how screening 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 change the frosted glass to mimic the effect of moving down a column of the periodic table? [Section 7.2]

Which of these spheres represents \(F\), which represents \(B r\)-, and which represents \(B r\)-? [Section 7.3] Equation Transcription: Text Transcription: F Br Br

Consider the \(\mathrm{Mg}^{2+}, \mathrm{Cl}^{-}, \mathrm{K}^{+}, \text {and } \mathrm{Se}^{2-}\) ions. The four spheres below represent these four ions, scaled according to ionic size. (a) Without referring to Figure 7.8, match each ion to its appropriate sphere. How do cations of the same charge change in radius as you move down a column in the periodic table? ? Figure 7.8 Cation and anion size. Radii, in angstroms, of atoms and their ions for five groups of representative elements. (b) In terms of size, between which of the spheres would you find the (i) \(\mathrm{Ca}^{2+}\) and (ii) \(\mathrm{S}^{2-}\) ions? [Section 7.3] Equation Transcription: Text Transcription: Mg2+, Cl-, K+, and Se2- Ca2+ S2-

In the following reaction which sphere represents a metal and which represents a nonmetal? Explain your answer. [Section 7.3]

Shown below is a qualitative diagram of the atomic orbital energies for an Na atom. The number of orbitals in each subshell is not shown. (a) Are all of the subshells for \(n=1, n=2, \text { and } n=3\) shown? If not, what is missing? (b) The \(2 s \text { and } 2 p\) energy levels are shown as different. Which of the following is the best explanation for why this is the case? (i) The \(2 s \text { and } 2 p\) energy levels have different energies in the hydrogen atom, so of course they will have different energies in the sodium atom, (ii) The energy of the \(2 p\) orbital is higher than that of the \(2 s\) in all many-electron atoms. (iii) The \(2 s\) level in Na has electrons in it, whereas the \(2 p\) does not. (c) Which of the energy levels holds the highest-energy electron in a sodium atom? (d) A sodium vapor lamp (Figure 7.23) operates by using electricity to excite the highest-energy electron to the next highest-energy level. Light is produced when the excited electron drops back to the lower level. Which two energy levels are involved in this process for the Na atom? [Section 7.7] If we had potassium vapor lamps, what color would they be? ? Figure 7.23 The characteristic yellow light in a sodium lamp results from excited electrons in the high-energy \(3 p\) orbital falling back to the lower-energy \(3 s\) orbital. Equation Transcription: Text Transcription: n=1, n=2, and n=3 2s and 2p 2s and 2p 2p 2s 2s 2p 2p 3p 3s

Consider the \(A_{2} X_{4}\) molecule depicted here, where \(A \text { and } X\) are elements. The \(A \neg A\) bond length in this molecule is \(d 1\), and the four \(A-X\) bond lengths are each \(d_{2}\). (a) In terms of \(d_{1} \text { and } d_{2}\), how could you define the bonding atomic radii of atoms \(A \text { and } X\)? (b) In terms of \(d_{1} \text { and } d_{2}\), what would you predict for the \(X-X\) bond length of an \(X 2\) molecule? [Section 7.3] Equation Transcription: Text Transcription: A_2X_4 A and X A \neg A d_1 A-X d_2 d1 and d2 A and X d1 and d2 X-X X2

(a) Which of the following charts below shows the general periodic trends for each of the following properties of the main-group elements (you can neglect small deviations going either across a row or down a column of the periodic table)? (1) Bonding atomic radius, (2) first ionization energy, and (3) metallic character. (b) Do any of the charts show the general periodic trends in the electron affinities of the maingroup elements? [Sections 7.2–7.6]

Problem 9E Periodic Table; Effective Nuclear Charge (Sections) (a) Evaluate the expressions 2x1,2x(1 +3), 2 x (1 + 3 + 5), and 2 x (1 + 3 + 5 + 7). (b) How do the atomic numbers of the noble gases relate to the numbers from part (a)? (c) What topic discussed in Chapter is the source of the number “2" in the expressions in part (a)?

An element \(X\) reacts with F21g2 to form the molecular product shown here. (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] Equation Transcription: Text Transcription: X X X

You might have expected that the elements would have been discovered in order of their relative abundance in the Earth’s crust (Figure 1.6), but this is not the case. Suggest a general reason. Relative abundances of elements.*Elements in percent by mass in Earth’s crust (including oceans and atmosphere) and the human body. Name two significant differences between the elemental composition of Earth’s crust and the elemental composition of the human body.

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 eka-manganese, 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 7.1, what do we now call the element Mendeleev called eka-manganese. Discovering the elements. Copper, silver, and gold have all been known since ancient times, whereas most of the other metals have not. Can you suggest an explanation?

