CHEMISTRY - BROWN 13E CH.7 - PERIODIC PROPERTIES OF THE ELEMENTS

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2 CONCEPT: EFFECTIVE NUCLEAR CHARGE & SLATER S RULES When looking at any particular electron within an atom it experiences two major forces. A(n) force from the nucleus and a(n) force from the surrounding electrons. Now the electron can become shielded from the full force of the nucleus because of the other surrounding electrons. Effective Nuclear Charge (Zeff) measures the force exerted onto an electron by the nucleus, and can be calculated using Slater s Rules. e - e - e - e - e - Z = Nuclear Charge e - e - Z eff = Z S S = Shielding Constant e - e - e - e - Guidelines for Determining S for an electron: 1. The atom s electronic configuration is grouped as follows, in terms of increasing n and l quantum numbers: (1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) (5d) etc. 2. Electrons in groups to the right of a given electron do not shield electrons to the left. 3. The shielding constant S for electrons in certain groups. For ns and np valence electrons: a) Each electron in the same group will contribute to the S value. A 1s electron contributes to the S value for another 1s electron. b) Each electron in n 1 group contributes to the S value. c) Each electron in n 2 group or greater contributes to the S value. For nd and nf valence electrons: a) Each electron in the same group will contribute to the S value. b) Each electron in groups to the left will contribute to the S value. EXAMPLE: Using Slater s Rules calculate the effective nuclear charge of a 3p electron in argon. Page 2

3 PRACTICE: EFFECTIVE NUCLEAR CHARGE & SLATER S RULES 1 EXAMPLE 1: Using Slater s Rules calculate the effective nuclear charge of the 4s electron in potassium. EXAMPLE 2: Using Slater s Rules calculate the effective nuclear charge of a 3d electron in bromine. Page 3

4 CONCEPT: TRENDS IN ATOMIC RADIUS Atomic radius is defined as half the distance between the nuclei in a molecule of two identical elements. Generally, it going from left to right across a period and going down a group. ATOMIC RADIUS EXAMPLE: If the sum of the atomic radii of diatomic carbon is 154 pm and of diatomic chlorine is 198 pm, what is the sum of the atomic radii between a carbon and a chlorine atom. PRACTICE: Which one of the following atoms has the largest atomic radius? A) K B) Rb C) Y D) Ca E) Sr Page 4

5 CONCEPT: TRENDS IN IONIC RADIUS Ionic Size estimates the size of an ion in an ionic compound. (POSITIVE IONS) tend to be smaller than their parent atoms. Lithium ( 3 Electrons) 1s 2s 1s 2s (NEGATIVE IONS) tend to be larger than their parent atoms. Fluorine ( 9 Electrons) 1s 2s 2p 1s 2s 2p The pattern for ionic size correlates with the following trend when comparing ions with the same number of electrons: -3 > -2 > -1 > 0 > +1 > +2 > +3 EXAMPLE: Rank each set of ions in order of increasing ionic size. a) K +, Ca 2+, Ar b) Sr 2+, Na +, I c) V 5+, S 2-, Cl Page 5

6 CONCEPT: TRENDS IN IONIZATION ENERGY Metals tend to lose electrons to become positive ions called. IONIZATION ENERGY Therefore they have ionization energies. Nonmetals tend to gain electrons to become negative ions called. Therefore they have ionization energies. Ionization energy (IE) is the energy (in kj) required to remove an electron from a gaseous atom or ion. Generally, it going from left to right of a period and going down a group. Exceptions: Atom (g) ion + (g) + e E = IE1 > 0 When in the same period, Group elements have lower ionization energy than elements in Group. O 1s 2s 2p 1s 2s 2p N 1s 2s 2p 1s 2s 2p When in the same period, Group elements have lower ionization energy than elements in Group. B 1s 2s 2p 1s 2s 2p Be 1s 2s 1s 2s Page 6

7 PRACTICE: TRENDS IN IONIZATION ENERGY EXAMPLE: Of the following atoms, which has the smallest second ionization energy? a. Al b. Li c. Rb d. Mg e. Be PRACTICE 1: Of the following atoms, which has the smallest third ionization energy? a. Al b. Ca c. K d. Ga e. Cs PRACTICE 2: Which of the following statements is/are true? a. Sulfur has a larger IE1 than phosphorus b. Boron has a lower IE1 than Magnesium c. Magnesium has a higher IE1 than Aluminum PRACTICE 3: Shown below are the numerical values for ionization energies (IE s). Match the numerical values with each of the following elements provided in the boxes. Na Mg Al Si P S Cl Ar Numbers: 496, 578, 738, 786, 1000, 1012, 1251 & Page 7

8 CONCEPT: TRENDS IN ELECTRON AFFINITY Electron Affinity (EA) is the energy change (in kj) from the addition of 1 mole of e to 1 mol of gaseous atoms or ions. Generally, it going from left to right across a period and going down a group. Atom (g) + e ion (g) E = - EA1 ELECTRON AFFINITY EXAMPLE: Rank the following elements in order of increasing electron affinity. a. Cs, Hg, F, S b. Se, S, Si PRACTICE: Shown below are the numerical values for electron affinities (EA s). Match the numerical values with each of the following elements provided in the boxes. Li Be B C N O F Ne Numbers: - 328, -141, -122, -60, -27, > 0, > 0, > 0. Page 8

