Chapter 5 – The Periodic Law

1. Chapter 5 – The Periodic Law Chuck Norris destroyed the periodic table, because he only recognizes the element of surprise.

2. Some people are REALLY into the elements. geek bling geek tattoo

3. The Periodic Table…so far • Displays the known elements in order of increasing atomic number • Arrangement is related to electron configurations of elements • Horizontal rows are periods (energy levels) • Vertical columns are groups or families (similar properties)

4. The periodic table also… • Allows us to predict the properties of elements. • Helps us to understand the chemical behavior of elements. • Allows us to detect trends in the properties of the elements. • Shows us that chemical properties are predictable.

5. 5-1 History of the Periodic Table • Dmitri Mendeleev, Russian chemist – hoped to organize the known elements according to their properties • At the time (1800s), nothing was known about the structure of the nucleus. • Mendeleev put elements with their properties on cards, arranged them according to their properties, and looked for patterns

6. 5-1 Mendeleev and Chemical Periodicity • Noticed that when elements placed in order of increasing atomic mass, certain similarities appeared at regular intervals • Repeating pattern is called PERIODIC • Published the first periodic table in 1869

7. 5-1 Mendeleev’s Periodic Table • Left empty spaces in order to keep elements with similar properties together • Suggested empty spaces represented elements not yet discovered AND predicted the properties of three undiscovered elements • by 1886, all three had been discovered – predictions were very accurate • Other chemists accept his arrangement.

8. 5-1 Mendeleev’s Periodic Table – A Problem • A few elements are out of atomic mass order • This was necessary to keep elements with similar properties together • Mendeleev thought atomic masses may have been incorrectly measured

9. 5-1 Moseley and Atomic Number • 1911, English physicist Henry Moseley was working with atomic spectra of metals • Noticed elements fit into patterns better when arranged in order of increasing NUCLEAR CHARGE • Led to modern definition of ATOMIC NUMBER and new basis for arrangement of periodic table

10. 5-1 The Modern Periodic Table • An arrangement of elements in order of their ATOMIC NUMBERS so that elements with similar properties fall in the same column or group • Properties repeat PERIODICALLY

11. 5-1 The Modern Periodic Table An interactive periodic table Even better interactive periodic table

12. 5-1 The Modern Periodic Table • Noble Gases – earliest tables omitted them because they hadn’t been discovered yet (1894 – Ar, 1895 – He, 1898 – Kr and Xe, 1900 – Rn) • Lanthanides and Actinides – added in early 1900s when their properties were better understood – belong in periods 6 and 7, between s block and p block

13. 5-2 Electron Configuration and the Periodic Table • The properties of an element are determined by the electron configuration of the atom’s highest occupied energy level (the valence electrons) • A filled outer energy level (s & p sublevels filled) gives an atom special stability – this is why the noble gases are so unreactive.

14. 5-1 The Periodic Table and Electron Configuration

15. 5-2 Using the Periodic Table to Determine Electron Configuration Ending • For s and p blocks: period gives energy level, block gives sublevel, and column gives number or electrons • Examples: K, Si, Rn, Cl, Sb, Sr, Mg, P

16. 5-2 The s-Block Elements: Groups 1 and 2 • Chemically REACTIVE metals (group 1 > group 2) • Group 1: ns1, alkali metals, silvery, soft, can be cut with a knife, so reactive they are not found free in nature (only in compounds) • Group 2: ns2, alkaline earth metals, harder, denser and stronger than alkali metals, higher melting points, less reactive than alkali metals but still too reactive to be found free in nature

17. 5-2 The Alkali Metals • Alkali metal reactivity

18. 5-2 Alkaline Earth Metals

19. 5-2 Hydrogen and Helium • Special cases • Hydrogen: 1s1, but is NOT an alkali metal – it is a nonmetal and a gas at room temperature • Helium: 1s2, placed with noble gases because of its properties (unreactive gas), has a filled outer energy level with only 2 electrons (1st energy level only holds 2 electrons)

20. 5-2 The d-Block Elements: Groups 3-12 • For energy level n, there are n possible sublevels, so n=3 has 3 sublevels • Because 3d has higher energy than 4s, the d-block first appears in the 4th period • ns(n-1)d • Transition metals – metals with typical metallic properties, good conductors of heat and electricity, high luster, less reactive than s-block metals, some so unreactive they exist in nature as free elements, palladium, platinum and gold are least reactive

21. 5-2 Exceptions to the Aufbau Principle • Atoms with a filled or half-filled sublevel have special stability • Some d-block elements have unusual electron configurations that reflect this fact • Cr, Cu

22. 5-2 The p-Block Elements – Groups 13-18 • Electrons only add to p sublevel once s is filled, so all p-block elements have 2 electrons in ns. • Along with s-block elements, p-block elements are called MAIN-GROUP ELEMENTS. • Number of valence electrons = group # - 10 • Properties vary greatly – metals (8), metalloids (6) and nonmetals (including halogens and noble gases)

23. 5-2 Noble Gases • ns2[(n-1)d10]np6 • Have an octet (a filled s and p sublevel in their highest energy level) • Have 8 valence electrons, filled outer energy level • Helium is an exception – has a filled outer energy level with only 2 electrons • Stable and inert (unreactive)

