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Atomic Electron Configurations and Chemical Periodicity

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1 Atomic Electron Configurations and Chemical Periodicity
Chapter 8 Atomic Electron Configurations and Chemical Periodicity

2 Chapter goals Understanding the role magnetism plays in determining and revealing atomic structure. Understand effective nuclear charge and its role in determining atomic properties. Write the electron configuration of neutral atoms and monatomic ions. Understand the fundamental physical properties of the elements and their periodic trends.

3 Electron Spin and the Fourth Quantum Number
The fourth quantum number is the spin quantum number which has the symbol ms. The spin quantum number only has two possible values. ms = +1/2 or −1/2 ms = ± 1/2 This quantum number tells us the spin and orientation of the magnetic field of the electrons. Wolfgang Pauli discovered the Exclusion Principle in 1925. No two electrons in an atom can have the same set of 4 quantum numbers, n, l, ml, and ms

4 Electron Spin Spin quantum number effects:
Every orbital can hold up to two electrons. Consequence of the Pauli Exclusion Principle. The two electrons are designated as having one spin up  ms = +1/2 and one spin down ms = −1/2 Spin describes the direction of the electron’s magnetic field.

5 Paramagnetism and Diamagnetism
Unpaired electrons have their spins aligned   or   (in diff. orbitals) This increases the magnetic field of the atom. Total spin  0, because they add up. Atoms with unpaired electrons are called paramagnetic . Paramagnetic atoms are attracted to a magnet.

6 Paramagnetism and Diamagnetism
Paired electrons have their spins unaligned . (in the same orbital) Paired electrons have no net magnetic field. Total spin = 0, because of cancellation, ½ − ½ = 0 Atoms with no unpaired electrons are called diamagnetic. Diamagnetic atoms are not attracted to a magnet.

7 Atomic Orbitals, Spin, and # of Electrons
Because two electrons in the same orbital must be paired (due to Pauli’s Exclusion Principle), it is possible to calculate the number of orbitals and the number of electrons in each n shell. The number of orbitals per n level is given by n2 (see table at end of chapter 7.) The maximum number of electrons per n level is 2n2 (two electrons per orbital.) The value is 2n2 because of the two paired electrons per orbital.

8 #orbitals Max n shell l subshell ml #e– 1 K s 1 1 2 2 L s 1 2 8 4 1 p
s 1 1 2 2 L s 1 2 8 4 1 p –1,0,1 3 6 3 M s 1 2 1 p 18 –1,0,1 3 9 6 2 d -2,-1,0,1,2 5 10 4 N s 1 2 1 p –1,0,1 3 6 32 16 2 d 10 -2,-1,0,1,2 5 3 f -3,-2,-1,0,1,2,3 7 14

9 Atomic Subshell Energies and Electron Assignments
The principle that describes how the periodic chart is a function of electronic configurations is the Aufbau Principle. The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available.

10 Penetrating and Shielding
the radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p the weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force, they are more shielded from the attractive force of the nucleus the deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively the result is that the electrons in the 2s sublevel are lower in (more negative) energy than the electrons in the 2p

11 Atomic Subshell Energies and Electron Assignments
The Aufbau Principle describes the electron filling order in atoms. This is product of the effective nuclear charge, Z*, Zeff For the same n, Z* is higher for s orbital: s > p > d > f Then, e− in s is the most attracted by nucleus and has the lowest energy

12 Atomic Subshell Energies and Electron Assignments
One mnemonic to remember the correct filling order for electrons in atoms is the increasing (n + ) value

13 Atomic Subshell Energies and Electron Assignments
or we can use this periodic chart

14 Atomic Electron Configurations
Now we will use the Aufbau Principle to determine the electronic configurations of the elements on the periodic chart. 1st row elements

15 Atomic Electron Configurations
Hund’s rule tells us that the electrons will fill the p and d orbitals by placing electrons in each orbital singly and with same spin until half-filled. That is the rule of maximum spin. Then the electrons will pair to finish the p orbitals. Electrons in orbitals of or same kind, such as p or d orbitals, in the same shell (n), have the same energy; the are said to be degenerate.