Problem 12E Periodic Table; Effective Nuclear Charge (Sections) (a) Moseley’s experiments on X rays emitted from atoms led to the concept of atomic numbers Where exactly do these X rayscome from? Draw an energy-level diagram to explain, (b) Why are chemical and physical properties of the elements more closely related to atomic numbers than they are to atomic weights?

Problem 14E Periodic Table; Effective Nuclear Charge (Sections) Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? (i) The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom, (ii) Effective nuclear charge increases going left to right across a row of the periodic table, (iii) Valence lectrons screen the nuclear charge more effectively than do core electrons. (iv) The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table, (v) The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table.

Problem 13E Periodic Table; Effective Nuclear Charge (Sections) (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?

Problem 15E Periodic Table; Effective Nuclear Charge (Sections) Detailed calculations show that the value of Zeff for the outermost electrons in 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? (e) Predict Zeff for the outermost electrons in the Rb atom based on the calculations for Na and K.

Problem 16E Periodic Table; Effective Nuclear Charge (Sections) Detailed calculations show that the value of Zeff for the outermost electrons in 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? (e) Predict Zeff for a valence electron in P. phosphorus, based on the calculations for Si and Cl.

Problem 17E Periodic Table; Effective Nuclear Charge (Sections) 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?

Problem 18E Periodic Table; Effective Nuclear Charge (Sections) 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. Ti.

Problem 19E Atomic and Ionic Radii (Section) (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? (d) If a free atom reacts to become part of a molecule, would you say that the atom gets smaller or larger?

Problem 20E Atomic and Ionic Radii (Section) (a) Why does the quantum mechanical description of manyelectron 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?

Problem 22E Which of the following statements about the bonding atomic radii in Figure is incorrect? (i) For a given period, the radii of the representative elements generally decrease from left to right across period, (ii) The radii of the representative elements for the n = 3 period are all larger than those of the corresponding elements in the n = 2 period, (iii) For most of the representative elements, the change in radius from the n = 2 to the n = 3 period is greater than the change in radius from n = 3 to n = 4. (iv) The radii of the transition elements generally increase moving from left to right within a period, (v) The large radii of the Group 1A elements are due to their relatively small effective nuclear charges.

Problem 21E Atomic and Ionic Radii (Section) Tungsten has the highest melting point of any metal in the periodic table: 3422 °C. The distance between W atoms in tungsten metal is 2.74 A. (a) What is the atomic radius of a tungsten atom in this environment? (This radius is called the metallic radius.) (b) If you put tungsten metal under high pressure, predict what would happen to the distance between W atoms.

The experimental \(B i-l\) bond length in bismuth triiodide, \(B i l_{3} \text { is } 2.81 \AA\). Based on this value and data in Figure 7.7, predict the atomic radius of \(B i\) Trends in bonding atomic radii for periods 1 through 5. Which part of the periodic table (top/bottom, left/right) has the elements with the largest atoms? Equation Transcription: Å Text Transcription: B i-l B i l_{3} \text { is } 2.81 \AA Bi

Problem 26E Atomic and Ionic Radii (Section) Using only the periodic table, arrange each set of atoms in order of increasing radius: (a) Ba, Ca. Na; (b) In. Sn. As; (c) Al. Be. Si.

Estimate the \(A s-l\) bond length from the data in Figure 7.7 and compare your value to the experimental \(A s-l\) bond length in arsenic triiodide, \(A s l_{3}, 2.55 \AA\) Trends in bonding atomic radii for periods 1 through 5. Which part of the periodic table (top/bottom, left/right) has the elements with the largest atoms? Equation Transcription: Å Text Transcription: As-l As-l A s l_{3}, 2.55 \AA

Problem 25E Atomic and Ionic Radii (Section) Using only the periodic table, arrange each set of atoms in order from largest to smallest: (a) K. Li, Cs; (b) Pb. Sn, Si; (c) F. O. N.

Problem 29E Atomic and Ionic Radii (Section) Which neutral atom is isoelectronic with each of the following ions? Ga3+, Zr4+, Mn7+, I-, Pb2+.

Problem 27E Atomic and Ionic Radii (Section) Identify each statement as true or false: (a) Cations are larger than their corresponding neutral atoms, (b) Li+ is smaller than Li. (c) Cl- is bigger than I-.

Problem 28E Atomic and Ionic Radii (Section) Explain the following variations in atomic or ionic radii: (a) I- >1 >1+ (b) Ca2+ >Mg2+ >Be2+ (c) Fe >Fe2+ >Fe3+

Problem 30E Atomic and Ionic Radii (Section) 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) CI-, (b) Sc3 + , (c) Fe2 + , (d) Zn2 + , (e) Sn4 + .