24. 5-2 Halogens • Group 17: fluorine, chlorine, bromine, iodine, astatine) • “Salt forming” • Most reactive nonmetals • F, Cl – gases at RT, Br – liquid at RT, I – solid at RT

25. 5-2 Metalloids • Semiconducting elements • Mostly brittle solids with some metallic properties and some nonmetallic properties • Intermediate conductivity • Top – As, bottom - Si

26. 5-2 p-Block Metals • Harder and denser than s-block metals but softer and less dense than d-block metals • Reactive – only found in nature in compounds (except Bi) • Once obtained as free metals, stable in air • Top – Pb, bottom – Bi

27. 5-3 Electron Configuration and Periodic Properties • Atomic Radius – distance from the center of the nucleus to the outer edge of the electron cloud • OR ½ the distance between the nuclei of two identical atoms that are bonded together

28. 5-3 Trends in Atomic Radii • L to R across a period – atomic radius decreases due to increasing nuclear charge • Elements in a period have outermost electrons in same energy level so might predict about same size, but increasing charge of nucleus pulls electron cloud in

29. 5-3 Trends in Atomic Radii • Top to bottom down a group – atomic radius increases • Electrons occupy higher energy levels located farther from nucleus

30. 5-3 Atomic Radii

31. 5-3 Atomic Radius as a Function of Atomic Number

32. 5-3 Valence Electrons and the Periodic Table

33. 5-3 Ions and Ion Formation • Ion – an atom or group of bonded atoms that has a positive or negative charge • Monatomic ions form when an atom gains or loses an electron or electrons • Cations – positive ions formed by the loss of electrons • Anions – negative ions formed by the gain of electrons

34. 5-3 The Octet Rule and Ion Formation • Octet rule – atoms will gain, lose or share electrons in order to obtain a filled outer energy level (usually 8 electrons) • Metals tend to lose electrons in order to gain an octet • Nonmetals tend to gain (or share) electrons in order to gain an octet

35. 5-3 The Octet Rule: Sodium • Sodium has one valence electron. • It loses one valence electron to reveal a filled shell. • A sodium ion has a charge of +1, Na+.

36. 5-3 The Octet Rule: Chlorine • Chlorine has 7 valence electrons. • It is very close to having a filled outer energy level. • It will gain one electron to fill its outer energy level. • A chloride ion has a charge of -1, Cl-.

37. 5-3 The Octet Rule and Ion Formation • In general… • Atoms with 1, 2 or 3 valence electrons (metals) will lose their valence electrons, forming cations. • Atoms with 5, 6 or 7 valence electrons (nonmetals) will gain enough electrons to complete their octet, forming anions.

38. 5-3 Ionization Energy • Neutral atoms can lose electrons • Ionization energy, IE – energy required to remove one electron from a neutral atom of an element, measured in kJ/mol A + energy  A+ + e-

39. 5-3 Trends in Ionization Energy • L to R across a period – ionization energy increases - it becomes harder to remove the valence electron(s) • The closer the valence electrons are to the nucleus and the greater the nuclear charge, the harder they are to remove – held more tightly • Nonmetals have higher ionization energies than metals

40. 5-3 Trends in Ionization Energy • Top to bottom down group – ionization energy decreases – it becomes easier to remove the valence electron(s) • Electrons removed from each successive element are farther from the nucleus so are held less tightly • More electrons lie between the nucleus and the valence electron(s), which partially shields the valence electron(s) from the pull of the nucleus

41. 5-3 Trends in Ionization Energy • Where on the periodic table are the elements most likely to LOSE electron?

42. 5-3 Multiple Ionization Energies • Electrons can be removed from positive ions as well as neutral atoms • These successive IEs are names 2nd ionization energy, 3rd ionization energy, etc • 2nd ionization energy is always higher than 1st • 3rd ionization energy is always higher than 2nd • As electrons are removed, fewer electrons remain to shield attractive force of nucleus, each successive electron removed feels greater effective nuclear charge

43. 5-3 Successive Ionization Energies of Selected Elements in kJ/mol

44. 5-3 Ionization Energy • Why do you think helium has the highest first ionization energy of any element?

45. 5-3 Electron Affinity • Neutral atoms can gain electrons. • electron affinity – the energy change that occurs when an electron is acquired by a neutral atom A + e- A- + energy The quantity of energy released is represented by a negative number The more energy lost, the more likely an atom is to gain an electron A value of zero means the atom has NO tendency to gain an electron

46. 5-3 Trends in Electron Affinity • Halogens gain electrons most readily • L to R across p block, electron affinities get more negative • Exception is between groups 14 and 15 because of stability of half-filled sublevel • Adding an electron to carbon is easier than adding an electron to nitrogen, which requires forcing an electron to pair up in an orbital of the already half-filled p sublevel • Note noble gases have electron affinities of zero.

47. 5-3 Trends in Electron Affinity • Top to bottom down group – generally decreases down a group, but there are some exceptions • Generally, as atomic radius increases, attraction of an atom for an electron decreases

48. 5-3 Ionization Energy, Electron Affinity and Reactivity

49. 5-3 Ionic Radii • Cations – smaller than their respective atoms – outer energy level is lost and remaining electrons are drawn closer to nucleus

50. 5-3 Ionic Radii • Anions – larger than their respective atoms – electron cloud spreads out because of greater repulsion between increased number of electrons