Chemistry part 12, Julia Burdge,2e (2009) potx

266 CHAPTER 7 Electron Configuration and the Periodic Table Type of compound Structure Melting point (0C) Boiling point (0C) Acid-base nature Certain oxides such as CO and NO are neutra l; that is, they do not react with water to produce acidic or basic solutions. In general, oxides of nonmetals are either acidic or neutral. Na20 MgO AI 2 0 3 Si0 2 P4 0 10 50 3 (1 2 0 7 ~ • 1275 ? • Basic Ionic Molecular • Extensive three-dimensional • Discrete molecular units 3 2800 2045 1610 580 16.8 -91.5 3600 2980 2230 ? 44.8 82 • Basic Amphoteric ~ Acidic • (Si0 2 ) also has a huge three-dimensional network, although it is not an ionic compound. The oxides of phosphorus, sulfur, and chlorine are molecular compounds composed of small discrete units. The weak attractions among these molecules result in relatively low melting points and boiling points. Most oxides can be classified as acidic or basic depending on whether they produce acids or bases when dissolved in water (or whether they react as acids or bases). Some oxides are ampho- teric, which means that they display both acidic and basic properties. The first two oxides of the third period, Na20 and MgO, are basic oxides. For example, Na20 reacts with water to form the base sodium hydroxide: Na20(S) + H 2 0(l) +. 2NaOH(aq) Magnesium oxide is quite insoluble; it does not react with water to any appreciable extent. How- ever, it does react with acids in a manner that resembles an acid-base reaction: MgO(s) + 2HCI(aq) +. MgCI 2 (aq) + H 2 0(I) The products of this reaction are a salt (MgCI 2 ) and water, the same kind of products that are obtained in an acid-base neutralization. Aluminum oxide is even less soluble than magnesium oxide. It, too, does not react with water, but it exhibits the properties of a base by reacting with acids: It also exhibits acidic properties by reacting with bases: Thus, Al 2 0 3 is classified as an amphoteric oxide because it has properties of both acids and bases. Other amphoteric oxides are ZnO, BeO, and Bi 2 0 3 . Silicon dioxide is insoluble and does not react with water. It has acidic properties, however, because it reacts with a very concentrated aqueous base: Si0 2 (s) + 20H-(aq) +. SiO~-(aq) + H 2 0(l) For this reason, concentrated aqueous, strong bases such as sodium hydroxide (NaOH) should not be stored in Pyrex glassware, which is made of Si0 2 . The remaining third-period oxides (P 4 0 IO , S03, and C1 2 0 7 ) are acidic. They react with water to form phosphoric acid, sulfuric acid, and perchloric acid, respectively: P 4 0 IO (S) + 6H 2 0(aq) +. 4H 3 P0 4 (aq) S03(g) + H 2 0(l) • 2H 2 S0 4 (aq) . . . . . . . . . . . , . CI 2 0il) + H 2 0(l) +. 2HCIOiaq) This brief examination of oxides of the third-period elements shows that as the metallic character of the elements decreases from left to right across the period, their oxides change from basic to amphoteric to acidic. Metal oxides are usually basic, and most oxides of nonmetals are acidic. The intermediate properties of the oxides (as demonstrated by the amphoteric oxides) are exhibited by elements whose positions are intermediate within the period. Because the metallic character of the elements increases from top to bottom within a group of main group elements, the oxides of elements with higher atomic numbers are more basic than the lighter elements. APPLYING WHAT YOU'VE LEARNED 267 Applying What You've Learned In 1949, the Australian psychiatrist John Cade published the results of his studies show- ing that lithium was useful in the treatment of "manic episodes," one of the phases of what is known today as bipolar disorder. Although the research had shown real prom- ise, its publication coincided with news of the lithium poisoning and resulting deaths of a group of cardiac patients that had used lithium chloride as a dietary salt substitute. Reports of this disaster prompted drug manufacturers to withdraw all lithium salts from the market, and for a time, any medical use of lithium was viewed as too dangerous even to consider. Additional research in Europe and the United States resulted in the gradual acceptance of lithium as a potentially valuable psychiatric therapy. The FDA approved lithium carbonate (Li 2 C0 3 ) in 1970 for the treatment of manic illness, and in 1974 for the treatment of bipolar disorder. - Problems: a) Without referring to a periodic table, write the electron configuration of lithium (Z = 3). [ ~~ Sample Problem 7.2] b) Referring only to a periodic table, arrange Li and the other alkali metals (not in- cluding Fr) in order of increasing atomic radius. [ ~~ Sample Problem 7.3] c) Again referring only to the periodic table, arrange the members of Group 1A (not including Fr) in order of increasing ionization energy (lEI). [ ~ Sample Problem 7.4] d) Write the electron configuration for each of the alkali metal cations. [ ~~ Sample Prob l em 7.6] e) For each alkali metal cation in part (d), identify an isoelectronic series consisting ;;;- • , of a noble gas and, where appropriate, one or more common ions (see Figure 2.14). [ ~~ Sample Problem 7.7] 268 CHAPTER 7 Electron Configuration and the Periodic Table CHAPTER SUMMARY Section 7.1 o The modern periodic table was devised independently by Dmitri Mendeleev and Lothar Meyer in the nineteenth century. The elements that were known at the time were grouped based on their physical and chemical properties. Using his arrangements of the elements, Mendeleev succe ssf ully predicted the existence of elements that had not yet been di scovered. o o Early in the twentieth century, Henry Mo seley refined the periodic table with the concept of the atomic number, thus resolving a few inconsistencies in the tables proposed by Mendeleev and Meyer. Elements in