Problem 31E Atomic and Ionic Radii (Section) 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?

Problem 32E Atomic and Ionic Radii (Section) Consider the isoelectronic ions Cl- and K+. (a) Which ion is smaller? (b) Using Equation 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? Zeff = Z - S [7.1]

Problem 34E Atomic and Ionic Radii (Section) Arrange each of the following sets of atoms and ions, in order of increasing size: (a) Se2-, Te2-, Se; (b) Co3+, Fe2+, Fe3+; (c) Ca. Ti4+, Sc3+; (d) Be2+, Na+, Ne.

Problem 33E Atomic and Ionic Radii (Section) 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.

Problem 35E Atomic and Ionic Radii (Section) Provide a brief explanation for each of the following: (a) 02-is larger than O. (b) S2- is larger than 02-. (c) S2- is larger than K+. (d) K+ is larger than Ca2+.

Problem 37E Ionization Energies; Electron Affinities (Sections) Write equations that show the processes that describe the first, second, and third ionization energies of an aluminum atom. Which process would require the least amount of energy?

Problem 38E Ionization Energies; Electron Affinities (Sections) Write equations that show the process for (a) the first two ionization energies of lead and (b) the fourth ionization energy of zirconium.

In the ionic compounds \(\text { Lif, } \mathrm{NaCl}, \mathrm{KBr} \text {, and } \mathrm{Rbl}\), the measured cation-anion distances are \(2.01 \AA(L i-F), 2.82 \AA(N a-C l), 3.30 \AA(K-B r), \text { and } 3.67 \AA(R b-l)\), respectively. (a) Predict the cation-anion distance using the values of ionic radii given in Figure 7.8. (b) Calculate the difference between the experimentally measured ion-ion distances and the ones predicted from Figure 7.8. Assuming we have an accuracy of \(0.04 \AA\) in the measurement, would you say that the two sets of ion-ion distances are the same or 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? Cation and anion size. Radii, in angstroms, of atoms and their ions for five groups of representative elements. How do cations of the same charge change in radius as you move down a column in the periodic table? Equation Transcription: ÅÅ Å Å Å Text Transcription: Lif, NaCl, KBr, and Rbl 2.01 \AA(L i-F), 2.82 \AA(N a-C l), 3.30 \AA(K-B r), \text { and } 3.67 \AA(R b-l) 0.04 \AA

Problem 39E Ionization Energies; Electron Affinities (Sections) (a) Why does Li have a larger first ionization energy than Na? (b) The difference between the third and fourth ionization energies of scandium is much larger than that of titanium. Why? (c) Why does Li have a much larger second ionization energy than Be?

Problem 41E Ionization Energies; Electron Affinities (Sections) (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?

Problem 40E Ionization Energies; Electron Affinities (Sections) Identify each statement as true or false (a) Ionization energies are always negative quantities, (b) Oxygen has a larger first ionization energy than fluorine, (c) The second ionization energy of an atom is always greater than its first ionization energy, (d) The third ionization energy is the energy needed to ionize three electrons from a neutral atom.

Problem 42E Ionization Energies; Electron Affinities (Sections) (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?

Problem 43E Ionization Energies; Electron Affinities (Sections) Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) Cl. Ar; (b) Be. Ca; (c) K. Co; (d) S. Ge; (e) Sn. Te.

Problem 44E Ionization Energies; Electron Affinities (Sections) For each of the following pairs, indicate which element has the smaller first ionization energy: (a) Ti, Ba: (b) Ag. Cu; (c) Ge, Cl; (d) Pb. Sb

Problem 45E Ionization Energies; Electron Affinities (Sections) Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) Co2 +. (b) Sn2 +, (c) Zr4 +, (d) Ag+, (e) S2-.

Problem 46E Ionization Energies; Electron Affinities (Sections) Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) Ru3 + , (b) As3 -, (c) Y3 + , (d) Pd2 + , (e) Pb2 + , (f) Au3 +.

Problem 47E Ionization Energies; Electron Affinities (Sections) Find three examples of ions in the periodic table that have an electron configuration of ndS (n 3, 4, 5,c)

Problem 48E Ionization Energies; Electron Affinities (Sections) Find three atoms in the periodic table whose ions have an electron configuration of nd6 (n = 3, 4, 5,c).

Problem 49E Ionization Energies; Electron Affinities (Sections) The first ionization energy and electron affinity of Ar are both positive values, (a) What is the significance of the positive value in each case? (b) What are the units of electron affinity?