16 Atomic Electron Configurations
3rd row elements…

17 Atomic Electron Configurations
4th row elements…

18 Atomic Electron Configurations
4th row elements…

19 Atomic Electron Configurations
4th row elements… The five d orbitals are degenerate

20 Atomic Electron Configurations
4th row elements…

21 Atomic Electron Configurations
4th row elements… The five d orbitals are degenerate

22 Atomic Electron Configurations
4th row elements…

23 Atomic Electron Configurations
4th row elements… The [Ar] 4s1 3d5 configuration of Cr is more stable than [Ar] 4s2 3d4 (expected)

24 Atomic Electron Configurations
4th row elements… The [Ar] 4s1 3d10 full d configuration of Cu is more stable than [Ar] 4s2 3d9 (expected)

25 Atomic Electron Configurations
4th row elements…

26 Atomic Electron Configurations
4th row elements… (remember Hund’s rule):   __ is better (lower energy) than  __ __ 4p p

27 Atomic Electron Configurations
Lanthanides (4f) 56Ba [Xe] 6s2 57La [Xe] 5d1 6s2 58Ce [Xe] 4f1 5d1 6s2 59Pr [Xe] 4f3 6s Praseodymium 70Yb [Xe] 4f14 6s Ytterbium 71Lu [Xe] 4f14 5d1 6s Lutetium

28 Periodic Table

29 s, p, d, and f-block in the Periodic Table

30 P 1 2 (P–1)d 3 4 5 6 7 (P)p (P)s (P–2)f Ba 56 Be 4 Mg 12 Ca 20 Sr 38
Ra 88 2A 1A H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Ga 31 In 49 Tl 81 B 5 Al 13 3A Ge 32 Sn 50 Pb 82 C 6 Si 14 4A As 33 Sb 51 Bi 83 N 7 P 15 5A Se 34 Te 52 Po 84 O 8 S 16 6A Br 35 I 53 At 85 F 9 Cl 17 7A Kr 36 Xe 54 Rn 86 Ne 10 Ar 18 He 2 8A 1 2 (P–1)d 3 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 21 22 23 24 25 26 27 28 29 30 4 Sc Ti V Cr Mn Fe Co Ni Cu Zn 39 40 41 42 43 44 45 46 47 48 5 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd 57 72 73 74 75 76 77 78 79 80 6 La Hf Ta W Re Os Ir Pt Au Hg 89 104 105 106 107 108 109 7 (P)p Ac Rf Db Sg Bh Hs Mt (P)s 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 90 91 92 93 94 95 96 97 98 99 100 101 102 103 (P–2)f Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

31 Valence Electrons They determine the chemical properties of an
electrons in shell with highest n, i.e., the outermost electrons, those beyond the core electrons 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2 3p2 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 1s2 2s2 1s2 2s2 2p6 3s2 3p6 4s2 3d7 They determine the chemical properties of an element. For the representative elements, they are the ns and np electrons; for transition elements they are the ns and (n−1)d electrons.

32 # of valence electrons = 1 4s1
P 1A 1 1s1 1 H 3 2 2s1 Li 11 3 3s1 Na # of valence electrons = 1 19 4 4s1 K 37 5 5s1 Rb 55 6 6s1 Cs 87 7 7s1 Fr

33 # of valence electrons = 2 4s2
Be 12 3s2 Mg # of valence electrons = 2 20 4s2 Ca 38 5s2 Sr 56 6s2 Ba 88 7s2 Ra

34 # of valence electrons = 3 4s2 4p1
5 2s2 2p1 B 13 3s2 3p1 Al # of valence electrons = 3 31 4s2 4p1 Ga 49 5s2 5p1 In 81 6s2 6p1 Tl

35 # of valence electrons = 7
9 2s2 2p5 F 17 3s2 3p5 Cl 35 4s2 4p5 Br 53 5s2 5p5 I 85 6s2 6p5 At For the representative elements, the # of valence electrons = # of group