the same group of the periodic table tend to have similar physical and chemical propertie s. Section 7.2 o The periodic table can be divided into the main group elements (also known as the representative elements) and the transition metals. It is further divided into smaller groups or columns of elements that all have the same configuration of valence electrons. o The 18 columns of the periodic table are labeled 1A through 8A (s - and p-block elements) and I B through 8B (d -block elements), or by the numbers 1 through 18. Section 7.3 o Effective nuclear charge (ZerrJ is the nuclear charge that is "felt" by the valence electrons. It is usually lower than the nuclear charge due to shielding by the core electrons. • According to Coulomb's law, the attractive force (F) between two oppositely charged particles (QI and Q 2) is directly proportional to the product of the charges and inversely proportional to the distance (d) between the objects squared: (F oc QJ . Q 21d\ Section 7.4 o Atomic radius is the distance between an atom's nucleus and its valence shell. The atomic radius of a metal atom is defined as the metallic radius, which is one-half the distance between adjacent, identical nuclei in a metal solid. The atomic radius of a nonmetal is defined as the covalent radius, which is one-half the distance between adjacent, identical nuclei in a molecule. In general, atomic radii decrease from left to right across a period of the periodic table and increase from top to bottom down a group. o Ionization energy (IE) is the energy required to remove an electron from an atom. The first ionization energy (IE I ) is smaller than subsequent ionization energies [e.g., second (lE 2 ), third (IE 3)' and so on]. The first ionization of any atom removes a valence electron. Ionization energies increase dramatically when core electrons are being removed. KEyWORDS o First ionization energies (IEI values) tend to increase across a period and decrease down a group. Exceptions to this trend can be explained based upon the electron configuration of the element. o Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron. EA is equal to -I1H for the process A(g) + e - • A - (g). o Electron affinities tend to increase across a period. As with first ionization energies, exceptions to the trend can be explained based on the electron configuration of the element. o Metals tend to be shiny, lustrous, malleable, ductile, and conducting (for both heat and electricity). Metals typically lose electrons to form cations, and they tend to form ionic compounds (including ba sic oxides). o Nonmetals tend to be brittle and not good conductors (for either heat or electricity). They can gain 'electrons to form anions but they commonly form molecular compounds (including acidic oxides). o In general, metallic character decreases across a period and increases down a group of the periodic table. Metalloids are elements with properties intennediate between metals and nonmetals. Section 7.5 o Ions of main group elements are isoelectronic with noble gases. When a d-block element loses one or more electrons, it loses them first from the shell with the highest principal quantum number (e.g., electrons in the 4s subshell are lost before electrons in the 3d subshe ll ). Section 7.6 o Ionic radius is the distance between the nucleus and valence shell of a cation or an anion. A cation is smaller than its parent atom. An anion is larger than its parent atom. o An isoelectronic series consists of one or more ions and a noble gas, all of which have identical electron configurations. Within an isoelectronic serie s, the greater the nuclear charge, the smaller the radius. Section 7.7 o Although members of a group in the periodic table exhibit similar chemical and physical properties, the first member of each group tends to be significantly different from the other members. Hydrogen is essentially a group unto itself. o The alkali metals (Group lA ) tend to be highly reactive toward oxygen, water, and acid. Group 2A metals are less reactive than Group lA metals, but the heavier members a ll react with water to produce metal hydroxides and hydrogen gas. Groups that contain both metals and nonmetals (e.g., Groups 4A, SA, and 6A) tend to show greater variability in their physical and chemical properties. o Amphoteric oxide s, such as A1 2 0 3 , are those that exhibit both acidic and basic behavior. Amphoteric, 266 Atomic radiu s, 246 Coul omb's law, 245 Covalent radiu s, 246 Effective nuclear charge (Zeff)' 244 Isoelectronic, 253 Metalloid, 252 Shielding, 244 Va lence electrons, 243 Electron affinity (EA), 250 Isoelectronic series, 255 Ionic radius, 254 Main group elements, 241 Ionization energy (IE), 247 Metallic radius, 246 QUESTIONS AND PROBLEMS 269 KEY EQUATION 7.1 Z eff = Z - (J" QUESTIONS AND PROBLEMS ================ ==========~~ Section 7.1: Development of the Periodic Table Review Questions 7.1 Briefly describe the significance of Mendeleev's periodic table. 7.2 What is Moseley's contribution to the modern periodic table? 7.3 Describe the general layout of a modern periodic table. 7.4 What is the most important relationship among elements in the same group in the periodic table? Section 7.2: The Modern Periodic Table Review Questions 7.5 Classify each of the following elements as a metal, a nonmetal, or a metalloid: As, Xe, Fe, Li, B, Cl, Ba, P, I, Si. 7.6 Compare the physical and chemical properties of metals and nonmetals. 