Problem 50E Ionization Energies; Electron Affinities (Sections) f 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 51E Ionization Energies; Electron Affinities (Sections) Although 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.

Problem 52E Ionization Energies; Electron Affinities (Sections) 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 53E Ionization Energies; Electron Affinities (Sections) 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 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 54E Ionization Energies; Electron Affinities (Sections) Consider the following equation; Ca+(gj + e—Ca (g) Which of the following statements are true? (i) The energy change for this process is the electron affinity of the Ca+ ion. (ii) The energy change for this process is the negative of the first ionization energy of the Ca atom, (iii) The energy change for this process is the negative of the electron affinity of the Ca atom.

Problem 55E Properties of Metals and Nonmetals (Section) (a) Does metallic character increase, decrease, or remain unchanged as one goes from left to right across a row of the periodic table? (b) Does metallic character increase, decrease, or remain unchanged as one goes down a column of the periodic table? (c) Are the periodic trends in (a) and (b) the same as or different from those for first ionization energy?

Problem 56E Properties of Metals and Nonmetals (Section) You read the following statement about two elements X and Y: One of the elements is a good conductor of electricity, and the other is a semiconductor. Experiments show that the first ionization energy of X is twice as great as that of Y. Which element has the greater metallic character?

Problem 57E Properties of Metals and Nonmetals (Section) You read the following statement about two elements X and Y: One of the elements is a good conductor of electricity, and the other is a semiconductor. Experiments show that the first ionization energy of X is twice as great as that of Y. Which element has the greater metallic character?

Problem 59E Properties of Metals and Nonmetals (Section) Discussing this chapter, a classmate says, "Since elements that form cations are metals and elements that form anions are nonmetals, elements that do not form ions are metalloids." Do you agree or disagree? Explain your answer

Problem 58E Properties of Metals and Nonmetals (Section) Discussing this chapter, a classmate says, "An element that commonly forms a cation is a metal." Do you agree or disagree? Explain your answer

Problem 61E Properties of Metals and Nonmetals (Section) Some metal oxides, such as Sc203. 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.

Problem 62E Properties of Metals and Nonmetals (Section) (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?

Problem 60E Properties of Metals and Nonmetals (Section) Predict whether each of the following oxides is ionic or molecular: Sn02. AI203, C02. Li20. Fe203, H20.

Problem 63E Properties of Metals and Nonmetals (Section) Arrange the following oxides in order of increasing acidity: C02. CaO. AI203, S03, Si02. P205.

An element \(\mathrm{X}\) 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 \(1000^{\circ} \mathrm{C}\)). 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 \(\mathrm{X}\) is a metal, nonmetal, or metalloid? (c) By using a sourcebook such as the CRC Handbook of Chemistry and Physics, try to determine the identity of element \(\mathrm{X}\) Equation Transcription: Text Transcription: X XO2 XCl4, XO2 1000°C XCl4 58°C XCl4 XO2 X X

Problem 65E Properties of Metals and Nonmetals (Section) An element X reacts with oxygen to form X02 and with chlorine to form XCI4. X02 is a white solid that melts at high temperatures (above 1000 °C). Under usual conditions, XCI4 is a colorless liquid with a boiling point of 58 °C. (a) XCI4 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? (c) By using a sourcebook such as the CRC Handbook of Chemistry and Physics, try to determine the identity of element X.

Problem 66E Properties of Metals and Nonmetals (Section) Write balanced equations for the following reactions: (a) barium oxide with water, (b) iron(ll) oxide with perchloric acid, (c) sulfur trioxide with water, (d) carbon dioxide with aqueous sodium hydroxide.

Problem 69E Group Trends in Metals and Nonmetals (Sections) Write a balanced equation for the reaction that occurs in each of the following cases: (a) Potassium metal is exposed to 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 reacts with molten sulfur.

Problem 67E Group Trends in Metals and Nonmetals (Sections) (a) Why is calcium generally more reactive than magnesium? (b) Why is calcium generally less reactive than potassium?

Problem 68E Group Trends in Metals and Nonmetals (Sections) Silver and rubidium both form +1 ions, but silver is far less reactive. Suggest an explanation, taking into account the ground-state electron configurations of these elements and their atomic radii.

Problem 70E Group Trends in Metals and Nonmetals (Sections) Write a balanced equation for the reaction that occurs in each of the following cases: (a) Cesium is added to water, (b) Strontium is added to water, (c) Sodium reacts with oxygen, (d) Calcium reacts with iodine.

