Given the following partial (valence-level) electron

Problem 70P The hot glowing gases around the Sun, the corona, can reach millions of degrees Celsius, high enough to remove many electrons from gaseous atoms. Iron ions with charges as high as 14+ have been observed in the corona. Which ions from Fe+ to Fe14+ are paramagnetic? Which would be most strongly attracted to a magnetic field?

Partial (valence-level) electron configurations for four different ions are shown below: Identify the elements from which the ions are derived, and write the formula of the oxide each ion forms.

Problem 73P Data from the planet Zog for some main-group elements are shown below (Zoggian units are linearly related to Earth units but are not shown). Radio signals from Zog tell that balloonium is a monatomic gas with two positive nuclear charges. Use the data to deduce the names that Earthlings give to these elements: Name Atomic Radius IE1 EA1 Balloonium 10 339 0 Inertium 24 297 +4.1 Allotropium 34 143 ?28.6 Brinium 63 70.9 ?7.6 Canium 47 101 ?15.3 Fertilium 25 200 0 Liquidi uin 38 163 ?46.4 Utilium 48 82.4 ?6.1 Crimsonium 72 78.4 ?2.9

Problem 1P What would be your reaction to a claim that a new element had been discovered, and it fit between tin (Sn) and antimony (Sb) in the periodic table?

Problem 4P Summarize the rules for the allowable values of the four quantum numbers of an electron in an atom.

Problem 6P State the exclusion principle. What does it imply about the number and spin of electrons in an atomic orbital?

Problem 7P What is the key distinction between sublevel energies in one- electron species, such as the H atom, and those in many-electron species, such as the C atom? What factors lead to this distinction? Would you expect the pattern of sublevel energies in Be3+ to be more like that in H or that in C? Explain.

Problem 5P Which of the quantum numbers relate(s) to the electron only? Which relate(s) to the orbital?

Problem 9P What is penetration? How is it related to shielding? Use the penetration effect to explain the difference in relative orbital energies of a 3p and a 3d electron in the same atom.

Problem 8P Define shielding and effective nuclear charge. What is the connection between the two?

Problem 11P How many electrons in an atom can have each of the following quantum number or sublevel designations? (a) n = 2, l= 1, ml = 0 ________________ (b) 5p ________________ (c) n = 4,l = 3

Problem 10P How many electrons in an atom can have each of the following quantum number or sublevel designations? (a) n = 2, l = 1 (b) 3d (c) 4s

How many electrons in an atom can have each of the following quantum number or sublevel designations? (a) \(4p\) (b) \(n = 3, l = 1, ml = + 1\) (c) \(n = 5, l = 3\) Equation Transcription: Text Transcription: 4p n = 3, l = 1, ml = + 1 n = 5, l = 3

Problem 14P State the periodic law, and explain its relation to electron configuration. (Use Na and K in your explanation.)

Problem 15P State Hund’s rule in your own words, and show its application in the orbital diagram of the nitrogen atom.

Problem 13P How many electrons in an atom can have each of the following quantum number or sublevel designations? (a) 2s ________________ (b) n = 3, l = 2 ________________ (c) 6d

Problem 17P For main-group elements, are outer electron configurations similar or different within a group? Within a period? Explain.

Problem 18P Write a full set of quantum numbers for the following: (a) The outermost electron in an Rb atom ________________ (b) The electron gained when an S– ion becomes an S2– ion ________________ (c) The electron lost when an Ag atom ionizes ________________ (d) The electron gained when an F– ion forms from an F atom.

Problem 16P How does the aufbau principle, in connection with the periodic law, lead to the format of the periodic table?

Problem 19P Write a full set of quantum numbers for the following: (a) The outermost electron in an Li atom ________________ (b) The electron gained when a Br atom becomes a Br– ion ________________ (c) The electron lost when a Cs atom ionizes ________________ (d) The highest energy electron in the ground-state B atom

Write the full ground-state electron configuration for each: (a) Rb (b) Ge (c) Ar

Problem 22P Draw an orbital diagram showing valence electrons, and write the condensed ground-state electron configuration for each: (a) Ti ________________ (b) CI ________________ (c) V

Problem 21P Write the full ground-state electron configuration for each: (a) Br ________________ (b) Mg ________________ (c) Se

From each partial (valence-level) orbital diagram, write the ground-state electron configuration and group number:

Problem 28P How many inner, outer, and valence electrons are present in an atom of each of the following elements? (a) O ________________ (b) Sn ________________ (c) Ca ________________ (d) Fe ________________ (e) Se

Problem 25P Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element: (a) [Kr] 5s24d10 ________________ (b) [Ar] 4s23d8

Problem 23P Draw an orbital diagram showing valence electrons, and write the condensed ground-state electron configuration for each: (a) Ba ________________ (b) Co ________________ (c) Ag

Problem 27P What experimental support did de Broglie’s concept receive?

