Lecture 08 (Chapter 8) The Periodic Table: Structure and Trends

1. Lecture 08 (Chapter 8)The Periodic Table: Structure and Trends

2. Periodic Trends of the Elements • The experimental trends in group properties on which the periodic table was based can now be explained by the arrangements of electrons in atoms. • The electron configurations of the valence electrons in each group member are similar. • Valence electrons: Electrons with highest principle quantum number in atom, and any electrons in an unfilled subshell from a lower shell. • Valence orbitals: Orbitals of the highest principle quantum number in atom, and (for d or f electrons), the orbitals of any partially-filled subshells of lower principle quantum numbers.

3. 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.

4. Electron Configurations and the Periodic Table • The ordering of orbitals with respect to energy is reflected in the periodic table (i.e., we can count electrons on the periodic table to determine addresses for all of the electrons in the electron configuration). • Elements with one electron in a new principal shell (ns1) start a new period in the periodic table.

5. Electron Configurations • The 4s orbital is lower in energy than the 3d orbital and fills first, starting the fourth period at potassium. • The 3d orbitals fill after the 4s. • Similar inversions occur in the remaining periods.

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

7. 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.

8. 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. • See Pb2+ as example.

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

10. Test Your Skill • Write the electron configurations of the following ions: (a) N3- (b) Co3+ (c) K+ • Answers: (a) 1s2 2s2 2p6 (b) [Ar] 3d6 (c) [Ar]

11. 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. • Would these species be stable, and why?

12. Atomic Radii • Since an electron cloud is a “fuzzy” probability function, the atomic radius can be calculated based on the half distance between adjacent atoms of the same element in a molecule. • This method does not work as well for metal atoms. 198/2 = 99 228/2 = 114 Sum = 213

13. Sizes of the Atoms and Their Cations • Atoms are always larger than their cations. • If the electrons are removed from an orbital, then there is less probability of finding an electron in that orbital. • If an atom makes more than one cation, the higher-charged ion has a smaller size.

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

15. Size Trends for an Isoelectronic Series

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

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

18. Sizes of Atoms • The sizes of atoms are impacted by the effective nuclear charge (Zeff) felt by the outermost electrons.

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

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

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

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

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

24. Property trends Atomic radius increases Ionization energy decreases Electronegativity decreases • Atomic radius decreases with addition of protons to nucleus which pulls e- closer to center (Electronegativity). • Decreasing radius makes it more difficult to remove e- (ionization energy is energy required to remove e-) Small radius e- held tightly Large radius e- held loosely

25. 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

26. 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. • Exceptions: • Group 3A. The np1 electron does not penetrate inner electrons as much as ns2 electrons. • Group 6A. First pairing of electrons in p orbital produces small repulsion between electrons.

27. 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 (therefore, the most positive charge) in its nucleus.

28. 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

29. 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. • Based on periodic table, do these numbers make sense? (all values in kJ/mol)

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

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

32. Electron Affinity • The electron affinity of an element is the energy change (in kJ/mol) 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 (i.e., these non-metal elements are more likely to gain electrons than lose them).

33. Electron Affinities

34. Alkali Metals – Group 1A (1) • Group 1A metals very reactivie, and reactivity increases down the group. Their chemistry is dominated by the formation of M+ ions (easier to remove a single electron as atomic radius increases). 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

35. Alkali Metal Reactions with O2 • Reactions of Group 1A metals with molecular oxygen do not necessarily follow assumed trends. • 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)

36. 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 (loss of 2 electrons results in a noble gas electron configuration).

37. The Halogens – Group 7A (17) • The halogens all exist as diatomic molecules, but they are very reactive, and tend to gain 1 electron to achieve a noble gas configuration. • The reactivity decreases as you go down the group. • The interhalogens are compounds formed from different halogens, like IF3 and BrCl.

38. Summary • As Atomic radius increases, electrons are held more distantly from protons, and are easier to remove as the effective nuclear charge is reduced. • Metals (with higher atomic radii) tend to lose enough electrons to achieve a noble gas electron configuration. • Once they have done so, it is more difficult to remove more electrons (requires higher ionization energy). • Non-metals (with smaller atomic radii) tend to gain enough electrons to achieve a noble gas electron configuration. • Once they have done so, it is more difficult to add more electrons (electron affinity decreases as they are already in a stable octet)