7.7 7.S 7.9 7.10 Draw a rough sketch of a periodic table (no details are required). Indicate regions where metals, nonmetals, and metalloids are located. What is a main group element? Gi ve names and symbols of four main group elements. Without referring to a periodic table, write the name and give the symbol for one element in each of the following groups: 1 A, 2A , 3A, 4A, SA, 6A, 7 A, SA, transition metals. Indicate whether the following elements exist as atomic species, molecular species, or extensive three-dimensional structures in their most stable states at room temperature, and write the molecular or empirical formula for each one: phosphorus, iodine, magnesium, neon, carbon, sulfur, ce sium, and oxygen. 7.11 You are given a sample of a dark, shiny solid and asked to determine whether it is the nonmetal iodine or a metallic element. What test could you do that would enable you to answer the question without destroying the sample? 7.12 What are valence electrons? For main group elements, the number of valence electrons of an element is equal to its group number. Show that this is true for the following elements: AI , Sr , K, Br, P, S, C. 7.13 Write the outer electron configurations for the (a) alkali metals, (b) alkaline earth metals, (c) halogens, (d) noble gases. 7.14 Use the first-row transition metals (Sc to Cu) as an example to illustrate the characteristics of the electron configurations of transition metals. 7.15 Arsenic is not an essential element for the human body. Ba s ed on its position in the periodic table, suggest a reason for its toxicity. Problems 7.16 7.17 7.1S 7.19 7.20 7.21 In the periodic table, the element hydrogen is sometimes grouped with the alkali metals and sometimes with the halogens. Explain why hydrogen can resemble the Group 1A and the Group 7 A elements. A neutral atom of a certain element has 16 electrons. Consulting only the periodic table, identify the element and write its ground- state electron configuration. Group the following electron configurations in pairs that would represent elements with similar chemical properties: (a) I s2 2/ 2p 6 3/ (b) Is22i2 p 3 (c) 1/ 2i 2l3 s2 3p64s2 3d J0 4p6 (d) li 2/ (e) 1/2 i 2p 6 (f) li 2s22p 6 3i 3 p 3 Group the following electron configurations in pairs that would represent elements with similar chemical properties: (a) 1i 2s 22 p 5 (b) l i 2s1 (c) li 2/2 p 6 (d) li2i2p 6 3s 2 3 p 5 (e) li2i2p 6 3i 3 p 64s1 (f) ls 22i 2p 6 3i3l4/3d 1 04 p 6 Without referring to a periodic table, write the electron configuration of elements with the following atomic numbers: (a) 9, (b) 20 , (c) 26, (d) 33. Specify the group of the periodic table in which each of the following elements is found: (a) [Ne]3s l , (b) [Ne]3s 2 3 p 3, (c) [Ne]3s 2 3 p 6, (d) [Ar]4s 2 3d 8 . Section 7.3: Effective Nuclear Charge Review Questions 7.22 Explain the term effective nuclear charge. 7.?3 Explain why the atomic radius of Be is smaller than that of Li. Problems 7.24 The electron configuration of B is ls22i2p I. (a) If each core electron (that is, the Is electrons) were totally effective in shielding the valence electrons (that i s, the 2s and 2p electrons) from the nucleus and the valence electrons did not shield one another, what would be the shielding constant (0') and the effective nuclear charge, (Z eff) for the 2s and 2p electrons? (b) In reality, the shielding constants for the 2s and 2p electrons in B are slightly different. They are 2.42 and 2.SS, respectively. Calculate Z eff for these electrons, and explain the differences from the values you determined in part (a). 270 CHAPTER 7 Electron Configuration and the Periodic Table 7.25 The electron configuration of Ci s 1 i2i 2p 2. (a) If each core electron (that is, the Is electrons) were totally effective in' screening the valence electrons (that is, the 2s and 2p electrons) from the nucleus and the valence electrons did not shield one another, what would be the shielding constant ( u) and the effective nuclear charge, (Z eff ) for the 2s and 2p electrons? (b) In reality, the shielding constants for the 2s and 2p electrons in C are slightly different. They are 2.78 and 2.86, respectively. Calculate Ze ff for these electrons, and explain the differences from the values you determined in part (a). Section 7.4: Periodic Trends in Properties of Elements Review Questions 7.26 Define atomic radius. Doe s the size of an atom have a precise meaning? 7.27 How does atomic radius change (a) from left to right across a period and (b) from top to bottom in a group? Define ionization energy. Explain why ionization energy measurements are usually made when atoms are in the ga seous state. Why is the secQnd ionization energy always greater than the first ionization energy for any element? 7.29 Sketch the outline of the periodic table, and show group and period trends in the first ionization energy of the elements. What types of elements have the highest ionization energies and what types the lowest ionization energies? 7.30 (a) Define electron affinity. (b) Explain why electron affinity measurements are made with gaseous atoms. (c) Ionization energy is always a positive quantity, whereas electron affinity may be either positive or negative. Explain. 7.31 Explain the trends in electron affinity from aluminum to chlorine (see Figure 7.lO). Problems 7.32 On the basis of their positions in the periodic table, select the atom with the larger atomic radius in each of the following pairs: (a) Na, Si; (b) Ba, Be; (c) N, F; (d) Br, CI; (e) Ne, Kr. 7.33 Arrange the following atoms in order of increasing atomic radius: Na, AI, P, CI, Mg. 7.34 Which is the largest atom in the third period of the periodic table? 7.35 Which is the smallest atom in Group 7 A? 7.36 Based on size, identify the spheres shown as Na, Mg , 0 , and S. 7.37 7.38 7.39 Based on size, identify the spheres shown as K, Ca , S, and Se. • Why is the radius of the lithium atom considerably larger than the radius of the hydrogen atom? Use the second period of the periodic table as an example to show that the size of atoms decreases as we move from left to right. Explain the trend. 