9 CONCEPT: BORN-HABER CYCLE The Born-Haber cycle is used a method to calculate the or of a compound. It looks mainly at the formation of an ionic compound from gaseous ions. The metal being from Groups or and the nonmetallic element being a or. M (s) X 2 o ΔH f MX (s) 1 = = M (g) 2 X (g) 4 3 = M + (g) + X (g) HX (s) 5 4 = ΔH o f = = Page 9

10 PRACTICE: BORN-HABER CYCLE EXAMPLE: Using the Born-Haber Cycle, demonstrate the formation of cesium chloride, CsCl, and calculate its heat of formation. ΔH Sublimation = 79 kj mol IE 1 = 376 kj mol ΔH Dissociation =122 kj mol EA = 349 kj mol U = 661 kj mol Page 10

11 8. Which of the following transitions (in a hydrogen atom) represent emission of the smallest or shortest wavelength? a. n = 4 to n = 2 b. n = 3 to n= 4 c. n = 1 to n = 2 d. n = 7 to n = 5 e. n = 2 to n = 5 Page 11

12 9. Which of the following transitions represent absorption of a photon with the highest frequency? a. n = 3 to n = 1 b. n = 2 to n = 4 c. n = 1 to n =2 d. n = 6 to n = 3 e. n = 1 to n = 3 Page 12

13 10. Provide the n, l and ml value for each of the given orbitals. a) 7s n = b) 5d n = l = l = ml = ml = c) 2p n = d) 4f n = l = l = ml = ml = Page 13

14 11. Which statement about the four quantum numbers is false? a. n = principal quantum number, n = 1 to b. l = azimuthal quantum number, l = 0,1,2,..., (n+1) c. ml = magnetic quantum number, ml = (-l),...,0,..., (+l) d. ms = spin quantum number, ms = or 1 2 e. The first three quantum numbers deal with the atomic orbitals except for the ms quantum number, which deals with the electrons in the atomic orbitals. Page 14

15 12. Each of the following sets of quantum numbers gives information on a specific orbital. Find the error in each. a. n = 4, l = 0, ml = 1, ms = 1 2 b. n = 5, l = 2, ml = - 1, ms = 1 c. n = 7, l = 7, ml = - 5, ms = 1 2 d. n = 0, l = 5, ml = - 3, ms = 1 2 Page 15

16 14. How many electrons can have the following quantum sets? a) n = 4, ml = -1 b) n = 5, ml = 0, ms = 1 2 c) n = 9, l = 4, ms = 1 2 d) n = 2, ms = 1 2 Page 16

17 19. For n = 2, what are the possible sublevels? a) 0 b) 0, 1 c) 0, 1, 2 d) 0, 1,2, 3 Page 17

18 16. Based on the following atomic orbital shape, which of the following set of quantum numbers is correct: a) n = 2, l = 1, ml = 0 b) n = 3, l = 2, ml = 1 c) n = 4, l = 0, ml = +1 d) n = 1, l = 1, ml = 0 Page 18

19 17. Based on the following atomic orbital shape, which of the following set of quantum numbers is correct: a) n = 3, l = 2, ml = 0, ms = 1 2 b) n = 3, l = 1, ml = - 3, ms = 1 c) n = 4, l = 0, ml = 0, ms = 1 2 d) n = 4, l = 2, ml = - 3, ms = 1 2 Page 19

20 18. Based on the following atomic orbital shape, which of the following set of quantum numbers is correct: a) n = 3, l = 3, ml = 0, ms = 1 2 b) n = 1, l = 3, ml = -3, ms = 1 c) n = 7, l = 3, ml = - 4, ms = 1 2 d) n = 6, l = 3, ml = -3, ms = 1 2 Page 20

21 25. Give the electron configuration for the following element and its ion. For the ion, state if it is paramagnetic or diamagnetic: a. Ag Ag + Page 21

22 26. Give the electron configuration for the following element and its ion. For the ion, state if it is paramagnetic or diamagnetic: a. Cl Cl + Page 22

23 27. Which of the following represents an excited state? a) Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 b) Be: 1s 2 2s 2 c) Na: 1s 2 2s 2-2p 6 3p 1 d) N: 1s 2 2s 2 2p 3 Page 23

24 28. Give the set of four quantum numbers that represent the indicated electron in the following element: a. Br (33 rd electron) n =, l =, ml =, ms = Page 24

25 29. Give the set of four quantum numbers that represent the indicated electron in the following element: a. Ca (19 th electron) n =, l =, ml =, ms = Page 25

26 30. Give the set of four quantum numbers that represent the indicated electron in the following element: a. Cu (27 th electron) n =, l =, ml =, ms = Page 26

27 31. Give the set of four quantum numbers that represent the indicated electron in the following element: a. Mo 3+ (38 th electron) n =, l =, ml =, ms = Page 27

28 32. For a multi-electron atom, arrange the electron subshells of the following listing in order of increasing energy: 6s, 4f, 2p, 5d. Page 28

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