1. 8 - 1 Main Group Elements Atomic radius is defined as being one half the distance between identical nuclei bonded in a molecule. Atoms get smaller as you proceed from left to right across a period (series).  The nucleus contains more protons and the electron cloud contains more electrons.  The increased charge results in a greater attraction making the atom smaller.

2. 8 - 2 Atomic Radius Across A Period. Na Mg Al Si P S Cl Ar Atomic Radius Decreases

3. 8 - 3 Atomic Radius

4. 8 - 4 Main Group Elements Atoms get larger as you proceed from top to bottom down a group (family).  There is one more principal energy level each time you go down a period.  The valence electrons get further from the nucleus and feel less of an attraction by the positive nucleus.

5. 8 - 5 Atomic Radius Down A Group. Li Na K Rb Cs Atomic Radius Increases

6. 8 - 6 Atomic Radius

7. 8 - 7 Ionic Radius for Cations in a Period Cations are smaller than the atoms from which there are formed. Metals give off their valence electrons, i.e. Ca → Ca 2+ + 2e - 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 → 1s 2 2s 2 2p 6 3s 2 3p 6 + 2e - The entire highest numbered principal energy is lost (n = 4) which also decreases the number of electrons which decreases the repulsion.

8. 8 - 8 Ionic Radius Across A Period. Na MgAlSiP S Cl Ar Atomic radius decreases for both cations and anions.

9. 8 - 9 Ionic Radii for Anions in a Period Anions are larger than the atoms from which there are formed. Nonmetals take in valence electrons to form a complete octet, i.e. N + 3e - → N 3- 1s 2 2s 2 2p 3 + 3e - → 1s 2 2s 2 2p 6 Adding electrons increases the repulsion between electrons residing in the same sublevel.

10. 8 - 10 Ionic Radii for Anions in a Period Compare nitrogen to oxygen: N + 3e - → N 3- 1s 2 2s 2 2p 3 + 3e - → 1s 2 2s 2 2p 6 O + 2e - → O 2- 1s 2 2s 2 2p 4 + 2e - → 1s 2 2s 2 2p 6 N 3- and O 2- are isoelectronic but O 2- has the greater nuclear charge making it the smaller anion.

11. 8 - 11 Ionic Radius Across A Period. Na MgAlSiP S Cl Ar Atomic radius decreases for both cations and anions.

12. 8 - 12 Transition Elements. Cr 2+ Cr 3+ Mn 2+ Fe 2+ Fe 3+ Ni 2+ Cu + Cu 2+ Co 2+ Co 3+ Transition elements tend to have multiple valence numbers. Almost all the transition elements of the fourth period form monatomic ions with a charge of +2.

13. 8 - 13 Transition Elements Because transition elements in the d-block have 2 electrons in their valence shell, they tend to react chemically the same. Both lanthanoids and actinoids are found in the f-block. Lanthanoids occur in trace amounts in nature and are called the rare earth elements. Actinoids usually have large unstable nuclei that undergo spontaneous radioactive decay.

14. 8 - 14 Ionization Energy First Ionization Energy Ionization energy or ionization potential is the minimum amount of energy needed to remove an electron from the valence shell of a gaseous atom. Na(g) + IE1 Na + (g) + e - IE indicates how easy it is for a metal to form a cation.

15. 8 - 15 Ionization Energy Generally, IE increases across a period (series) because of an increase in nuclear charge. IE increases as the size of the atom decreases. Nonmetals easily accept electrons causing them to have a high IE.

16. 8 - 16 Ionization Energy As you move from left to right in a period, the IE also depends on half-filled and completely-filled orbitals.  When an s sublevel is filled with 2 electrons, there is an increase in its stability.  When a p sublevel is half-filled with 3 electrons, there is an increase in its stability.

17. 8 - 17 Ionization Energy  When a p sublevel is completely filled with 6 electrons, there is an even greater increase in its stability.  The same is true for a d sublevel except it is for 5 electrons and 10 electrons.  Remember the two exceptions in Period 4, Cr (3d 5 ) and Cu (3d 10 ) which follows this same tendency.

18. 8 - 18 Ionization Energy The highest IE occurs for the noble gases because they have a complete octet. Generally, IE decreases from the top to the bottom in a group or family because of the addition of a principal energy level.  IE decreases as the size of the atom increases.

19. 8 - 19 Ionization Energy The second IE is larger than the first because the second electron is being removed from a cation rather than a neutral atom. The third IE is larger than the second because the third electron is being removed from a cation with a +2 charge.

20. 8 - 20 First Ionization Energy He Ne Ar Kr Xe Rn

21. 8 - 21 Electronegativity Electronegativity is a measure of the attraction of an element for a shared pair of electrons. H Cl Comparing the electronegativity values of hydrogen and chlorine, chlorine has a value of 3.2 and that of hydrogen is 2.2... δ-δ- δ+δ+

22. 8 - 22 Electronegativity The origin of these values is unimportant and the atom with the higher value is more electronegative. The most electronegative element is 9 F because it has the smallest atomic radius with very few of it electrons shielding the nucleus. The least electronegative is 87 Fr.

23. 8 - 23 Electronegativity Generally, electronegativity increases from left to right within a period or series. Generally, electronegativity decreases from top to bottom within a group or family. Electronegativity values are not assigned to the noble gases because they are inactive. The explanation for the trends in electronegavity is the same as for ionization energy.