36 The element X has the valence shell electron configuration, ns2 np4.
X belongs to what group? chalcogens Ba 56 Be 4 Mg 12 Ca 20 Sr 38 Ra 88 2A 1A H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Ga 31 In 49 Tl 81 B 5 Al 13 3A Ge 32 Sn 50 Pb 82 C 6 Si 14 4A As 33 Sb 51 Bi 83 N 7 P 15 5A Se 34 Te 52 Po 84 O 8 S 16 6A Br 35 I 53 At 85 F 9 Cl 17 7A Kr 36 Xe 54 Rn 86 Ne 10 Ar 18 He 2 8A Hg La 57 Sc 21 Y 39 89 3B Ti 22 Zr 40 Hf 72 104 4B V 23 Nb 41 Ta 73 105 5B Cr 24 Mo 42 W 74 106 6B Mn Tc Re 25 43 75 107 7B Fe Os 26 Ru 44 76 108 8B Ir 45 77 Co 27 Rh 109 Ni 28 Pd 46 Pt 78 Cu 29 Ag 47 Au 79 1B 30 Cd 48 80 2B Zn Ac Unq Unp Unh Uns Uno Une

37 Energy (Orbital) Diagram
4p 4s 3d 3p 3s E 2p 2s Be 1s2 2s2 1s

38 Orbital Box Diagrams Be 1s 2s 2p 3s

39 Orbital Box Diagrams N 1s 2s 2p

40 Formation of Cations electrons lost from subshell with highest n and l first (from valence electrons) examples K 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1 K+ 1s2 2s2 2p6 3s2 3p6 [Ar]

41 Ca 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2 Ca2+ 1s2 2s2 2p6 3s2 3p6 [Ar] Al 1s2 2s2 2p6 3s2 3p1 Al3+ [Ne] In [Kr] 4d10 5s2 5p1 In3+ [Kr] 4d10

42 Transition Metal Cations
In the process of ionization transition metals the ns electrons are lost before the (n-1)d Fe: [Ar] 3d6 4s2  Fe2+: [Ar] 3d6 Fe2+: [Ar] 3d6  Fe3+: [Ar] 3d5 Cu: [Ar] 3d10 4s1  Cu+: [Ar] 3d10 Cu+: [Ar] 3d  Cu2+: [Ar] 3d9 Fe, Fe2+, Fe3+, Cu, and Cu2+ are paramagnetic

43 Two problems of ions, charge, and electron configuration
An anion has a 3− charge and electron configuration 1s2 2s2 2p6 3s2 3p6. What is the symbol of the ion? The neutral atom has gained 3e- to form the ion, then the neutral atom had 15 e-. In the neutral atom the # e- = # p+ = Atomic number, that is 15. The element is, then, phosphorus (phosphorus). Symbol of ion is P3−. A cation has a 2+ charge and its electron configuration is [Ar] 3d7. What is the symbol of the ion? Here, the neutral atom has lost 2e-. It is a transition metal, due to the 3d electrons. Remember they firstly lose e-s in 4s orbital. Symbol of ion is Co2+. Neutral atom has = 27 e- = 27 p+ = atomic # [Ar] 3d7 lost

44 Atomic Properties and Periodic Trends
Periodic Properties of the Elements Atomic Radii Ionization Energy Electron Affinity Ionic Radii

45 Atomic Properties and Periodic Trends
Establish a classification scheme of the elements based on their electron configurations. Noble Gases All of them have completely filled electron shells. They are not very reactive. Since they have similar electronic structures, their chemical reactions are similar. He 1s2 Ne [He] 2s2 2p6 Ar [Ne] 3s2 3p6 Kr [Ar] 4s2 4p6 Xe [Kr] 5s2 5p6 Rn [Xe] 6s2 6p6

46 Atomic Properties and Periodic Trends
Representative Elements are the elements in A groups on periodic chart. These elements will have their “last” electron in an outer s or p orbital. These elements have fairly regular variations in their properties. Metallic character, for expl, increases from right to left and top to bottom.

