THE PERIODIC TABLE

1. THE PERIODIC TABLE Objectives: Examine the progression of periodicity

2. Alkali metals alkaline earth metals “s” group Transition metals “d” block Inner transition metals “f” group Metalloids (semimetals) Nonmetals “p” block Halogens Noble gases

3. Periodic Patterns • ALKALI METALS (part of the “s” group of elements) ~ all are in shiny solid form but are quite soft ~ form the 1st group of metals on the periodic table ~ highly reactive elements based upon their electron configurations (ns1) Other Characteristics: • malleable and ductile; low density and melting points • good conductors of electricity; very soluble as comps.

4. Alkaline Earth Metals PART OF THE “S” GROUP JUST LIKE ALKALI METALS • Belong to the second group of metals on the periodic table. • Harder, more dense, and stronger than there group 1 counterparts • Not as reactive as the Alkali metals due to the metals in this group having 2 electrons in their valence shell. • This gives them the configuration of “ns2” for these metals. Be Mg Ca Sr Ba Ra

5. Transition Metals The “d” - block elements • Transition metals begin in the 4th period after the alkaline earth metals. • Metallic elements with varying properties. • Not nearly as reactive as group 1 and 2 elements. • Fill their sublevels differently than do the Main group elements. Important in living organisms! Valuable as structurally useful materials!

6. LANTHANOIDS Composed of the elements with atomic numbers 58 through 71 Electrons are being added to the “ 4f “ sublevel Shiny reactive metals with practical uses ie. dots in TV tubes ACTINOIDS Composed of the elements with the atomic numbers 90 through 103 They fill the “ 5f “ sublevel All are radioactive with an unstable nucleus Lanthanoids & Actinoids The ‘f’-group is broken into two classifications “Y”

7. Nonmetals & Metalloids (semi-metals) NONMETALS! • Generally are gases at room temperature (or brittle solids) • Poor conductors of heat and electricity • Have more electrons in their outer level than metals METALLOIDS! • Properties of both metals and nonmetals • Will give up (electron donor) electron(s) when reacted with a nonmetal, and will accept (electron acceptor) electron(s) when reacted with a metal • In general, more like nonmetals than metals • Considered semiconductors

8. Periodic Trends • Trends (we will study) – atomic radius (ionic radius), ionization energy, electronegativity, electron affinity • Trends are looked at from top to bottom of a column and from left to right in a period (row) • Trends show patterns of atoms properties (relationships among elements)

9. Trend Number 1 Li Na K Rb Cs Fr ATOMIC RADIUS • Atomic radius is the half the distance between the nuclei of two like atoms. TREND: • the trend for atomic radius shows us the size of the atom will increase as we move down a column WHY: more levels and orbitals, greater distances from the nucleus

10. ATOMIC RADIUS TREND: • the atomic radii will decrease from the left to the right in a period WHY: Effective Nuclear Charge (also applies to what takes place from top to bottom of a column) • positive charge felt by the outermost electrons of an atom • atomic # - # of inner complete level electrons • The larger the ENC, the greater the attraction of electrons to the nucleus Shielding - the ability of other electrons,especially inner electrons, to lessen the nuclear charge of the outer electron(s)

11. THE TREND! • The trend shows the increase of radii down a group and decrease of radii across a period.

12. NUMBER TWO! IONIZATION ENERGY Magnesium IONIC BOND • bond formed between two ions by the transfer of electrons Ions: How do they form? • In certain types of bonding, the atom will “lose” or “gain” an electron(s) • When an atom loses or gains electrons, it is called an ion

13. Magnesium • Atoms that lose electrons have a positive charge • Atoms that gain electrons have a negative charge BOINK! BOINK!

14. CHLORINE • For the most part, the metals will lose electrons and the nonmetals will accept the electrons • The atoms gain or lose electrons to reach outer shell (valence) stability Electron from magnesium

16. Ionic Bonds: One Big Greedy Thief Dog!

17. Ion Sizes Does the size go up or down when losing an electron to form a cation?

18. + + Li , 78 pm 2e and 3 p Ion Sizes Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p

19. Ion Sizes Does the size go up or down when gaining an electron to form an anion?

20. - - F, 71 pm F , 133 pm 9e and 9p 10 e and 9 p Ion Sizes Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The electron/proton attraction has gone DOWN and so size INCREASES.