Problem 71E Group Trends in Metals and Nonmetals (Sections) (a) As described in Section, the alkali metals react with hydrogen to form hydrides and react with halogens to form halides. Compare the roles of hydrogen and halogens in these reactions. How are the forms of hydrogen and halogens in the products alike? (b) Write balanced equations for the reaction of fluorinewith calcium and for the reaction of hydrogen with calcium. What are the similarities among the products of these reactions?

Potassium and hydrogen react to form the ionic compound potassium hydride. (a) Write a balanced equation for this reaction. (b) Use data in Figures 7.10 and 7.12 to determine the energy change in kJ/mol for the following two reactions: \(K(g)+H(g) \rightarrow K^{+}(g)+H^{-}(g)\) \(K(g)+H(g) \rightarrow K^{-}(g)+H^{+}(g)\) (c) Based on your calculated energy changes in (b), which of these reactions is energetically more favorable (or less unfavorable)? (d) Is your answer to (c) consistent with the description of potassium hydride as containing hydride ions? The value for astatine, At, is missing in this figure. To the nearest \(100 \mathrm{~kJ} / \mathrm{mol}\), what estimate would you make for the first ionization energy of At? ? Figure 7.10 Trends in first ionization energies of the elements. Why are the electron affinities of the Group 4A elements more negative than those of the Group 5A elements? ? Figure 7.12 Electron affinity in kJ ,mol for selected s- and p-block elements. Equation Transcription: Text Transcription: K(g)+H(g) \rightarrow K^{+}(g)+H^{-}(g) K(g)+H(g) \rightarrow K^{-}(g)+H^{+}(g) 100 \mathrm{~kJ} / \mathrm{mol}

Problem 73E Group Trends in Metals and Nonmetals (Sections) Compare the elements bromine and chlorine with respect to the following properties: (a) electron configuration, (b) mostcommon ionic charge, (c) first ionization energy, (d) reactivity toward water, (e) electron affinity, (f) atomic radius. Account for the differences between the two elements.

Problem 74E Group Trends in Metals and Nonmetals (Sections) 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) Doyou expect the element to be a gas, liquid, or solid at room temperature? Explain. (b) Would you expect At to be a metal, nonmetal, or metalloid? Explain. (c) What is the chemical formula of the compound it forms with Na?

Problem 76E Group Trends in Metals and Nonmetals (Sections) (a) Why does xenon react with fluorine, whereas neon does not? (b) Using appropriate reference sources, look up the bond lengths of Xe—F bonds in several molecules. How do these numbers compare to the bond lengths calculated from the atomic radii of the elements?

Problem 75E Group Trends in Metals and Nonmetals (Sections) Until the early 1960s the group 8A elements were called the inert gases. (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?

Problem 77E Group Trends in Metals and Nonmetals (Sections) 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.

Figure 7.4 shows the radial probability distribution functions for the \(2 s\) orbitals and \(2 p\) orbitals. (a) Which orbital, \(2 s \text { or } 2 p\), has more electron density close to the nucleus? (b) How would you modify Slater’s rules to adjust for the difference in electronic penetration of the nucleus for the \(2 s \text { and } 2 p\) orbitals? Equation Transcription: Text Transcription: 2s 2p 2s or 2p 2s and 2p

Problem 81AE (a) If the core electrons were totally effective at screening the valence electrons and the valence electrons provided no screening 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?

Problem 78E Group Trends in Metals and Nonmetals (Sections) 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.

Problem 79AE Consider the stable elements through lead (Z = 82). In how many instances are the atomic weights of the elements out of order relative to the atomic numbers of the elements?

Problem 82AE 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 5A hydrides, of general formula \(M H_{3}\), the measured bond distances are \(P-H, 1.1419 \AA \text { A } ; A s-H, 1.519 \AA \text {; } S b-H, 1.707 \AA\) (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 electron configurations of the M atoms. Equation Transcription: ÅÅÅ Text Transcription: M H_{3} P-H, 1.1419 \AA \text { A } ; A s-H, 1.519 \AA \text {; } S b-H, 1.707 \AA M-H

Problem 84AE In Table the bonding atomic radius of neon is listed as 0.58 Å, whereas that for xenon is listed as 1.40 Å. A classmate of yours states that the value for Xe is more realistic than the one for Ne. Is she correct? If so, what is the basis for her statement? Table Some Properties of the Noble Gases Element Electron Configuration Boiling Point (K) Density (g/L) Atomic Radius* (Å) I1 (kJ/mol) Helium 1s2 4.2 0.18 0.28 2372 Neon [He]2s22p6 27.1 0.90 0.58 2081 Argon [Ne]3s23p6 87.3 1.78 1.06 1521 Krypton [Ar]4s23d104p6 120 3.75 1.16 1351 Xenon [Kr]5s24d105p6 165 5.90 1.40 1170 Radon [Xe]6s24f145d106p6 211 9.73 1.50 1037

Problem 87AE Elements in group 7 A in the periodic table are called the halogens; elements in group 6A are called the chalcogens (a) What is the most common oxidation state of the chalcogens compared to the halogens? (b) For each of the following periodic properties, state whether the halogens or the chalcogens have larger values: atomic radii, ionic radii of the most common oxidation state, first ionization energy, second ionization energy.

