Electron Configuration and Chemical Periodicity. Chapter Eight. AP Chemistry

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1 Electron Configuration and Chemical Periodicity Chapter Eight AP Chemistry

2 General Review Information about the Periodic Table

3 Organization of the Elements 1869: Dmitri Mendeleev - Published an organizational scheme for the elements. He called it the Periodic Table. Organized his elements by atomic mass in two rows of 7 elements and two rows of 17 elements. Noticed that the properties of the elements repeated so he put elements that behaved the same in the same vertical column.

4 Left blank spots on his periodic table and predicted the discovery and properties of several undiscovered elements. Problems were found with the table. Ar -- K Te -- I Co -- Ni

5 If Mendeleev s Table was to be correct, these elements should have been reversed. Elements could not be reversed because of their properties. This lead to a new version of the Periodic Table.

6 The Modern Periodic Law Henry Mosely - The properties of the elements are periodic when they are arranged in increasing order of their Atomic Number or Protons.

7 Patterns on the Periodic Table. Columns: Group or family - elements with similar properties. Several Groups have descriptive names.

8 Alkali Metals - Group IA React with water to form an alkaline or basic solution. Base Na (S) + H 2 O (L) --> NaOH (aq) + H 2(g)

9 Alkaline Earth Metals - Group IIA Elements also react with water to form a base.

10 Halogens - Group VIIA Salt Formers Na (S) + Cl 2(g) --> NaCl (S) Salt

11 Noble Gases - Group VIIIA Gases that are highly unreactive because their electron configuration is very stable.

12 Transition Metals Elements in the center section of the periodic table.

13 Inner Transition Metals Lanthanide Series Elements Ce - Lu Actinide Series Elements Th - Lr

14 Period Row on the periodic table. There are 7 periods on the periodic table. One period for each layer or energy level of electrons.

15 Characteristics of Multi-Electron Atoms Electrostatics plays a huge role in determining the energy stats of multi-electron atoms. By definition: electrons in the ground state occupy the orbitals of lowest energy. In multi-electron atoms, the energy states are determined by both the nucleus/electron attractions and the electron/electron repulsions. When the electron is far from the nucleus, the energy is higher (the system is less stable).

16 The higher the opposite charges, the stronger the attraction. - When a high charge nucleus attracts an electron, the energy is lower (the system is more stable) than when a lower nucleus of lower charge attracts an electron. Orbital energy is determined by finding the energy needed to remove an electron from that orbital in an atom. kj/mole energy needed to remove 1 mole of electrons from 1 mole of atoms.

17 More energy is needed to remove an electron from a low energy (more stable) orbital than is needed to remove an electron from a high energy (less stable) orbital. By definition, an atom s energy has negative value. Hence the lower (more negative) the energy is, the more stable the orbital.

18 The Effect of Nuclear Charge (Z) on Orbital Energy The high the nuclear charge lowers orbital energy by increasing the nucleus-electron attractions. 1s in H = kj/mole 1s in He +1 = kj/mole

19 Shielding: Effect of Electron Repulsions on Orbital Energy 1. Additional electron same orbital He to He +1 = kj/mole He +1 to He +2 = kj/mole Additional electrons raises the orbitals energy because of electron/electron repulsions. Each electron shields the others from receiving the full attraction to the nucleus. It is reduced somewhat to what is called an effective nuclear charge (Z eff ).

20 Shielding by other electrons makes an electron easier to remove. 2. Electrons located in inner orbitals shield electrons in outer orbitals. 3. Penetration or the probable time that the electrons spend close to the nucleus increases the attraction of the nucleus for those electrons in that orbital. Penetration, and the resulting shielding, causes an energy level to split into sublevels.

21 Therefore, the lower the value of l the lower the sublevel energy. s < p < d < f Energy Increases

22 Quantum Numbers A Quick Review n = 1 1 sublevel. 1s

23 n = 2 2s n = 3 2 sublevels. 2p 3 sublevels. 3s 3p 3d n = 4 4 sublevels. 4s 4p 4d 4f

24 Putting Them All Together: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f These sublevels are not needed. 6s 6p 6d 7s 7p

25 Diagonal Rule: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f Aufbau principle: to build up gives ground state electron configurations. 6s 6p 6d 7s 7p

28 Summarize l Letter Orbits Value of m 0 s 1 1 p 3 2 d 5 3 f 7 0-1, 0, +1-2, -1, 0, +1, +2-3, -2, -1, 0, +1, +2, +3

29 The Relation Between Orbital Filling and the Periodic Table

30 Hund s Rule Electrons will remain unpaired until all orbits in that sublevel have an electron. Aufbau Principle: add electrons 1 at a time to the lowest energy orbital available.