Problem 24P Draw the partial (valence-level) orbital diagram, and write the symbol, group number, and period number of the element: (a) [He] 2s2 2p4 ________________ (b) [Ne] 3s23p3

Problem 29P How many inner, outer, and valence electrons are present in an atom of each of the following elements? (a) Br ________________ (b) Cs ________________ (c) Cr ________________ (d) Sr ________________ (e) F

Problem 31P Identify each element below, and give the symbols of the other elements in its group: (a) [Ar] 4s23d104p4 ________________ (b) [Xe] 6s24f145d2 ________________ (c) [Ar] 4s23d5

Problem 37P In a plot of IE1 for the Period 3 elements (see Figure 8.17, p. 321), why do the values for elements in Groups 3A(13) and 6A(16) drop slightly below the generally increasing trend?

Problem 30P Identify each element below, and give the symbols of the other elements in its group: (a) [He] 2s2 2p1 ________________ (b) [Ne] 3s23p4 ________________ (c) [Xe] 6s25d1

Given the following partial (valence-level) electron configurations, (a) identify each element, (b) rank the four elements in order of increasing atomic size, and (c) rank them in order of increasing ionization energy:

Problem 34P In what region of the periodic table will you find elements with relatively high EEs? With relatively low IEs?

Problem 38P Which group in the periodic table has elements with high (endothermic) IE1 and very negative (exothermic) first electron affinities (EA1)? Give the charge on the ions these atoms form.

Problem 32P One reason spectroscopists study excited states is to gain information about the energies of orbitals that are unoccupied in an atom’s ground state. Each of the following electron configurations represents an atom in an excited state. Identify the element, and write its condensed ground-state configuration: (a) 1s22s22p63s13p1 ________________ (b) 1s22s22p63s23p44s1 ________________ (c) 1s22s22p63s23p64s23d44p1 ________________ (d) 1s22s22p53s1

Problem 39P How does d- electron shielding influence atomic size among the Period 4 transition elements?

Problem 42P Arrange each set of atoms in order of increasing IE1: (a) Sr,Ca,Ba ________________ (b) N,B,Ne ________________ (c) Br,Rb,Se ________________ (d) As,Sb,Sn

Problem 41P Arrange each set in order of decreasing atomic size: (a) Ge, Pb, Sn ________________ (b)Sn,Te,Sr ________________ (c) F,Ne,Na ________________ (d) Be, Mg, Na

Problem 40P Arrange each set in order of increasing atomic size: (a) Rb, K, Cs ________________ (b) C, O, Be ________________ (c) Cl,K,S ________________ (d) Mg,K,Ca

Problem 43P Arrange each set of atoms in order of decreasing IE1: (a) Na, Li, K ________________ (b) Be,F,C ________________ (c) Cl,Ar,Na ________________ (d) Cl,Br,Se

Problem 45P Write the full electron configuration of the Period 3 element with the following successive IEs (in kJ/mol): IE, = 738 ________________ IE2 = 1450 ________________ IE3 = 7732 ________________ IE4 = 10,539 ________________ IE5 = 13,628

Problem 44P Write the full electron configuration of the Period 2 element with the following successive IEs (in kJ/mol): IE1 = 801 ________________ IE2 = 2427 ________________ IE3 = 3659 ________________ IE4 = 25,022 ________________ IE5 = 32,822

Problem 47P Which element in each of the following sets would you expect to have the lowest IE3? (a) Na, Mg, Al ________________ (b) K,Ca,Sc ________________ (c) Li,Al,B

Problem 49P Summarize the trend in metallic character as a function of position in the periodic table. Is it the same as the trend in atomic size? Ionization energy?

Problem 48P List three ways in which metals and nonmetals differ.

Problem 46P Which element in each of the following sets would you expect to have the highest IE2? (a) Na, Mg, A1 ________________ (b) Na, K, Fe ________________ (c) Sc, Be, Mg

Problem 50P Summarize the acid-base behavior of the main-group metal and nonmetal oxides in water. How does oxide acidity in water change down a group and across a period?

What is a pseudo-noble gas configuration? Give an example of one ion from Group 3A(13) that has it.

Problem 52P The charges of a set of isoelectronic ions vary from 3+ to 3–. Place the ions in order of increasing size.