24. 8 - 24 Electronegativity F Cl Br I At

25. 8 - 25 Electron Affinity A measure of an atom’s tendency to gain electrons in the gas phase. A(g) + e - A - (g) + thermal energy Electron affinity is an irregular periodic function of atomic number. In general, it increases from left to right. Noble gases are not included since they have little or no tendency to gain electrons.

26. 8 - 26 Periodic Trends in Density Generally, the density increases as you proceed from top to bottom in a group of metals or nonmetals.  The atomic mass increases more rapidly than the atomic radius.

27. 8 - 27 Density (g/cm 3 ). ElementDensityElementDensity Li0.53F1.31 Na0.97 Cl1.56 K0.86 Br3.12 Rb1.53 I4.92 Cs1.90

28. 8 - 28 Periodic Trends in Density Generally, the density increases as you proceed from left to right in a period until you reach the metalloids.  There is a big drop off in density in the nonmetals (gases) but then starts to increase.

29. 8 - 29 Trends in Boiling and Melting Points The boiling and melting points generally decrease as you proceed from top to bottom in the metals.  This results from metallic bonding. The boiling and melting points generally increase as you proceed from top to bottom in the nonmetals.  This results from Van der Waals forces.

30. 8 - 30 Boiling and Melting Points. ElementBP (°C)MP (°C) Li1372179 Na 892 98 K 774 64 Rb 679 39 Cs 690 28 Alkali Alkali Metals Metallic Bonds

31. 8 - 31 Boiling and Melting Points. Element BP (°C) MP (°C) F-187-223 Cl -35-101 Br 59 -72 I 185 114 Halogens Van der Waal Forces

32. 8 - 32 Metallic Characteristics Metals are good conductors of heat and electricity due to their “sea of electrons”. Metals have shiny surfaces that are both malleable and ductile due to their d-electrons. Metals are malleable because they can be hammered into a thin foil without breaking. Metals are ductile because they can be stretched into a thin wire without breaking.

33. 8 - 33 Metallic Characteristics Metals have three or fewer valence electrons which they donate during chemical reactions. Generally, metallic character increases as you go down a group or family.  The valence electrons are further from the nucleus and are more shielded from the nucleus.

34. 8 - 34 Metallic Characteristics Generally, metallic character decreases as you proceed from left to right in a period (series).  Metals → Metalloids → Nonmetals

35. 8 - 35 Metallic Characteristics When a metal donates electrons it is said to undergo oxidation. Metals that are more easily oxidized will react more readily in the presence of a nonmetal. Group I and Group II are very active metals.  Because they have such a strong tendency to form compounds they are not found in their elemental or free state.

36. 8 - 36 Alkali Metals The Group IA metals have an outer electron configuration of ns 1. The loss of an electron to form a 1+ ion is the basis of almost all reactions of the alkali metals. M → M + + e -

37. 8 - 37 Alkaline Earth Metals The Group IIA metals have an outer electron configuration of ns 2. The Group II metals are not as reactive as the alkali metals because they need to lose two electrons from a completely filled s-sublevel in order to achieve a noble gas configuration. M → M 2+ + 2e -

38. 8 - 38 Nonmetallic Characteristics Nonmetals are poor conductors of heat and electricity. Nonmetals have dull surfaces and are brittle. Nonmetals have 4-7 valence electrons, therefore they gain electrons. Generally, nonmetallic character decreases from top to bottom within a group or family.

39. 8 - 39 Nonmetallic Characteristics The atomic radius is a very important factor in determining the reactivity in nonmetals. F and Cl are the smallest halogens and will more readily accept electrons in their valence shell. According to Coulomb’s Law, the positive nucleus will attract valence electrons more in a smaller atom.

40. 8 - 40 Nonmetallic Characteristics Generally, nonmetallic character increases as you proceed from left to right in a period (series).  Metals → Metalloids → Nonmetals

41. 8 - 41 Nonmetallic Characteristics When a nonmetal accepts electrons it is said to undergo oxidation. Nonmetals that are more easily oxidized will react more readily in the presence of a nonmetal. Group VII are very reactive nonmetals.  Because they have such a strong tendency to form compounds they are not found in their elemental or free state.

42. 8 - 42 Halogens The common group VIIA elements are all nonmetals. Each only needs a single electron to achieve a noble gas configuration. When reacting with metals, they form 1- ions. 2Na(s) + Cl 2 (g) 2NaCl(s) When they have no other elements to react with, they are found as diatomic molecules. 2F(g) F 2 (g)

43. 8 - 43 Noble Gases Each noble gas has filled s and p sublevels except for helium (1s 2 ). All are very unreactive. A limited number of compounds have been produced using xenon and krypton. Xe(g) + F 2 (g) XeF 2 (g)

44. 8 - 44 The Anomoly of Hydrogen Hydrogen is a nonmetallic gas at room temperature. While it may lose an electron to form H +, it also can gain an electron to form H - (hydride). 2Na(l) + H 2 (g)2NaH(s) Hydrogen is placed in Group IA. Where else could it go?

45. 8 - 45 Semimetals (Metalloids) These elements (B, Si, Ge, As, Sb, Te, and At) along the “stairway” exhibit properties of both metals and nonmetals. They have some similarities with metals because they are shiny and conduct electricity. They are similar to nonmetals because they are brittle.

46. 8 - 46 Allotropes Allotropes are elements having more than one form because of structural differences (the way in which their atoms or molecules are arranged). The element oxygen has 3 forms:  O – monatomic oxygen  O 2 – molecular or diatomic oxygen  O 3 - ozone