31 Short Hand Electron Configurations 07 Electron Configuration Polka.wma H 1s 1 Represents the number of electrons in the sublevel. He 1s 2

32 Li 1s 2 2s 1 B 1s 2 2s 2 2p 1 N 1s 2 2s 2 2p 3 Ne 1s 2 2s 2 2p 6

33 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Elements in the same group finish with an electron in the same position in the sublevel.

34 Cd 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 Hg 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10

35 What element ends in? 2p 3 N 5p 5 I 6s 2 Ba 4d 6 Ru 3d 5 Mn 4p 2 Ge

36 How does end? Cr 3d 4 Rn 6p 6 Mg 3s 2 Eu 4f 7 * Cl 3p 5 Sb 5p 3

37 Short- Short Hand (Condensed) Electron Configurations This method uses the Noble Gas that is just prior to it on the Periodic Table as a base and build from there. Mg {Ne} 3s 2

38 Cu {Ar} 4s 2 3d 9 * Te {Kr} 5s 2 4d 10 5p 4 Pd {Kr} 5s 2 4d 8 At {Xe} 6s 2 4f 14 5d 10 6p 5

39 Long Hand Electron Configurations or Orbital Diagrams

43 Orbital Box Diagram - IV : Sc Zn 4s 3d Z = 21 Sc [Ar] 4s 2 3d 1 Z = 22 Ti [Ar] 4s 2 3d 2 Z = 23 V [Ar] 4s 2 3d 3 Z = 24 Cr [Ar] 4s 1 3d 5 Z = 25 Mn [Ar] 4s 2 3d 5 Z = 26 Fe [Ar] 4s 2 3d 6 Z = 27 Co [Ar] 4s 2 3d 7 Z = 28 Ni [Ar] 4s 2 3d 8 Z = 29 Cu [Ar] 4s 1 3d 10 Z = 30 Zn [Ar] 4s 2 3d 10

44 H Rb Sr Cs Ba Fr Ra The Periodic Table of the Elements Anomalies to Electron Filling Li Be NaMg K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn He B C N O F Ne Al Si P S Cl Ar Ga Ge As Se Br Kr Y Zr NdMo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Ac Rf Du Sg Bo HaMe Ce Pr Nd PmSmEu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np PuAmCm Bk Cf Es FmMd No Lr Anomalous Electron Filling

45 Determining Quantum Numbers from Orbital Diagrams Write a set of quantum numbers for the third electron and a set for the eighth electron of the F atom. Use the orbital diagram to find the third and eighth electrons. 9 F 1s 2s 2p The third electron is in the 2s orbital. Its quantum numbers are n = 2 l = 0 m l = 0 m s = +1/2 The eighth electron is in a 2p orbital. Its quantum numbers are n = 2 l = 1 m l = -1 m s = -1/2

46 Give the full and shorthand electrons configurations, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements: (a) potassium (K: Z = 19) full configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

47 shorthand configuration: [Ar] 4s 1 partial orbital diagram: 4s 1 n = l = m l = m s = There are 18 inner electrons.

48 (b) for Mo (Z = 42) full configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 5 shorthand configuration: [Kr] 5s 1 4d 5

49 partial orbital diagram: 5s 1 4d 5 n = l = m l = m s = There are 36 inner electrons and 6 valence electrons.

50 (c) for Pb (Z = 82) full configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 2 shorthand configuration [Xe] 6s 2 4f 14 5d 10 6p 2 partial orbital diagram 6s 2 6p 2 n = l = m l = m s = There are 78 inner electrons and 4 valence electrons.

51 Orbital Occupancy for the First 10 elements, H through Ne. Atoms are Spherical in Shape.

52 Trends in Some Important Periodic Atomic Properties

53 All physical and chemical behavior of the elements is based ultimately on the electron configurations of their atoms. Trends in Atomic Size Trends in Ionization Energy Trends in Electron Affinity These trends are periodic.

54 Trends in Atomic Size Definition is based on how closely one atom lies next to another identical atom. Generally, the distance between the atoms is measured and then divided by two to give the atomic radius.

55 Metallic radius: is one half the distance between the nuclei of adjacent atoms in a crystal of the element. I ii

56 For elements commonly occurring as molecules, mostly nonmetals, atomic size is defined by the covalent radius. Covalent radius: is one half the distance between the nuclei of identical covalently bonded atoms.