7.40 7.41 7.42 7.43 7.44 7.45 7.46 7.47 7.48 7.49 7.51 Arrange the following in order of increasing first ionization energy: Na, CI, AI, S, and Cs. Arrange the following in order of increasing first ionization energy: F, K, P, Ca, and Ne. Use the third period of the periodic table as an example to illustrate the change in first ionization energies of the elements as we move from left to right. Explain the trend. In general, the first ionization energy increases from left to right across a given period. Aluminum, however, has a lower first ionization energy than magnesium. Explain. The first and second ionization energies of K are 419 and 3052 kl lmol, and those of Ca are 590 and 1145 kllmol, respectively. Compare their values and comment on the differences. Two atoms have the electron configurations 1i2i2 p 6 and IS 22i 2l3s 1 • The first ionization energy of one is 2080 kJ/mol, and that of the other is 496 kJ/mo!. Match each ionization energy with one of the given electron configurations. Justify your choice. A hydrogen-like ion is an ion containing only one electron. The energies of the electron in a hydrogen-like ion are given by E = -(218 X lO-18 J )Z 2 ( L 1/ · 2 n where n is the principal quantum number and Z is the atomic number of the element. Calculate the ionization energy (in kl lmol) of the He + ion. Plasma is a state of matter consisting of positive gaseous ions and electrons. In the plasma state, a mercury atom could be stripped of its 80 electrons and therefore would exist as Hg8o +. Use the equation in Problem 7.46 to calculate the energy required for the last ionization step, that is, Hg 79 +(g) • Hg 8o+ (g) + e- Arrange the elements in each of the following groups in order of increasing electron affinity: (a) Li, Na, K; (b) F, CI, Br, 1. Specify which of the following elements you would expect to have the greatest electron affinity: He, K, Co, S, Cl. Considering their electron affinities, do you think it is possible for the alkali metals to form an anion like M- , where M represents an alkali metal? Explain why alkali metals have a greater affinity for electrons than alkaline earth metals. Section 7.5: Electron Configuration of Ions Review Questions How does the electron configuration of ions derived from main group elements give them stability? 7.53 What do we mean when we say that two ions or an atom and an ion are isoelectronic? 7.54 7.55 Is it possible for the atoms of one element to be isoelectronic with the atoms of another element? Explain. Give three examples of first-row transition metal (Se to Cu) ions that are isoelectronic with argon. Problems 7.56 A M 2 + ion derived from a metal in the first transition metal series has four electrons in the 3d subshell. What element might M be ? 7.57 A metal ion with a net + 3 charge has five electrons in the 3d subshell. Identify the metal. 7.58 Write the ground-state electron configurations of the following ions: (a) Li+, (b) W, (c) N 3 - , (d) F- , (e) S 2- , (f) Al 3 +, (g) Se 2 -, (h) Br - , (i) Rb +, (j) Sr 2 +, (k) Sn 2+ , (I) Te 2 - , (m) Ba 2 +, (n) Pb 2+ , (0) In 3 +, (p) Tl +, (q) T1 3 + 7.59 Write the ground-state electron configurations of the following ions, which play important roles in biochemical processes in our bodies: (a) Na +, (b) Mg2+, (c) Cl - , (d) K+, (e) Ca 2 +, (f) Fe 2 +, (g) Cu 2+ , (h) Zn 2+ . 7.60 Write the ground-state electron configurations of the following transition metal ions: (a) Sc 3+ , (b) Ti 4 +, (c) V5+, (d) Cr 3+ , (e) Mn 2+, (f) Fe 2+ , (g) Fe 3+ , (h) Co 2+ , (i) Ni 2+ , U) Cu+, (k) Cu 2 +, (1) Ag+, (m) Au+, (n) Au3+, (0) Pt U 7.61 Name the ions with three charges that have the following electron configurations: (a) [ArJ3d 3 , (b) [Ar], (c) [KrJ4d 6 , (d) [XeJ4/45d 6 . 7.62 Which of the following species are isoelectronic with each other: C, CI- , Mn2+, B-, Ar, Zn, Fe 3+ , Ge 2+ ? 7.63 Group the species that are isoelectronic: Be 2 +, F- , Fe 2+ , N 3 - , He, S2- , Co 3+ , Ar. 7.64 Thallium (Tl) is a neurotoxin and exists mostly in the TI(I) oxidation state in its compounds. Aluminum (AI), which causes anemia and dementia, is only stable in the Al(ill) form. The first, second, and third ionization energies ofTI are 589, 1971, and 2878 kJ/mol, respectively. The first, second, and third ionization energies of Al are 577.5,1817, and 2745 kJ/mol, respectively. Plot the ionization energies of Al and TI versus atomic number and explain the trends. Section 7.6: Ionic Radius Review Questions 7.65 Define ionic radius. How does the size of an atom change when it is converted to (a) an anion and (b) a cation? 7.66 Explain why, for isoelectronic ions, the anions are larger than the cations. Problems 7.67 7.68 7.69 7.70 Indicate which one of the two species in each of the following pairs is smaller: (a) CI or CC (b) Na or Na +; (c) 0 2 - or S 2- ; (d) Mg 2+ or Al 3+ ; (e) Au + or Au 3 + List the following ions in order of increasing ionic radius: N 3 -, N + F- M 2+ 0 2 - a, , g, . Explain which of the following cations is larger, and why: Cu + orCu 2 + Explain which of the following anions is larger, and why: Se 2 - or Te 2 - . QUESTIONS AND PROBLEMS 271 7.71 Both Mg 2+ a nd Ca 2 + are important biological ions. One of their functions is to bind to the phosphate group of ATP molecules or amino acids of proteins. For Group 2A metals in general, the tendency for binding to the anions increases in the order Ba 2+ < Sr 2+ < Ca 2+ < Mg 2+ . Explain this trend. Section 7.7: Periodic Trends in Chemical Properties of the Main Group Elements Review Questions 7.72 Why do member s of a group exhibit similar chemical properties? 7.73 Which elements are more likely to form acidic oxides? Basic oxides? Amphoteric oxides? Problems 7.74 Give the physical states (gas, liquid, or solid) of the main group elements in the fourth period (K, Ca, Ga, Ge, As, Se , Br) at room temperature. 