Chapter 8 The Periodic Table: Structure and Trends

1. Chapter 8The Periodic Table: Structure and Trends

2. Electron Configurations and the Periodic Table • The periodic table can be divided into four blocks of elements: elements with highest energy electrons in s, p, d, or f subshells. • The arrangement of the elements in the periodic table correlates with the subshell that holds the highest energy electron.

3. Example: Electron Configurations • Using only the periodic table, determine the electron configurations of Al, Ti, Br, and Sr.

4. Electron Configurations of Anions • For anions, the additional electrons fill orbitals following the same rules that applies to atoms. Cl: [Ne] 3s2 3p5 Cl-: [Ne] 3s2 3p6 As: [Ar] 4s2 3d10 4p3 As3-: [Ar] 4s2 3d10 4p6 • Many stable anions have the same electron configuration as a noble gas atom.

5. Electron Configurations of Cations • For the electron configurations of cations, electrons of highest n value are removed first. For cases of the same n level, electrons are first removed from the subshell having highest l. As: [Ar] 4s2 3d10 4p3 As3+: [Ar] 4s2 3d10 Mn: [Ar] 4s2 3d5 Mn2+: [Ar] 3d5 • NOTE: For d-block atoms, the ns electrons are removed before the (n-1)d electrons.

6. Test Your Skill • Write the electron configurations of the following ions: (a) N3- (b) Co3+ (c) K+

7. Isoelectronic Series • An isoelectronic series is a group of atoms and ions that contain the same number of electrons. • The species S2-, Cl-, Ar, K+, and Ca2+ are isoelectronic – they all have 18 electrons.

8. Atomic Radii • An atomic radius is one half the distance between adjacent atoms of the same element in a molecule. 198/2 = 99 228/2 = 114 Sum = 215

9. Size Trends for an Isoelectronic Series

10. Sizes of the Atoms and Their Cations • Atoms are always larger than their cations.

11. Sizes of the Atoms and Their Cations • If an atom makes more than one cation, the higher-charged ion has a smaller size.

12. Atomic and Ionic Radii • Anions are always larger than their atoms.

13. Test Your Skill • Identify the larger species of each pair: (a) Mg or Mg2+ (b) Se or Se2-

14. Atomic Radii of Main Group Elements

15. Sizes of Atoms • The sizes of atoms are impacted by the effective nuclear charge felt by the outermost electrons.

16. Effective Nuclear Charge & Size • The sizes of atoms increase going down a group.

17. Sizes of Atoms • The increase in effective nuclear charge causes a size decrease across the period.

18. Test Your Skill • Identify the larger species of each pair: (a) Mg or Na (b) Si or C

19. Ionization Energy • The ionization energy is the energy required to remove an electron from a gaseous atom or ion in its electronic ground state.

20. Ionization Energies • An atom has as many ionization energies as it has electrons. • Example: Mg(g) → Mg+(g) + e- I1 = first ionization energy Mg+(g) → Mg2+(g) + e- I2 = second ionization energy

21. Trends in 1st Ionization Energies • The increase in the effective nuclear charge across a period causes an increase in the ionization energy as you go across that period.

22. Trends in 1st Ionization Energies

23. Trends in 1st Ionization Energies • The slight dip in ionization energy for O is because the fourth p electron now pairs with another electron, slightly repelling each other.

24. Trends in First Ionization Energies

25. Ionization Energy Trends in Isoelectronic Series • Isoelectronic species with the greatest charge in the nucleus will have the largest ionization energy. • For the isoelectronic series S2-, Cl-, and Ar, Ar has the largest ionization energy because it has the most protons in its nucleus.

26. Ionization Energy • Predict which species in each pair has the higher ionization energy. (a) Ca or As (b) K+ or Ca2+ (c) N or As

27. Successive Ionization Energies • Successive ionization energies always increase because of the increasing hold the nucleus has on remaining electrons. I1I2I3I4 Mg 738 1450 7734 10550 Al 578 1817 2745 11600 • A much larger increase is seen when an electron comes from a lower-energy subshell. (all values in kJ/mol)

28. Test Your Skill • Which element, magnesium or sodium, has the greater second ionization energy?

29. Electron Affinity • The electron affinity of an element is the energy change that accompanies the addition of an electron to a gaseous atom to form an anion. A(g) + e- → A-(g) • Electron affinities are generally favorable (exothermic) for elements on the right side of the periodic table.

30. Electron Affinities

31. Alkali Metals – Group 1A (1) • The reactivity of the Group 1A metals increases down the group. Their chemistry is dominated by the formation of M+ ions. 2M(s) + H2O(l) → 2MOH(aq) + H2(g) 2M(s) + H2(g) → 2MH(s) 2M(s) + X2(g) → 2MX(s) X = F, Cl, Br, I

32. Alkali Metal Reactions with O2 • Only lithium reacts with O2 to give the expected product, lithium oxide. 4Li(s) + O2(g) → 2Li2O(s) • Sodium reacts mainly to yield sodium peroxide. 2Na(s) + O2(g) → Na2O2(s) • Potassium reacts to yield mixtures of the oxide, peroxide, and superoxide. K(s) + O2(g) → KO2(s)

33. Flame Colors of the 1A Elements

34. The Alkaline Earth Metals – Group 2A (2) • The Group 2A metals are not as reactive as the Group 1A metals. Reactivity increases down the group, and they all form M2+ ions. • Magnesium alloys are useful in aeronautical applications, where low density and high strength are important.

35. Flame Colors of 2A Elements Calcium Strontium Barium

36. The Halogens – Group 7A (17) • The halogens all exist as diatomic molecules, but they are very reactive. • The reactivity decreases as you go down the group. Their chemistry is dominated by the formation of X- ions. • The interhalogens are compounds formed from different halogens, like IF3 and BrCl.