21. Trends in Ion Sizes Figure 8.13

22. IONIZATION ENERGY • the energy required to remove the most loosely held electron from an atom • ionization energy decreases as the size of the atom increases (top to bottom of a column) “Y”? Because the outer most electron is farther from the nucleus and the electrical attraction to the protons.

23. More Details! • Energy is absorbed by the atom to free the electron(s) • Ionization is endothermic, meaning that the atom or molecule increases its internal energy ( takes energy from an outside source) A + energy A+ + e-

24. Ionization Energy is affected by three factors: 1. Effective Nuclear Charge 2. Number of Energy Levels 3. Shielding

25. Ionization Energies • The first ionization energy, I1, is the energy needed to remove the first electron from the atom: • Mg  Mg+ + 1e-

26. The second ionization energy, I2, is the energy needed to remove the next (i.e. the second) electron from the atom • Mg+ Mg2+ + 1e- • The higherthe value of the ionization energy, the more difficultit is to remove the electron

27. Ionization Energies in kJ/mol

28. Within each period ( row) the ionization energy increases with atomic number. • Y? • Electrons are being added to the same energy level (ENC) • increasing valence electrons as approaching the nonmetals

29. The Trend

30. Electronegativity • The tendency for an atom to attract electrons to itself when in combination with another atom • Defined differences in electronegativity determine the bonding character of a compound • Ionic or Covalent bonds Linus Pauling scale is used to determine electronegativity differences

31. COVALENT BOND bond formed by the sharing of electron clouds • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs

32. Covalent Bonds

33. NONPOLAR COVALENT BONDS when electron clouds are shared equally H2 or Cl2

34. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O2)

35. POLAR COVALENT BONDS when electron clouds are shared but shared unequally H2O

36. Polar Covalent Bonds: Unevenly matched, but willing to share.

37. - water is a polarmolecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

39. Electronegativity Differences and Bond Type 0 • If the electronegativity difference is less than 0.2 then the bond is a nonpolar covalent • If the difference is between 0.2 and 1.6, the bond is polar covalent • If the difference is greater than 2, the bond is ionic nonpolar covalent 0.2 polar covalent 1.6 ? 2 ?  between 1.6 and 2, if a metal is involved, the bond is ionic. If only nonmetals are involved the bond is polar covalent ionic 4

40. Trend of EN decrease increase

41. Electron Affinity elements GAIN electrons to form anions. Electron affinity is the energy change when an electron is added: A(g) + e- ---> A-(g) E.A. = ∆E

42.      -  [He] O ion + electron       [He] O atom Electron Affinity of Oxygen ∆E is EXOthermic because O has an affinity for an e-. EA = - 141 kJ

43. Trends in Electron Affinity • Affinity for electron increases across a period (EA becomes more negative). • Affinity decreases down a group (EA becomes less negative). Atom EA F -328 kJ Cl -349 kJ Br -325 kJ I -295 kJ

44. Trends in Electron Affinity

45. Practice with Comparing Ionization Energies • For each of the following sets of atoms, decide which has the highest and lowest ionization energies and why. • a. Mg, Si, S • b. Mg, Ca, Ba • c. F, Cl, Br • d. Ba, Cu, Ne • e. Si, P, N

46. Answers to Comparing Ionization Energies • Here are answers to the exercises above. • a. Mg, Si, S • All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear charge. • b. Mg, Ca, Ba • All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of energy levels.

47. c. F, Cl, Br All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of energy levels. d. Ba, Cu, Ne All are in different groups and periods, so both factors must be considered. Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear charge and uses the lowest number of energy levels.

48. e. Si, P, N Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for having the highest effective nuclear charge.

49. BECAUSE... The relative stability of an atom can be predicted by its electron configuration

50. Rule of Thumb • As a general rule, elements with three or fewer electrons in their outer level are considered to be metals.