7.75 The boiling points of neon and krypton are -245.9 °C and - 152.9°C, respectively. Using these data, estimate the boiling point of argon. 7.76 Use the alkali metals and alkaline earth metals as examples to show how we can predict the chemical properties of elements simply from their electron configurations. 7.77 Based on your knowledge of the chemistry of the alkali metals, predict some of the chemical properties of francium, the last member of the group. 7.78 As a group, the noble gases are very stable chemically (only Kr and Xe are known to form compounds). Why? 7.79 Why are Group 1B elements more stable than Group 1A elements even though they seem to have the same outer electron configuration, nsl, where n is the principal quantum number of the outermost shell? 7.80 How do the chemical properties of oxides change from left to right across a period? How do they change from top to bottom within a particular group? 7.81 7.82 Wri te balanced equations for the reactions between each of the following oxide s and water: (a) Li 2 0 , (b) CaO, (c) S0 3' Write formulas for and name the binary hydrogen compounds of the second-period elements (Li to F). Describe how the physical and chemical properties of these compounds change from left to right across the period. 7.83 Which oxide is more basic, MgO or BaO? Why? Additional Problems 7.84 State whether each of the following properties of the main group elements generally increases or decreases (a) from left to right acro ss a period and (b) from top to bottom within a group: metallic character, atomic size, ionization energy, acidity of oxides. 7.85 Referring to the periodic table, name (a) the halogen in the fourth period, (b) an element similar to phosphorus in chemical properties, ( c) the most reactive metal in the fifth period, (d) an element that has an atomic number smaller than 20 and is similar t6 strontium. 272 CHAPTER 7 Electron Configuration and the Periodic Table 7.86 7.87 7.88 7.89 7.90 7.91 Write equations representing the following processes: (a) The electron affinity of S- (b) The third ionization energy of titanium (c) The electron affinity of Mg 2+ (d) The ionization energy of 0 2 - An'ange the following isoelectronic species in order of increasing . 0 2 - F- N + M 2+ lOTIlZatlOn energy: , , a, g . Write the empirical (or molecular) formulas of compound s that the elements in the third period (sodium to chlorine) should form with (a) molecular oxygen and (b) molecular chlorine. In each case indicate whether you would expect the compound to be ionic or molecular in character. Element M is a shiny and highly reactive metal (melting point 63°C), and element X is a highly reactive nonmetal (melting point -7.2 °C). They react to form a compound with the empirical formula MX, a colorless, brittle white solid that melts at 734 °C. When dissolved in water or when in the molten state, the substance conducts electricity. When chlorine gas is bubbled through an aqueous solution containing MX , a reddish-brown liquid appears and CC ions are formed. From these observations, identify M and X. (You may need to consult a handbook of chemistry for the melting-point values.) Match each of the elements on the right with its description on the left: (a) A dark-red liquid (b) A colorless gas that bums in oxygen gas (c) A metal that reacts violently with water (d) A shiny metal that is used in jewelry (e) An inert gas Calcium (Ca ) Gold (Au) Hydrogen (H 2 ) Argon (Ar ) Bromine (Brz) Arrange the following species in isoelectronic pairs: 0 +, Ar , S2- , Ne, Zn, Cs +, N 3 -, AS H, N, Xe. 7.92 In which of the following are the species written in decreasing order by size ofradius? (a) Be , Mg , Ba, (b) N 3 - , 0 2 - , F- , (c) TI H, TIZ +, Tl+. 7.93 Which of the following properties show a clear periodic variation: (a) first ionization energy, (b) molar mass of the elements, (c) number of isotopes of an element , (d) atomic radius? 7.94 When carbon dioxide is bubbled through a clear calcium hydroxide solution, the solution appears milky. Write an equation for the reaction, and explain how this reaction illustrates that CO 2 is an acidic oxide. 7.95 7.96 7.97 You are given four substances: a fuming red liquid, a dark metallic-looking solid, a pale-yellow gas, and a yellow-green gas that attacks glass. You are told that these substances are the first four members of Group 7 A, the halogens. Name each one. For each pair of elements listed, give three properties that show their chemical similarity: (a) sodium and potassium and (b) chlorine and bromine. Name the element that forms compounds , under appropriate conditions, with every other element in the periodic table except He, Ne, and Ar. 7.98 Explain why the first electron affinity of sulfur is 200 kJ/mol but the second electron affinity is - 649 kJ/mo!. 7.99 The H- ion and the He atom have two Is electrons each. Which of the two species is larger? Explain. 7.100 Predict the products of the following oxides with water: Na zO, BaO, CO 2 , NzOs, P 4 0 10 , S0 3' Write an equation for each of the reactions. Specify whether the oxides are acidic, basic, or amphoteric. 7.101 Write the formulas and name s of the oxides of the second- period elements (Li to N). Identify the oxides as acidic, basic, or amphoteric. Use the highest oxidation state of each element. 7.102 State whether each of the following elements is a gas, liquid, or solid under atmospheric conditions. Also state whether it exists in the elemental form as atoms, molecules, or a three-dimensional network: Mg , CI, Si, Kr, 0 , J, Hg, Br. 7.103 What factors account for the unique nature of hydrogen? 