The \(\mathrm{As}-\mathrm{As}\) bond length in elemental arsenic is \(2.48 \AA\). The \(\mathrm{Cl}-\mathrm{Cl}\) bond length in \(\mathrm{Cl}_{2} \text { is } 1.99 \AA\). (a) Based on these data, what is the predicted \(\mathrm{As}-\mathrm{Cl}\) bond length in arsenic trichloride, \(\mathrm{AsCl}_{3}\), in which each of the three \(\mathrm{Cl}\) atoms is bonded to the As atom? (b) What bond length is predicted for \(\mathrm{AsCl}_{3}\), using the atomic radii in Figure 7.7? Equation Transcription: Å Å Text Transcription: As-As 2.48 \AA Cl-Cl Cl2 is 1.99 \AA As-Cl AsCl3 Cl AsCl3

Problem 86AE The following observations are made about two hypothetical elements A and B: The A—A and B—B bond lengths in elemental A and B are 2.36 and 1.94 A. respectively. A and B react to form the binary compound AB2. which has a linear structure (that is^B-A-B = 180°). Based on these statements, predict the separation between the two B nuclei in amolecule of AB2.

Note from the following table that there is a significant increase in atomic radius upon moving from \(Y\) to La whereas the radii of \(\mathrm{Zr}$ to $\mathrm{Hf}\) are the same. Suggest an explanation for this effect. Equation Transcription: Text Transcription: Y Zr to Hf

Problem 89AE (a) Which ion is smaller, Co3+ or Co+? (b) In a lithium-ion battery that is discharging to power a device, for every Li+ that inserts into the lithium cobalt oxide electrode, a Co4+ ion must be reduced to a Co3+ ion to balance charge. Using the CRCHandbook of Chemistry and Physics or other standard reference, find the ionic radii of Li+, Co3+, and Co4+. Order these ions from smallest to largest, (c) Will the lithium cobalt electrode expand or contract as lithium ions are inserted? (d) Lithium is not nearly as abundant as sodium. If sodium ion batteries were developed that function as lithium ion ones, do you think “sodium cobalt oxide" would still work as the electrode material? Explain, (e) If you don't think cobalt would work as the redox-active partner ion in the sodium version of the electrode, suggest an alternative metal ion and explain your reasoning.

The ionic substance strontium oxide, \(\mathrm{SrO}\), forms from the reaction of strontium metal with molecular oxygen. The arrangement of the ions in solid \(\mathrm{SrO}\) is analogous to that in solid \(\mathrm{NaCl}\): (a) Write a balanced equation for the formation of \(\mathrm{SrO}(\mathrm{s})\) from its 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 corner to the center of an atom at a neighboring corner). (c) The density of \(\\mathrm{SrO} \text { is } 5.10 \mathrm{~g} / \mathrm{cm}^{3}). Given your answer to part (b), how many formula units of \(\mathrm{SrO}\) are contained in the cube shown here? Equation Transcription: Text Transcription: SrO SrO NaCl SrO(s) SrO is 5.10 g/cm3 SrO

Group 4A elements have much more negative electron affinities than their neighbors in groups 3A and 5A (see Figure 7.12). Which of the following statements best explains this observation? (i) The group 4A elements have much higher first ionization energies than their neighbors in groups 3A and 5A. (ii) The addition of an electron to a group 4A element leads to a half-filled \(n p^{3}\) outer electron configuration. (iii) The group 4A elements have unusually large atomic radii. (iv) The group 4A elements are easier to vaporize than are the group 3A and 5A elements. Equation Transcription: Text Transcription: n p^{3}

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

Problem 94AE (a) Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons. (b) Why does 03 - not exist?

Problem 93AE 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) 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?

Problem 96AE Identify two ions that have the following ground-state electron configurations: (a) [Ar], (b) [Ar]3d5, (c) [Kr]5s24c(10

Problem 95AE 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.