47 Atomic Properties and Periodic Trends
d-Transition Elements Elements on periodic chart in B groups. Sometimes called transition metals. Each metal has d electrons. nsx (n-1)dy configurations These elements make the transition from metals to nonmetals. Exhibit smaller variations from row-to-row than the representative elements.

48 Atomic Properties and Periodic Trends
f - transition metals Sometimes called inner transition metals. Electrons are being added to f orbitals. Electrons are being added two shells below the valence shell! Consequently, very slight variations of properties from one element to another.

49 Atomic Properties and Periodic Trends
Outermost electrons (valence electrons) have the greatest Influence on the chemical properties of elements.

50 Atomic Properties and Periodic Trends
Atomic radii describe the relative sizes of atoms. Atomic radii increase within a column going from the top to the bottom of the periodic table. The outermost electrons are assigned to orbitals with increasingly higher values of n. The underlying electrons require some space, so the electrons of the outer shells must be further from the nucleus.

51 Atomic Properties and Periodic Trends
Atomic radii decrease within a row going from Left to right on the periodic table. This last fact seems contrary to intuition. How does nature make the elements smaller even though the electron number is increasing?

52

53 Atomic Radii The reason the atomic radii decrease across a period is due to shielding or screening effect. Effective nuclear charge, Zeff, experienced by an electron is less than the actual nuclear charge, Z. The inner electrons block the nuclear charge’s effect on the outer electrons. Moving across a period, each element has an increased nuclear charge and the electrons are going into the same shell (2s and 2p or 3s and 3p, etc.). Consequently, the outer electrons feel a stronger effective nuclear charge. For Li, Zeff ~ +1 For Be, Zeff ~ — For B, Zeff ~ +3

54 Atomic Radii Example: Arrange these elements based on their increasing atomic radii. Se, S, O, Te O < S < Se < Te In the same group atomic size increases as n (and Z) increases ─ Br, Ca, Ge, F F < Br < Ge < Ca same group same period

55 Ionization Energy First ionization energy (IE1)
The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1+ ion. Symbolically: Atom(g) + energy  ion+(g) + e Endothermic Mg(g) + 738kJ/mol  Mg+ + e IE1= 738kJ/mol

56 Ionization Energy Second ionization energy (IE2)
The amount of energy required to remove the second electron from a gaseous 1+ ion. Symbolically: ion+ + energy  ion2+ + e- Mg kJ/mol  Mg2+ + e IE2= 1451 kJ/mol Atoms can have 3rd (IE3), 4th (IE4), etc. ionization energies. The values are consecutively getting larger.

57 Ionization Energy Periodic trends for Ionization Energy:
1) IE2 > IE1 It always takes more energy to remove a second electron from an ion than from a neutral atom. 2) IE1 generally increases moving from IA elements to VIIIA elements. Important exceptions at Be & B, N & O, etc. due to s and p and half-filled subshells. 3) IE1 generally decreases moving down a family. IE1 for Li > IE1 for Na, etc

58 First Ionization Energies of Some Elements

59 Sr < Ca < Mg < Be
Ionization Energy Example: Arrange these elements based on their (increasing) first ionization energies. Sr, Be, Ca, Mg Sr < Ca < Mg < Be Al, Cl, Na, P Na < Al < P < Cl O, Ga, Sr, Se Sr < Ga < Se < O

60 Ionization Energy The reason Na forms Na+ and not Na2+ is that the energy difference between IE1 and IE2 is so large. Requires more than 9 times more energy to remove the second electron than the first one. The same trend is persistent throughout the series. Thus Mg forms Mg2+ and not Mg3+. Al forms Al3+ and not Al4+.