7.104 The air in a manned spacecraft or submarine needs to be purified of exhaled carbon dioxide. Write equations for the reactions between carbon dioxide and (a) lithium oxide (Li 2 0), (b) sodium peroxide ( Na 2 0 z), and (c) potass ium superoxide (K0 2 ). 7.105 The formula for calculating the energies of an electron in a hydrogen-like ion is given in Problem 7.46. This equation can be applied only to one-electron atoms or ions. One way to modify it for more complex species is to replace Z with Z - CT or Zeff' Calculate the value of CT if the first ionization energy of helium is 3.94 x 10- 18 J per atom. (Disregard the minus sign in the given equa tion in your calculation.) 7.106 Why do noble gases have negative electron affinity values? 7.107 The atomic radius of K is 227 pm and that of K+ is 138 pm. Calculate the percent decrease in volume that occurs when K(g) is conve11ed to K+(g). (The volume of a sphere is :h rr 3, where r is the radius of the sphere.) 7.108 The atomic radius of F is 72 pm and that of F- is 133 pm. Calculate the percent increase in volume that occurs when F(g) is converted to F- (g). (See Problem 7.107 for the volume of a sphere.) 7.109 A technique called photoelectron spectroscopy is used to measure the ionization energy of atoms. A gaseous sample is irradiated. with UV light, and electrons are ejected from the valence shel!. The kinetic energies of the ejected electrons are measured. Because the energy of the UV photon and the kinetic energy of the ejected electron are known, we can write 7.11 0 hv = IE + ~mu 2 where v is the frequency of the UV light, and m and u are the mass and velocity of the electron, respectively. In one experiment the kinetic energy of the ejected electron from potassium is found to be 5.34 x 10- 19 J using a UV source of wavelength 162 nm. Calculate the ionization energy of potassium. How can you be sure that this ionization energy corresponds to the electron in the valence shell (that is, the most loosely held electron)? The energy needed for the following process is 1.96 X 10 4 kJ/mol: If the first ionization energy of lithium is 520 kJ/mol, calculate the second ionization energy of lithium, that is, the energy required for the process Li+(g) • Li 2+ (g) + e- (Hint: You need the equation in Problem 7.46.) 7.111 A student is given samples of three elements, X, Y, and Z, which could be an alkali metal, a member of Group 4A, or a member of Group SA. She makes the following observations: Element X has a metallic luster and conducts electricity. It reacts slowly with hydrochloric acid to produce hydrogen gas. Element Y is a light yellow solid that does not conduct electricity. Element Z has a metallic luster and conducts electricity. When exposed to air, it slowly forms a white powder. A solution of the white powder in water is ba sic. What can you conclude about the elements from these observations? 7.112 What is the electron affinity of the Na+ ion? 7.113 The ionization energies of sodium (i n kJ/mol), starting with the first and ending with the eleventh, are 496, 4562, 6910, 9543, 13,354,16 ,613,2 0,117 ,2 5, 496,28,932,141,362,159,075. Plot the log of ionization energy (y axis) versus the number of ionization (x axis); for example, log 496 is plotted versus 1 (labeled 1Eb the first ionization energy), log 4562 is plotted versus 2 (labeled lE 2 , the second ionization energy), and so on. (a) LabellE] through IE] ] with the electrons in orbitals such as Is, 2s, 2p, and 3s. (b) What can you deduce about electron shells from the breaks in the curve? 7.114 Experimentally, the electron affinity of an element can be determined by using a laser light to ionize the anion of the element in the gas phase: X-(g) + hv -_. X( g) + e- Referring to Figure 7.10, calculate the photon wavelength (in nm) corresponding to the electron affinity for chlorine. In what region of the electromagnetic spectrum does this wavelength fall? 7.115 Explain, in terms of their electron configurations, why Fe 2+ is more easily oxidized to Fe 3+ than Mn2 + to Mn 3+ . 7.116 Write the formulas and names of the hydrides of the following second-period elements: Li, C, N, 0, F. Predict their reactions with water. 7.117 Ba sed on knowledge of the electronic configuration of titanium, state which of the following compounds of titanium is unlikely to exist: K3 TiF6' K2 Ti 2 0 s , TiCI 3 , K2 Ti0 4 , K2 TiF 6 · 7.118 In halogen displacement reactions a halogen element can be generated by oxidizing its anions with a halogen element that lies above it in the periodic table. This means that there is no way to prepare elemental fluorine, because it is the first member of Group 7 A. Indeed, for years the only way to prepare elemental fluorine was to oxidize F- ions by electrolytic mean s. Then, in 1986, a chemist reported that by combining potass ium hexafluoromanganate(IV) (K2 MnF 6) with antimony pentafluoride (S bF s ) at 150°C, he had generated elemental fluorine. Balance the following equation representing the reaction: QUESTIONS AND PROBLEMS 273 7.119 Write a balanced equation for the preparation of (a) molecular oxygen, (b) ammonia, (c) carbon dioxide, (d) molecular hydrogen, (e) calcium oxide. Indicate the physical state of the reactants and products in each equation. 7.120 Write chemical formulas for oxides of nitrogen with the following oxidation numbers: + 1, +2, +3, +4, +5. (Hint: There are two oxides of nitrogen with a +4 oxidation number.) 7.121 Most transition metal ions are colored. For example, a solution of CUS04 is blue. How would you show that the blue color is due to the hydrated Cu 2 + ions and n ot the SO ~ - ions? 7.122 In general, atomic radius and ionization energy have opposite periodic trends. Why? 7.123 Explain why the electron affinity of nitrogen is approximately zero, while the elements on either side, carbon and oxygen, have substantial positi ve electron affinities. 