Problem 97AE Which of the following chemical equations is connected to the definitions of (a) the first ionization energy of oxygen, (b) the second ionization energy of oxygen, (c) the electron affinity of oxygen? (i) O(g) + e-?O-(g) (ii) O(g)?O+(g) + e- (iii) O(g)+ 2e-?O2-(g) (iv) O(g)?O2+(g) + 2e-(v) O+(g)?O2+(g) + e–

The electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group 1B and group 2B metals are as follows: (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 affinities of other groups as we proceed down the periodic table.] Equation Transcription: Text Transcription: kJ/mol Cu-119 Zn>0 Ag-126 Cd>0 Au-223 Hg>0

Problem 99AE Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. 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, (d) The hydride ion is H-. Write out the process corresponding to the first ionization energy of the hydride ion. (e) How does the process in part (d) compare to the process for the electron affinity of a neutral hydrogen atom?

Problem 100AE The first ionization energy of the oxygen molecule is the energy required for the following process: 02(g)—>02+(g) + e- The energy needed for this process is 1175 kJ/mol, very similarto the first ionization energy of Xe. Would you expect 02 to react with F2? If so. suggest a product or products of this reaction.

Problem 101AE It is possible to define metallic character as we do in this book and base it on the reactivity of the element and the ease with which it loses electrons. Alternatively, one could measure how well electricity is conducted by each of the elements to determine how “metallic" the elements are. On the basis of conductivity, there is not much of a trend in the periodictable: Silver is the most conductive metal, and manganese theleast. Look up the first ionization energies of silver and manganese: which of these two elements would you call more metallic based on the way we define it in this book?

Problem 102AE Which of the following is the expected product of the reaction of K(s) and H2(g)? (i) KH(s), (ii) K2H(s), (iii) KH2(s), (iv) K2H2(s), or (v) K(s) and H2(g) will not react with one another

Problem 103AE Elemental cesium reacts more violently with water than does elemental sodium. Which of the following best explains this difference in reactivity? (i) Sodium has greater metallic character than does cesium, (ii) The first ionization energy of cesium is less than that of sodium, (iii) The electron affinity of sodium is smaller than that of cesium, (iv) The effective nuclear charge for cesium is less than that of sodium, (v) The atomic radius of cesium is smaller than that of sodium.

Problem 104AE (a) 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? (b) Write a balanced chemical equation for the reaction of the white substance with water.

Problem 105AE Zinc in its 2+ oxidation state is an essential metal ion for life. Zn2+ is found bound to many proteins that are involved in biological processes, but unfortunately Zn2+ is hard to detect by common chemical methods. Therefore, scientists who are interested in studying Zn2+-containing proteins frequently substitute Cd2+ for Zn2+, since Cd2+ is easier to detect. (a) On the basis of the properties of the elements and ions discussed in this chapter and their positions in the periodic table, describe the pros and cons of using Cd2+ as a Zn2+ substitute (b) Proteins that speed up (catalyze) chemical reactions are called enzymes. Many enzymes are required for proper metabolic reactions in the body. One problem with using Cd2+ to replace Zn2+ in enzymes is that Cd2+ substitution can decrease or even eliminate enzymatic activity. Can you suggest a different metal ion that might replace Zn2+ in enzymes instead of Cd2+? Justify your answer.

A historian discovers a nineteenth-century notebook in which some observations, dated 1822, were recorded on a substance thought to be a new element. Here are some of the data recorded in the notebook: “Ductile, silver-white, 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 the CRC Handbook of Chemistry and Physics, 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 7.1, might this nineteenth-century investigator have been the first to discover a new element? Equation Transcription: Text Transcription: 153°C 7.3 g/cm3 20% 1% 4.20 g 5.08 g 800°C

Problem 108AE We will see in Chapter that semiconductors are materials that conduct electricity better than nonmetals but not as well as metals. The only two elements in the periodic table that are technologically useful semiconductors are silicon and germanium. Integrated circuits in computer chips today are based on silicon. Compound semiconductors are also used in the electronics industry. Examples are gallium arsenide. GaAs: gallium phosphide, GaP; cadmium sulfide. CdS; and cadmium selenide. CdSe. (a) What is the relationship between the compound semiconductors' compositions and the positions of their elements on the periodic table relative to Si and Ge? (b) Workers in the semiconductor industry refer to “II—VI” and “III—V” materials, using Roman numerals. Can you identify which compound semiconductors are II—VI and which are III—V? (c) Suggest other compositions of compound semiconductors based on the positions of their elements in the periodic table.

Problem 107AE In April 2010, a research team reported that it had made Element 117. This discovery was confirmed in 2012 by additional experiments. Write the ground-state electron configuration for Element 117 and estimate values for its first ionization energy, electron affinity, atomic size, and common oxidation state based on its position in the periodic table.

