The Modern Periodic Table Chapter 6

1. The Modern Periodic TableChapter 6

2. Objectives for Friday: • 4. Define period, group, and family. • 5. Explain the different systems for numbering the groups on the Periodic table. • 6. Name and locate on the periodic table: metals, nonmetals, metalloids, noble gases, alkali metals, alkaline earth metals, halogens, chalcogens, Lanthanides, Actinides, Transition metals, and Inner transition metals.

3. Objectives: • 7. Apply the octet rule in predicting the stability of an element. • 8. Predict the electron configuration of the outer energy level of any element from its position on the Periodic table. • 9. Recognize exceptions to the Aufbau order due to half-full or full d-sublevels. • 10. Define atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity. • 11. Predict trends in atomic radius, ionic radius, oxidation number, ionization energy, electron affinity, and electronegativity. • 12. Explain trends in terms of nuclear charge and shielding effects.

4. Arrangement and Nomenclature • Rows are called periods • Columns are designated as groups • Each column in the main table and each row at the bottom is also designated an individual family • Groups 1A, 2A, and 3-8A are the main groups, or representative elements • Groups 1B-8B are called the transition elements

5. The Periodic Table With Atomic Symbols, Atomic Numbers, and Partial Electron Configurations

6. Broad Periodic Table Classifications • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O) • Transition Elements: filling dorbitals (Fe, Co, Ni) • Lanthanide and Actinide Series(inner transition elements): filling 4fand 5forbitals (Eu, Am, Es)

7. Information Contained in the Periodic Table • Each group member has the same valence electron configuration (these electrons primarily determine an atom’s chemistry). • The electron configuration of any representative element.

8. Information Contained in the Periodic Table • Certain groups have special names (alkali metals, alkaline earth metals, chalcogens, halogens, etc). • Metals and nonmetals are characterized by their chemical and physical properties.

9. Special Names for Groups in the Periodic Table

10. Metals • Metals makeup more than 75% of the elements in the periodic table. Metals are characterized by the following physical properties: • They have metallic shine or luster. • They are usually solids at room temperature. • They are malleable. Malleable means that metals can be hammered, pounded, or pressed into different shapes without breaking. • They are ductile meaning that they can be drawn into thin sheets or wires without breaking. • They are good conductors of heat and electricity.

11. Metals (cont) • All B and most A elements are metals. • The B  At stairstep designates the border between metals and non-metals • 1A elements are alkali metals • They are soft shiny metals that usually combine with group VIIA nonmetals in chemical compounds in a 1:1 ratio. • 2A elements are the alkaline earth metals • Both alkali and alkaline earth metals are chemically reactive, but 2A metals are less reactive than 1As. • They combine with the group VIIA nonmetals in a 1:2 ratio.

12. Transition Metals & Metalloids • Transition metals • The remaining 1-8B elements are all transition elements • The transition elements also have valence electrons in two shells instead of one. • Inner transition metals • The lanthanide and actinide series comprise the inner transition metals

13. Metalloids • Metalloids have characteristics of both metals and nonmetals and so can’t be classified as either, but something in between. • They are good conductors of heat and electricity • They are not good conductors or insulators. • The six metalloids are B, Si, Ge, As, Sb, and Te.

14. Nonmetals • There are 17 nonmetals in the periodic table, and they are characterized by four major physical properties. • They rarely have metallic luster. • They are usually gases at room temperature. • Nonmetallic solids are neither malleable nor ductile. • They are poor conductors of heat and electricity. • The elements above the B  At stairstep are nonmetals

15. Nonmetals (cont) • Group 6A contains the chalcogen elements • Group 7A contains the highly reactive halogen elements • They are fluorine, chlorine, bromine, and iodine. • The halogens exist as diatomic molecules in nature. • Group 8A comprises the completely non-reactive noble gases • The noble gases are also called rare gas elements, and they all occur in nature as gases. • The noble gases fulfill the octet rule by having a full outer level with 8 valence electrons. • Therefore, they do not undergo chemical reactions because they do not accept any electrons.