Problem 55P Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each: (a) Cl ________________ (b) Na ________________ (c) Ca

Problem 56P Write the charge and full ground-state electron configuration of the monatomic ion most likely to be formed by each: (a) Rb ________________ (b) N ________________ (c) Br

Problem 53P Which element would you expect to be more metallic? (a) Ca or Rb ________________ (b) Mg or Ra ________________ (c) Br or I

Problem 59P Write the condensed ground-state electron configurations of these transition metal ions, and state which are paramagnetic: (a) V3+ ________________ (b) Cd2+ ________________ (c) Co3+ ________________ (d) Ag+

Problem 58P How many unpaired electrons are present in the ground state of an atom from each of the following groups? (a) 4A(14) ________________ (b) 7A(17) ________________ (c) 1A(1) ________________ (d) 6A(16)

Problem 57P How many unpaired electrons are present in the ground stale of an atom from each of the following groups? (a) 2A(2) ________________ (b) 5A(15) ________________ (c) 8A(18) ________________ (d) 3A(13)

Problem 60P Write the condensed ground-state electron configurations of these transition metal ions, and state which are paramagnetic: (a) Mo3+ ________________ (b) Au+ ________________ (c) Mn2+ ________________ (d) Hf2+

Problem 61P Palladium (Pd; Z = 46) is diamagnetic. Draw partial orbital diagrams to show which of the following electron configurations is consistent with this fact: (a) [Kr] 5s24d8 ________________ (b) [Kr] 4d10 ________________ (c) [Kr]5s14d9

Rubidium and bromine atoms are depicted at right. (a) What monatomic ions do they form? (b) What electronic feature characterizes this pair of ions, and which noble gas are they related to? (c) Which pair best represents the relative ionic sizes?

Problem 62P Niobium (Nb: Z = 41) has an anomalous ground-state electron configuration for a Group 5B(5) element: [Kr] 5s14d4. What is the expected electron configuration for elements in this group? Draw partial orbital diagrams 10 show how paramagnetic measurements could support niobium’s actual configuration.

Problem 63P Rank the ions in each set in order of increasing size, and explain your ranking: (a) Li+, K+, Na+ ________________ (b) Se2–, Rb+, Br– ________________ (c) O2–. F–, N3–

Problem 67P When a nonmetal oxide reacts with water, it forms an oxoacid with the same nonmetal oxidation state. Give the name and formula of the oxide used to prepare each of these oxoacids:(a) hypochlorous acid; (b) chlorous acid; (c) chloric acid; (d) perchloric acid; (e) sulfuric acid; (f) sulfurous acid; (g) nitric acid; (h) nitrous acid; (i) carbonic acid; (j) phosphoric acid.

Problem 65P Name the element described in each of the following: (a) Smallest atomic radius in Group 6A(16) ________________ (b) Largest atomic radius in Period 6 ________________ (c) Smallest metal in Period 3 ________________ (d) Highest IE1, in Group 4A(14] ________________ (e) Lowest IE1, in Period 5 ________________ (f) Most metallic in Group 5A(15) ________________ (g) Group 3A(13) element that forms the most basic oxide ________________ (h) Period 4 element with filled outer level ________________ (i) Condensed ground-state electron configuration of [Ne] 3s23p2 ________________ (j) Condensed ground-state electron configuration of [Kr] 5s24d6 ________________ (k) Forms 2+ ion with electron configuration [Ar] 3d3 ________________ (l) Period 5 element that forms 3+ ion with pseudo-noble gas configuration ________________ (m) Period 4 transition element that forms 3+ diamagnetic ion ________________ (n) Period 4 transition element that forms 2+ ion with a hall-filled d sublevel ________________ (o) Heaviest lanthanide ________________ (p) Period 3 element whose 2– ion is isoelectronic with Ar ________________ (q) Alkaline earth metal whose cation is isoelectronic with Kr ________________ (r) Group 5A(15) metalloid with the most acidic oxide

Problem 64P Rank the ions in each set in order of decreasing size, and explain your ranking: (a) Se2–, S2?, O2– ________________ (b) Te2–, Cs+, I– ________________ (c) Sr2+, Ba2+, Cs+

Problem 68P The energy difference between the 5d and 6s sublevels in gold accounts for its color. Assuming this energy difference is about 2.7 eV (electron volt; 1 eV = 1.602×10–19 J), explain why gold has a warm yellow color.

Problem 69P Write the formula and name of the compound formed from the following ionic interactions;(a) The 2+ ion and the 1– ion are both isoelectronic with the atoms of a chemically unreactive Period 4 element, (b) The 2+ ion and the 2? ion are both isoelectronic with the Period 3 noble gas. (c) The 2+ ion is the smallest with a filled d subshell; the anion forms from the smallest halogen, (d) The ions form from the largest and smallest ionizable atoms in Period 2.