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: (a) Calculate the frequency, \(v\), of the X rays emitted by each of the elements, in Hz. (b) Plot the square root of 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 \(0.980 \AA\). What element do you think it is? Equation Transcription: Å Text Transcription: v 0.980 \AA

(a) Write the electron configuration for \(L i \) 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 \(\left(-2.18 \times 10^{-18} J\right)\left(\frac{Z^{2}}{n^{2}}\right)\), where Z is the nuclear charge and n is the principal quantum number of the electron. Estimate the first ionization energy of \(L i\) (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 part (c)? Equation Transcription: Text Transcription: Li (-2.18 X 10^-18J) (Z^2 over n^2) Li

One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect. (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 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 \(10.75 \mathrm{eV}\). What is the first ionization energy of \(\mathrm{Hg}, \text { in } \mathrm{kJ} / \mathrm{mol}\)? (d) Using Figure 7.10, determine which of the halogen elements has a first ionization energy closest to that of mercury. Equation Transcription: Text Transcription: 58.4 nm 10.75eV Hg, in kJ/mol

Problem 112IE Mercury in the environment can exist in oxidation states 0, +1, and +2. One major question in environmental chemistry research is how to best measure the oxidation state of mercury in natural systems; this is made more complicated by the fact that mercury can be reduced or oxidized on surfaces differently than it would be if it were free in solution. XPS, X-ray photoelectron spectroscopy, is a technique related to PES (see Exercise), but instead of using ultraviolet light to eject valence electrons, X rays are used to eject core electrons. The energies of the core electrons are different for different oxidation states of the element. In one set of experiments, researchers examined mercury contamination of minerals in water. They measured the XPS signals that corresponded to electrons ejected from mercury's 4f orbitals at 105 eV, from an X-ray source that provided 1253.6 eV of energy. The oxygen on the mineral surface gave emitted electron energies at 531 eV, corresponding to the 1s orbital of oxygen. Overall the researchers concluded that oxidation states were +2 for Hg and -2 for O (a) Calculate the wavelength of the X rays used in this experiment (b) Compare the energies of the 4f electrons in mercury and the 1s electrons in oxygen from these data to the first ionization energies of mercury and oxygen from the data in this chapter (c) Write out the ground-state electron configurations for Hg2+ and 02-; which electrons are the valence electrons in each case (d) Use Slater's rules to estimate Zeff for the 4f and valence electrons of Hg2+ and 02-; assume for this purpose that all the inner lectrons with (n - 3) or less screen a full +1. Exercise One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect, coo (Section) 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 energyof 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) Using Figure determine which of the halogen elements has a first ionization energy closest to that of mercury.

(a) The measured \(\mathrm{Bi} \neg \mathrm{Br}\) bond length in bismuth tribromide, \(\mathrm{BiBr}_{3} \text {, is } 2.63 \AA\). Based on this value and the data in Figure 7.8, predict the atomic radius of \(\mathrm{Bi}\). (b) Bismuth tribromide is soluble in acidic solution. It is formed by treating solid bismuth(III) oxide with aqueous hydrobromic acid. Write a balanced chemical equation for this reaction. (c) While bismuth(III) oxide is soluble in acidic solutions, it is insoluble in basic solutions such as \(\mathrm{NaOH}(a q)\). Based on these properties, is bismuth characterized as a metallic, metalloid, or nonmetallic element? (d) Treating bismuth with fluorine gas forms \(\mathrm{BiF}_{5}\). Use the electron configuration of Bi to explain the formation of a compound with this formulation. (e) While it is possible to form \(\mathrm{BiF}_{5}\) in the manner just described, pentahalides of bismuth are not known for the other halogens. Explain why the pentahalide might form with fluorine but not with the other halogens. How does the behavior of bismuth relate to the fact that xenon reacts with fluorine to form compounds but not with the other halogens? Equation Transcription: Å Text Transcription: Bi \neg Br BiBr_3 is 2.63 \AA Bi NaOH(aq) BiF5 BiF5

When magnesium metal is burned in air (Figure 3.6), two products are produced. One is magnesium oxide, \(\mathrm{MgO}\). 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 \(\mathrm{MgO}\) 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 \(0.486 \mathrm{~g} \text { of } \mathrm{MgO}\). 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 completion, 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. Equation Transcription: Text Transcription: MgO MgO 0.470 g 0.486 g of MgO 6.3-g Mg 2.57 g NH_3(g) H_2(g) -461.08 kJ/mol

Problem 115IE Potassium superoxide. K02, 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 002(g). (a) The products of the reaction are K2C03(s) and 02(g). Write a balanced equation for the reaction between K02(sj and C02(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(sj is needed to consume 18.0 g 002(g)? What mass of 02(g) is produced during this reaction?