16. Valence Electrons and the Periodic Table • Valence Electrons and Group • Atoms in the same group have the same chemical properties because they have the same number of valence electrons. • Moreover, they have the same outermost orbital structure • E.g. 1A elements all have s1 valence electrons • E.g. 2A elements all have s2 valence electrons • Valence Electrons and Period • The primary quantum number (n) for an element’s valence electrons is the same its period. • E.g. Lithium’s valence electron is n=2 and Li is found in the 2nd period

17. The Octet Rule • Atoms tend to lose, gain, or share electrons until they are surrounded by 8 valence electrons

18. Exceptions to Aufbau Order • Subshell degeneracies occur in elements larger than Vanadium • i.e. different 4s and 3d orbitals have nearly the same energy • Also, it turns out that full and half-full sublevels have the most stability.

19. Exceptions: Copper (Cu), Silver (Ag), Gold (Au) • Strict Aufbau ordering of Cu would be [Ar]4s23d9 • experimental observation shows this to be an excited state • the ground state has a configuration of [Ar]4s13d10 • The observed configuration for Cu creates a ½-full s and a full d, which is more stable than a full s and a partial d • Ag is NOT [Kr]5s24d9, but [Kr]5s14d10 • Au is NOT [Xe]6s25d9, but [Kr]6s15d10

20. Exceptions: Lanthanum and Actinium • Aufbau would place them in the inner transition series, but instead they are in the scandium family • i.e. La is [Xe]6s25d1 • i.e. Ac is [Rn]7s26d1

21. Exceptions: Chromium (Cr), Molybdenum (Mo), but NOT Tungsten (W) • Cr is [Ar]4s13d5, NOT [Ar]4s23d4 • Mo is [Kr]5s14d5, NOT [Kr]5s24d4 • W IS [Xe]6s15d5

22. Ionization Energy • The quantity of energy required to remove an electron from the gaseous atom or ion.

23. For Aluminum • Al (g) Al+ (g) + e- I1 = 580 kJ/mol • Al + (g) Al2+ (g) + e- I2 = 1850 kJ/mol • Al2+ (g) Al3+ (g) + e- I3 = 2740 kJ/mol • Al3+ (g) Al4+ (g) + e- I4 = 11,600 kJ/mol

24. Periodic Trends • First ionization energy: • increasesfrom left to right across a period Why? • decreases going down a group. Why?

25. Trends in Ionization Energies for the Representative Elements

26. The Values of First Ionization Energy for the Elements in the First Six Periods

27. Question • The first ionization energy for the group IIA elements are significantly higher than those of the Group IA elements in the same periods. Why?

28. Question • The first ionization energy of the Group IIIA elements are lower than the IIA elements in the same period. Why?

29. Question • Group VIA elements have slightly lower first ionization energies than Group VA elements in the same period. Why?

30. Electron Affinity • The energy change associated with the addition of an electron to a gaseous atom. • X(g) + e X(g) • Note: the more negative the electron affinity, the more energy is released

31. The Electronic Affinity Values for Atoms Among the First 20 Elements that Form Stable, Isolated X- Ions

32. Questions • Helium and Beryllium do not form stable isolated negative ions. Why? • Nitrogen does not form a stable, isolated N- (g) ion, whereas carbon forms C-(g). Why? • In contrast to nitrogen, oxygen can add an electron to form the stable O- ion. Why?

33. Periodic Trends • Atomic Radii: • decrease going from left to right across a period; Why? • increasegoing down a group. Why?

34. The Radius of an Atom

35. Atomic Radii for Selected Atoms

36. Ionic Radii • What is the trend for ionic radii? • Which of the Period 3 ions would be the smallest? • Na+, Mg2+, Al3+, S2-, Cl-

37. Sizes of Ions Related to Positions of the Elements in the Periodic Table

38. Electronegativity Increases Up and To the Right

39. What is electronegativity? • How tightly an atom holds on to its valence electrons. • Essentially, this value depends on • the number of positively charged protons in the atom’s nucleus • the radius of the outermost electron shell

40. What is electronegativity? • The more positive the nucleus • The smaller the valence electron shell around it • The greater the attraction between nucleus and electrons • Thus, the more electronegative the atom! • Thus, a high electronegativity value implies that the valence electrons are tightly held and require a large amount of energy to remove.

41. Oxidation Numbers • The Octet Rule states that atoms want to have their valence shell filled with electrons. • This means that, ideally, atoms are most stable with 8 valence electrons • N.B. This is not true for Period 1. Why? • Atoms will gain or lose electrons to form ions in order to fulfill the Octet Rule. • The charge they take on in this process is called the valence. • The oxidation state is, for ions, equal to the valence.