61 H 1312 Ionization Energies (kJ/mole)
He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K

62 Electron Affinity (EA)
Electron affinity is the amount of energy absorbed or emitted when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. Sign conventions for electron affinity. If EA > 0 energy is absorbed (difficult) If EA < 0 energy is released (easy) Electron affinity is a measure of an atom’s ability to form negative ions. Symbolically: atom(g) + e-  ion-(g) EA (kJ/mol)

63 Electron Affinity General periodic trend for electron affinity is
the values become more negative from left to right across a period on the periodic chart (affinity for electron increases). the values become more negative from bottom to top at a group on the periodic chart. −Noble gases have EA > 0 (full electron confg) An element with a high ionization energy generally has a high affinity for an electron, i.e., EA is largely negative. That is the case for halogens (F, Cl, Br, I), O, and S.

64 Electron Affinity F (Z= 9) and Cl (Z = 17) have the most negative EA Noble gases, He (2), Ne (10), and Ar (18), EA > 0; also Be, Mg, N They are all first Electron Affinity. A(g)- + e-  A2-(g) EA2(kJ/mol) is the 2nd

65 Electron Affinity Two examples of electron affinity values:
Mg(g) + e kJ/mol  Mg-(g) EA = 231kJ/mol Br(g) + e-  Br-(g) kJ/mol EA = -323 kJ/mol Br has a larger affinity for e− than Mg. The greater the affinity an atom has for an e− , the more negative EA is, the smaller it is.

66 Ionic Radii Cations (positive ions) are always smaller than their respective neutral atoms. When one or more electrons are removed, the attractive force of the protons is now exerted on less electrons. Element Na 11 p+, 11e- Mg 12p+, 12 e- Al 13 p+, 13e- Atomic Radius (Å) 1.86 1.60 1.43 Ion Na+ 11 p+, 10e- Mg2+ 12 p+, 10 e- Al3+ 13 p+, 10e- Ionic Radius (Å) 1.16 0.85 0.68

67 Ionic Radii nine electrons ten electrons
Anions (negative ions) are always larger than their neutral atoms. F 1s2 2s2 2p5 + e−  F− 1s2 2s2 2p6 same Z nine electrons ten electrons Element N 7 p+, 7e- O F Atomic Radius(Å) 0.75 0.73 0.72 Ion N3- 7 p+, 10e- O2- 8 p+, 10e- F− 9 p+, 10e- Ionic 1.71 The three 1.26 ions are 1.19 isoelectronic

68 Ionic Radii Cation (positive ions) radii decrease from left to right across a period. Increasing nuclear charge attracts the electrons and decreases the radius. Rb+ and Sr2+ are isoelectronic, same # of e-s Ion Rb+ Z = 37 p+ Sr2+ Z = 38 p+ In3+ Z = 49 p+ Ionic Radii(Å) 1.66 1.32 0.94

69 Ionic Radii Anion (negative ions) radii decrease from left to right across a period. Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius. For these isoelectronic anions… 10 e− and 7 p p p+ Ion N3- O2- F− Ionic Radii(Å) 1.71 1.26 1.19

70 Ionic Radii Example: Arrange these ions in order of decreasing radius.
Ga3+, K+, Ca2+ K+ > Ca2+ > Ga3+ Cl−, Se2−, Br−, S2− Se2− > Br− > S2− > Cl− isoelectronic isoelectronic, same # of electrons Se2−(34 p+) > Br−(35 p+); they have 36 e− each. S2−(16 p+) > Cl−(17 p+); they have 18 e− each. Br− > S2− because Br− is in the 4th period, S2− is in the 3rd.

71 Ionic Radii of isoelectronic species
Isoelectronic species have the same number of electrons. Here are some examples with the number of (protons) and + or − charges N3−(Z=7) > O2−(Z=8) > F−(Z=9) > Ne(Z=10) neutral > Na+(Z=11) > Mg2+(Z=12) > Al3+ (Z=13) all have 10e− The nuclear charge (+) increases from left to right, so does attraction force to electrons: r decreases. S2−(Z=16) > Cl− (Z=17) > Ar0 (Z=18) > K+ (Z=19) > Ca2+ (Z=20) > Sc3+ (Z=21) all of them have 18e−


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