7.124 Consider the halogens c hl orine, bromine, and iodine. The melting point and boiling point of chlorine are -101.0 °C and -3 4.6°C and those of iodine are 113.5°C and 184.4°C, respectively. Thus chlorine is a gas and iodine is a solid under room conditions. Estimate the melting point and boiling point of bromine. Compare your values with those from the webelements.com website. 7.125 Although it is possible to determine the second, third, and higher ionization energies of an element, the sa me cannot usually be done with the electron affinities of an element. Explain. 7.126 Little is known of the chemistry of astatine, the last member of Group 7 A. Describe the physical characteristics that you would expect this halogen to have. Predict the products of the reaction between sodium astatide ( NaAt ) and sulfuric acid. (Hint: Sulfuric acid is an oxidizing agent.) 7.127 As discussed in the chapter, the atomic mass of argon is greater than that of potassium. This observation created a problem in the early development of the periodic table because it meant that argon should be placed after potassium. (a) How was this difficulty resolved? (b) From the following data, calculate the average atomic ma sses of argon and potassium: Ar-36 (35.9675 amu, 0.337 percent ), Ar-38 (37.9627 amu, 0.063 percent), Ar- 40 (39.9624 amu, 99.60 percent), K-39 (38.9637 amu, 93.258 percent), K-40 (3 9.9640 amu, 0.0117 percent), K-41 (40.9618 amu, 6.730 percent ). 7.128 Calculate the maximum wavelength of light (in nm) required to ionize a single sodium atom. 7.129 Predict the atomic number and ground-state electron configuration of the next member of the alkali metals after francium. 7.130 Why do elements that have high ionization energies also have more positive electron affinities? Which group of elements would be an exception to this generalization? . 7.131 The first four ionization energies of an element are approximately 738,1450,7.7 X 10 3 , and 1.1 X 10 4 kJ/mol. To which periodic group does this element belong? Explain your answer. 7.132 Some chemists think that helium should properly be called "he lon." Why? What does the ending in helium (-ium) suggest? 274 CHAPTER 7 Electron Configuration and the Periodic Table 7.133 7.134 7.135 (a) The formula of the simplest hydrocarbon is CH 4 ( methane ). Predict the formulas of the s imple st compound s formed bet w een hydrogen and the following element s: silicon, germanium , tin, and lead. (b) Sodium hydride ( NaH ) is an ionic compound. Would you expect rubidium hydride (RbH) to be more or less ionic than NaB ? (c) Predict the reaction between · radium (Ra ) and water. (d) When expos ed to air, aluminum forms a tenacious oxide (AI 2 0 3 ) coating that protects the metal from corrosion. Which metal in Group 2A would you expect to exhibit similar properties? (See the margin note on p. 257.) Match each of the elements on the right with its description on the left: (a) A pale yellow gas that reacts with water. (b) A soft metal that reacts with water to produce hydrogen. (c) A metalloid that is hard and has a high melting point. (d) A colorless, odorless gas. (e) A metal that is more reactive than iron, but doe s not corrode in air. Nitrogen (N 2 ) Boron (B ) Fluorine (F 2 ) Aluminum ( AI ) Sodium (Na) Write at least two paragraphs de scribing the importance of the periodic table. Pay particular attention to the significance of the position of an element in the table and how the position relates to the chemical and physical properties of the element. 7.136 On one graph, plot the effective nuclear charge (shown in parentheses) and atomic radius (see Figure 7.6) versus atomic number for the second-period elements: Li(1.28), Be(1.91), B(2.42), C(3.14), N(3.83), 0(4.45) , F(5.1O), Ne(5.76). Comment on t he trends. 7.137 One allotropic form of an element X is a colorless crystalline solid. The reaction of X with an exce ss amount of oxygen produces a colorless gas. This gas dissolves in water to yield an acidic solution. Choose one of the following elements that matches X: (a) sulfur, (b) phosphorus, (c) carbon, (d) boron, (e) silicon. 7.138 The ionization energy of a certain element is 412 kJ/mo!. When the atoms of this element are in the first excited state, howe ver, the ionization energy is only 126 kJ/mo!. Based on this information, calculate the wavelength of light emitted in a transition from the first excited state to the ground state. 7.139 One way to estimate the effective charge (Ze ff) of a many-electron atom is to use the equation lE I = (1312 kJ/mol)(Z~ffln2), where lE I is the first ionization energy and n is the principal quantum number of the shell in which the electron resides. Use this equation to calculate the effective charges of Li, Na, and K. Also calculate Ze ffin for each meta!. Comment on your results. PRE-PROFESSIONAL PRACTICE EXAM PROBLEMS: PHYSICAL AND BIOLOGICAL SCIENCES These questions are not based on a de scriptive passage. 1. A halogen ha s valence electrons in which orbitals? a) s b) s andp c)p d) s, p, and d 2. How many subshells does a shell with principal quantum number n contain? a) n b) n 2 c) n - 1 d) 2n - 1 • 3. In a shell that contains anjsubshell, what is the ratio ofjorbitals to s orbitals? a) 14:1 b)7:1 c) 7:3 d) 7:5 4. What is the maximum number of electrons that can be in the n = 3 shell? a) 2 b) 6 c) 8 d) 18 ANSWERS TO IN-CHAPTER MATERIALS 275 ANSWERS TO IN-CHAPTER MATERIALS Answers to Practice Problems 7.1A Ge. 7.1B Bi, and less so Sb and As. 7.2A (a) Is2 2i 2 p 6 3i 3 p 3, p-block, (b) Ii 2S 2 2l3s 2 3l4s 2 , s-block, (c) Is 2 2S 2 2l3 i 3p6 4i 3dlO 4i, p-block. 7.2B (a) AI, (b) Zn, (c) Sr. 7.3A P < Se < Ge. 7.3B P and Se. 7.4A Mg, Mg. 7.4B Rb has a smaller Ze ff , IE2 for Rb corresponds to the removal of a core electron. 