58 Trends in the Main Group Elements 1. As n increases, the outer electrons spend more time farther from the nucleus. Hence, the atoms are larger. 2. As the effective nuclear charge (Z eff )- the positive charge felt by an electron increases, outer electrons are pulled closer to the nucleus. Hence the atoms are smaller.

59 The net effect: 1. Down a group, n dominates. Inner electrons shield outer electrons very effectively. Hence, atomic radius generally increases down a group. 2. Across a period, Z eff dominates. As electrons are added to the same outer level, the shielding of the inner electrons does not change. Z eff rises and the outer electrons are pulled closer. Hence, atomic radius generally decreases across a period (L to R).

60 i

61 Ranking Elements by Atomic Size Using only the periodic table, rank each set of main group elements in order of decreasing atomic size: (a) Ca, Mg, Sr These elements are in Group 2A(2). Sr > Ca > Mg (b) K, Ga, Ca K > Ca > Ga These elements are in Period 4.

62 (c) Br, Rb, Kr Rb > Br > Kr (d) Sr, Ca, Rb Rb > Sr > Ca Rb has a higher energy level and is far to the left. Br is to the left of Kr. Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

64 Ranking Elements by Size Problem: Rank the following elements in each group according to decreasing size ( largest first!): a) Na, K, Rb b) Sr, In, Rb c) Cl, Ar, K a) Rb > K > Na These elements are all alkali metals and the elements increase in size as you go down the group. b) Rb > Sr > In These elements are in Period 5 and the size decreases as you go across the period. c) K > Cl > Ar These elements border a noble gas, and it is the smallest diameter.

65 Trends in Ionization Energy Ionization Energy (IE) is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms or ions. Removing an electron always requires energy to break the attraction of the electron towards the nucleus. IE is always positive since energy flows into the system like an endothermic reaction.

66 IE 1 removes an outermost electron (highest energy sublevel). Atom (g) ion + (g) + e- E = IE 1 > 0 IE 2 removes a second electron. Ion (g) ion 2+ (g) + e- E = IE 2 (always > IE 1 ) Atoms with a low IE 1 tend to form positive ions (cations) during reactions while those with a high IE 1 (except the Noble Gases) often form negative ions (anions)

67 Variations in IE 1 Generally, as the size of the atom increases, it takes more energy to remove it. 1. Down a group: As the distance between the nucleus and the outer electron increases, IE decreases. Exception: IE in group IIIA Filling of the 4d 6d sublevels causes a greater than expected Z eff which holds the electrons more tightly in the larger 3A elements

68 2. Across a period: Z eff increases so atomic size decreases. This causes the attraction between the nucleus and the outer electron to increase. Hence, IE generally increases across a period. Exception: Stability of full and half-full sublevels. 1. Between Group 2A - Group 3A Full s sublevel to a p 1 configuration. p sublevels are at a higher energy state than s sublevels and are hence removed easier.

69 2. Between Group 5A - Group 6A p 3 configuration p 4 configuration. p 4 is the first to pair up and the electron/electron repulsion raises the energy level of the orbital. Removing p 4 yields p 3 which relieves the repulsion and leaves a stable half-filled sublevel.

70 Periodicity of First Ionization Energy (IE 1 )

71 IE 1 of the maingroup elements

72 Ranking Elements by First Ionization Energy Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE 1 : (a) Kr, He, Ar IE decreases as you proceed down in a group; IE increases as you go across a period. He > Ar > Kr Group 8A- IE decreases down a group.

73 (b) Sb, Te, Sn Te > Sb > Sn Period 5 elements - IE increases across a period. (c) K, Ca, Rb Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I, Xe, Cs I is to the left of Xe; Cs is further to the left and down one period.

74 Variations in Successive Ionization Energies Successive I.E. (IE 1, IE 2, ) of a given element increase because each electron is pulled away from an ion with a higher and higher positive charge.

75 Note: Increases in I.E. is not smooth. A jump appears after the outer (valence) electrons are removed. Much energy is needed to remove an inner (core) electron. The first three ionization energies of beryllium (MJ/mole)

77 Identifying an Element from Successive Ionization Energies Name the Period 3 element with the following ionization energies (in kj/mol) and write its electron configuration: IE 1 IE 2 IE 3 IE 4 IE 5 IE ,230 Look for a large increase in energy which indicates that all of the valence electrons have been removed. The largest increase occurs after IE 5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s 2 3p 3 and the element must be phosphorous, P (Z = 15). The complete electron configuration is 1s 2 2s 2 2p 6 3s 2 3p 3.

78 Trends in Electron Affinity Electron Affinity (EA) is the energy change accompanying the ADDITION of 1 mole of electrons to 1 mole of gaseous atoms or ions. Atom (g) + e- ion - (g) E = EA 1 In most cases, energy is released when the first electron is added because of the attraction to the nucleus. EA 1 is usually negative just like H is negative for exothermic reactions.

79 Electron Affinities of the Main-Group Elements

80 EA 2 is always positive (energy absorbed) because electrons are being added to something that is already negatively charged. There is no clear cut trend for EA. Reactive nonmetals: Have IEs and very negative (exothermic) EAs. These elements gain electrons with ease and lose them with difficulty. Reactive metals: Have IEs and slightly negative or slightly positive (endothermic) EAs. These elements lose electrons with ease and gain them with relative difficulty.

81 Noble Gases: Have IEs and slightly positive (endothermic) EAs. These elements tend to neither lose or gain electrons.

82 Trends in Three Atomic Properties

83 Trends in Metallic Behavior Metals tend to lose electrons during chemical reactions because of their low IEs as compared to nonmetals. Hence, the elements with the most metallic behaviors are those on the left and towards the bottom of the periodic table.

84 Trends in Metallic Behavior

85 Acid - Base Behavior of Element Oxides Most main group metals transfer electrons to oxygen so their oxides are ionic. When added to water, these oxides act as bases. CaO (s) + H 2 O (l) Ca(OH) 2(aq) + H 2(g)

86 Nonmetals share electrons with oxygen so nonmetal oxides are covalent. When added to water, these oxides act as acids. P 4 O 10(s) + H 2 O (l) H 3 PO 4(aq)

87 The change in metallic behavior in Group 5A and Period 3. Basic Oxides The trend in acid-base behavior of element oxides. Acidic Oxides

88 Isoelectronic Atoms and Ions Isoelectronic: same with the nearest noble gas. H - 1 { He } Li + Be +2 N - 3 O - 2 F - { Ne } Na + Mg +2 Al +3 P - 3 S - 2 Cl - { Ar } K + Ca +2 Sc +3 Ti +4 As - 3 Se - 2 Br - { Kr } Rb + Sr +2 Y +3 Zr +4 Sb - 3 Te - 2 I - { Xe } Cs + Ba +2 La +3 Hf +4

90 Metals in Groups 3A(13) to 5A(15) typically can not lose enough electrons to become isoelectronic with a noble gas. Pseudo-noble gas configurations: elements that empties its outer s and p sublevels so that only a full d sublevel remains. Sn ([Kr]5s 2 4d 10 5p 2 ) Sn 4+ ([Kr]4d 10 ) + 4e - Inert pair: full s sublevel electrons are very difficult to remove. (stability of full s and d sublevels. Sn ([Kr]5s 2 4d 10 5p 2 ) Sn 2+ ([Kr]5s 2 4d 10 ) + 2e -

91 Electron Configurations of Main-Group Ions Using condensed electron configurations, write reactions for the formation of the common ions. (a) Iodine Group VIIA gains 1 electron to become isoelectronic with xenon. I [Kr] 5s 2 4d 10 5p 5 + e - I - [Kr] 5s 2 4d 10 5p 6 (b) Potassium Group IA loses 1 electron to become isoelectronic with argon. K [Ar] 5s 1 K + [Ar] + e -

92 (c) Indium Group IIIA loses 3 electrons to become In 3+ (pseudo-noble gas) or it can lose 1 electron to form In 1+ (an inert pair). In [Kr] 5s 2 4d 10 5p 1 In 3+ [Kr]4d e - In [Kr] 5s 2 4d 10 5p 1 In 1+ [Kr]5s 2 4d e -