7.SA AI. 7.SB Adding an electron to As involves pairing. 7.6A (a) [Ne], (b) [Ar], (c) [Kr]. 7.6B N 3 - , 0 2 - , P- , Ne, Na+, Mg 2+ , AI H . 7.7A (a) [Ar]3d 6 , (b) [Ar]3d 9 , (c) [Kr]3d lO. 7.7B Cu+. 7.8A Rb+ < Kr < Br - < Se 2 - . 7.8B P- , 0 2 - , N 3 - , Na+, Mg 2 +, Mg 2 +, AIH . Answers to Checkpoints 7.1.1 c. 7.1.2 a. 7.2.1 b. 7.2.2 a, d, e. 7.4.1 b. 7.4.2 c. 7.4.3 e. 7.4.4 a. 7.5.1 b, c, e. 7.5.2 b, d. 7.5.3 d. 7.5.4 b. 7.6.1 d, e. 7.6.2 a, c. Answers to Applying What You've Learned a) Ii 2S I . b) In order of increasing atomic radius: Li < Na < K < Rb < Cs. c) In order of increasing ionization energy ( IE): Cs < Rb < K < Na < Li. d) Li+: Ii or [He] ; Na +: I s2 2i 2 p 6 or [Ne]; K+: li 2 s2 2p 6 3i 3 p 6 or [Ar]; Rb+: li 2i 2p 6 3i 3 p 64 s2 3d lO 4l or [Kr]; Cs+: Is 22s2 2l3 i3p 6 4i 3d lO 4l Ss 2 4d lO Sp 6 or [Xe]. e) Isoelectronic with Li+: He. Isoelectronic with Na +: Mg 2+, AI H , Ne , P- , 0 2 - , and N 3 Isoelectronic with K+: Ca 2+ , Ar, Cl -, S2- , and p 3 Isoelectronic with Rb +: Sr 2+, Kr, Br- , and Se 2 - . Isoelectronic with Cs+: Ba 2+, Xe, r, and Te 2 - . [...]... =QX r Equation 8.1 The distance, r, between partial charges in a polar diatomic molecule is the bond length expressed in meters Bond lengths are usually given in angstroms (A), or picometers (pm), so it is generally necessary to convert to meters In order for a diatomic molecule containing a polar bond to be electrically neutral, the partial positive and partial negative charges must have the same... Calculated partial charges should always be less than 1 If a "partial" charge were 1 or greater, it would indicate that at least one electron had been transferred from one atom to the other Remember that polar bonds involve unequal sharing of electrons, not a complete transfer of electrons 1 9.2XIO- 1Im In units of electronic charge: 6.598 X 10- 20 C X l e9 = 0.41e 1.6022 X 10- 1 C Therefore, the partial... + 0.41 H _ :-0.41 F •• Practice Problem A Using data from Table 8.5, determine the magnitude of the partial charges in HBr Practice Problem B Given that the partial charges on C and 0 in carbon monoxide are +0.020 and -0.020, respectively, calculate the dipole moment of CO (The distance between the partial charges, I; is 113 pm.) • Checkpoint 8.4 8.4.1 Electronegativity and Polarity Which of the following... greater injury than more concentrated ones by penetrating more deeply before causing injury, thus delaying the onset of symptoms and preventing timely treatment Determine the magnitude of the partial positive and partial negative charges in the HF molecule Strategy Rearrange Equation 8.1 to solve for Q Convert the resulting charge in coulombs to charge in units of electronic charge o Setup According... most of the time, so they do not participate in bond formation Thus, each F atom has seven valence electrons (the two 2s and five 2p electrons) According to Figure 8.1, there is only one unpaired electron on F, so the formation of the F2 molecule can be represented as follows: •• •• •• •• •• •• :F· + ·F: - _ :F:F: or :F-F: •• •• •• •• • • •• Only two valence electrons participate in the bond that forms... William Debye (1884-1 966) American chemist and physicist of Dutch origin Debye made many significant contributions to the study of molecular structure, polymer chemistry, X-ray analysis, and electmlyte solutions He was awarded the Nobel Prize in Chemistry in 1936 We usually express the charge on an electron as -1 This refers to units of electronic charge However, remember that the charge on an electron... hand, where the two bonding atoms are different, the electrons are not shared equally They spend more time in the vicinity of the F atom than in the vicinity of the H atom (The 0 symbol is used to denote partial charges on the atoms.) In NaP, the electrons are not shared at all but rather are transferred from sodium to fluorine One way to visualize the distribution of electrons in species such... the H-O bonds in H 2 0 2 , and (c) the 0-0 bond in H 20 2 Practice Problem B Using data from Figure 8.6, list all the main group elements that can form purely ionic compounds with N Dipole Moment and Partial Charges The shift of electron density in a polar bond is symbolized by placing a crossed arrow (a dipole arrow) above the Lewis structure to indicate the direction of the shift Por example, • H-P:... chemist Alfred Nobel died in 1896, his will specified that the bulk of his considerable fortune was to be used to establish the prizes that bear his name The prizes, given annually in five categories (Chemistry, Physics, Physiology or Medicine, Literature, and Peace), are intended to recognize significant contributions to the betterment of humankind In life, Nobel had been a prolific scientist and entrepreneur... containing a polar bond to be electrically neutral, the partial positive and partial negative charges must have the same magnitude Therefore, the Q term in Equation 8.1 refers to the magnitude of the partial charges and the calculated value of jJ- is always positive 3 Dipole moments are usually expressed in debye units (D), named for Peter Debye In terms of more familiar SI units, 1 D = 3.336 X 10- . of the alkali metal cations. [ ~~ Sample Prob l em 7.6] e) For each alkali metal cation in part (d), identify an isoelectronic series consisting ;;;- • , of a noble gas and, where appropriate,. • According to Coulomb's law, the attractive force (F) between two oppositely charged particles (QI and Q 2) is directly proportional to the product of the charges and inversely. Calculate Z eff for these electrons, and explain the differences from the values you determined in part (a). 270 CHAPTER 7 Electron Configuration and the Periodic Table 7.25 The electron