93 Pseudo - Noble Gas Electron Configurations Elements in groups 3A, 4A, and 5A can form cations by losing enough electrons to leave a pseudo noble gas configuration. By losing electrons and leaving a filled d orbital, which is quite stable! Sn [Kr] 5s 2 4d 10 5p 2 Sn 4+ [Kr] 4d e - Sn [Kr] 5s 2 4d 10 5p 2 Sn 2+ [Kr] 5s 2 4d e - Pb [Xe] 4f 14 5d 10 6s 2 6p 2 Pb +2 [Xe] 4f 14 5d 10 6s e - Pb [Xe] 4f 14 5d 10 6s 2 6p 2 Pb +4 [Xe] 4f 14 5d e - As [Ar] 3d 10 4s 2 4p 3 As 3+ [Ar] 3d 10 4s e - As [Ar] 3d 10 4s 2 4p 3 As 5+ [Ar] 3d e - Sb [Kr] 4d 10 5s 2 5p 3 Sb 3+ [Kr] 4d 10 5s e - Sb [Kr] 4d 10 5s 2 5p 3 Sb 5+ [Kr] 4d e -

94 Magnetic Properties Paramagnetic - An atom or ion which has unpaired electrons, which add up to give a spin vector. They are thereby attracted by a magnetic field. Diamagnetic - An atom or ion with all electrons paired and with no net spin. These atoms or ions are not attracted by a magnetic field.

95 Apparatus for Measuring the Magnetic Behavior of a Sample

96 Elements and Ions that are Para/dia-magnetic. Spectral analysis says that Ti and Ti 2+ should have the following electron configurations. Notice, the two s electrons are gone.

97 Tests for paramagnetism supports both electron configurations. Ti loses the two 4s electrons instead of the two unpaired d electrons.

98 An increase in paramagnetism occurs when Fe metal forms Fe 3+ compounds. Fe loses the two 4s electrons and the paired electron in the 3d sublevel.

99 Cu is paramagnetic, but Zn is diamagnetic, as are Cu + and Zn 2+ ions. (The two ions are isoelectronic)

100 Writing Electron Configurations and Predicting Magnetic Behavior of Transition Metal Ions Use condensed electron configurations to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic.

101 Write the electron configuration and remove electrons starting with ns to match the charge on the ion. If the remaining configuration has unpaired electrons, it is paramagnetic. Mn 2+ : Mn([Ar]4s 2 3d 5 ) Mn 2+ ([Ar] 3d 5 ) + 2e - paramagnetic

102 Cr 3+ Cr([Ar]4s 2 3d 6 ) paramagnetic Cr 3+ ([Ar] 3d 5 ) + 3e - Hg 2+ : Hg([Xe]6s 2 4f 14 5d 10 ) Hg 2+ ([Xe] 4f 14 5d 10 ) + 2e - not paramagnetic (is diamagnetic)

103 Ionic Size vs Atomic Size Cations (+) are smaller than in their atom form. Electrons are removed from their outer level. Less electron repulsions allow the nucleus to pull the electrons closer.

104 Anions (-) are larger than in their atom form. Electrons are added to their outer level. Increased electron repulsions cause the electrons cloud to enlarge.

105

106 Ionic vs. atomic radius

107 Ionic size increases down a group. Ionic size decreases across a period but increases from cation to anion. Ionic size decreases with increasing positive or decreasing negative charge. Ionic size decreases as charge increases for different cations of an element.

108 Ranking Ions by Size Rank each set of ions in order of decreasing size, and explain your ranking: Ca 2+, Sr 2+, Mg 2+ Sr 2+ > Ca 2+ > Mg 2+ These are members of the same Group (2A/2) and therefore decrease in size going up the group.

109 K +, S 2-, Cl - S 2- > Cl - > K + The ions are isoelectronic; S 2- has the smallest Z eff and therefore is the largest while K + is a cation with a large Z eff and is the smallest. Au 3+, Au + Au + > Au 3+ The higher the + charge, the smaller the ion.

110 Ranking Ions According to Size Rank each set of Ions in order of increasing size. a) K +, Rb +, Na + b) Na +, O 2-, F - c) Fe +2, Fe +3 i) size increases down a group, ii) size decreases across a period but increases from cation to anion. iii) size decreases with increasing positive (or decreasing negative) charge in an isoelectronic series. iv) cations of the same element decreases in size as the charge increases. a) since K +, Rb +, and Na + are from the same group (1A), they increase in size down the group: Na + < K + < Rb + b) the ions Na +, O 2-, and F - are isoelectronic. O 2- has lower Z eff than F -, so it is larger. Na + is a cation, and has the highest Z eff, so it is smaller: Na + < F - < O 2- c) Fe +2 has a lower charge than Fe +3, so it is larger: Fe +3 < Fe +2

111 The End

Made the FIRST periodic table

Made the FIRST periodic table 1869 Mendeleev organized the periodic table based on the similar properties and relativities of certain elements Later, Henri Moseley organized the elements by increasing