1. Periodic Trends Chapter 6

2. Octet Rule Atoms tend to achieve electron configuration of Noble Gases Octet = Eight Noble Gases have eight electrons in their highest energy level General Equation for Noble Gases is S 2 P 6

3. IONS Ion- is an atom or group of atoms that have a positive or negative charge A typical atom is electrically neutral because it has an equal amount of protons and electrons Positive and Negative Ions are formed when at atom donates or receives an electron An Ion with a positive charge is called a Cation An Ion with a negative charge is called an Anion

4. Effective Nuclear Charge Force of attraction between an electron and the nucleus depends on the magnitude of the net nuclear charge acting on the electron and the average distance between the nucleus and the electron. Force of attraction increases as the nuclear charge increases and decreases as the electron moves farther from the nucleus.

5. Effective Nuclear Charge Cont. Valence electron in an atom is attracted to the nucleus of the atom and is repelled by the other electrons. Inner electrons (core) partially shield the outer electrons from the attraction of the nucleus

6. The effective nuclear charge increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller. Periodic Properties: Effective Nuclear Charge Mg has a greater effective nuclear charge than Na, and is smaller than Na.

7. Trends in Atomic Size Atomic Radius- ½ the distance between the nuclei of two like atoms in a diatomic molecule Radius is measured in Picometers 1pm = 1 picometer = 1 x 10 -12 m

8. Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

9. Group Trends of Atomic Size Atomic Size generally increases as you move down a group on the periodic table As you descend, electrons are added to higher principle energy levels and the nuclear charge increases The outermost orbital is also larger as you move down a group The shielding of the nucleus by electrons also increases as you move down a group

10. Group trends As we go down a group Each atom has another energy level, So the atoms get bigger. H Li Na K Rb

11. Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus

12. Shielding The electron on the outside energy level has to look through all the other energy levels to see the nucleus. A second electron has the same shielding.

13. Increasing Atomic Size

14. Periodic Trends Atomic Size generally decreases as you move from left to right across a period As you go across a period, the energy level remains Each element has one more proton and electron then the preceding The electrons are added to the same principle energy level

15. The effect of the increasing nuclear charge on the outermost electrons is to pull them closer to the nucleus Atomic Size therefore decreases

16. Periodic Trends As you go across a period the radius gets smaller. Same energy level. More nuclear charge. Outermost electrons are closer. NaMgAlSiPSClAr

17. Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

19. Trends in Ionization Energy When an atom gains or loses an electron, it becomes an ion Ionization Energy- The energy required to overcome the attraction of the nuclear charge and remove an electron from an atom The energy required to remove the first outermost electron is called the first ionization energy The energy required to remove the second outermost electron is called the 2 nd ionization energy Ect……

20. Ionization Energy The second ionization energy is the energy required to remove the second electron. Always greater than first IE. The third IE is the energy required to remove a third electron. Greater than 1st of 2nd IE.

21. The Noble Gases are at the top showing they don’t want to form an Ion The Alkali are at the bottom of the peaks, showing their ease to form an Ion

22. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11 810 14840 3569 4619 4577 5301 6045 6276

26. What determines IE The greater the nuclear charge the greater IE. Distance form nucleus increases IE Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. Shielding

27. Group Trends 1 st Ionization energy generally decreases as you move down a group The size of the atoms increases as you descend, so the outermost electron is farther from the nucleus The outermost electron should be more easily removed and the element should have a lower ionization energy

28. Periodic Trends The 1 st ionization energy generally increases as you move from left to right across a period The nuclear charge increases and the shielding effect is constant as you move across A greater attraction of the nucleus for the electron leads to the increase in ionization energy Exceptions at Full and ½ fill orbitals

29. First Ionization energy Atomic number He He has a greater IE than H. same shielding greater nuclear charge H

30. First Ionization energy Atomic number H He l Li has lower IE than H l more shielding l further away l outweighs greater nuclear charge Li

31. First Ionization energy Atomic number H He l Be has higher IE than Li l same shielding l greater nuclear charge Li Be

32. First Ionization energy Atomic number H He l B has lower IE than Be l same shielding l greater nuclear charge l By removing an electron we make s orbital half filled Li Be B

33. First Ionization energy Atomic number H He Li Be B C

34. First Ionization energy Atomic number H He Li Be B C N

35. First Ionization energy Atomic number H He Li Be B C N O Breaks the pattern because removing an electron gets to 1/2 filled p orbital

36. First Ionization energy Atomic number H He Li Be B C N O F

37. First Ionization energy Atomic number H He Li Be B C N O F Ne Ne has a lower IE than He Both are full, Ne has more shielding Greater distance

38. First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

39. First Ionization energy Atomic number

40. Driving Force Full Energy Levels are very low energy. Noble Gases have full orbitals. Atoms behave in ways to achieve noble gas configuration.

41. 2nd Ionization Energy For elements that reach a filled or half filled orbital by removing 2 electrons 2nd IE is lower than expected. True for s 2 Alkali earth metals form +2 ions.

42. 3rd IE Using the same logic s 2 p 1 atoms have an low 3rd IE. Atoms in the aluminum family form + 3 ions. 2nd IE and 3rd IE are always higher than 1st IE!!!

43. Electron Affinity The energy change associated with adding an electron to a gaseous atom. Easiest to add to group 7A. Gets them to full energy level. Increase from left to right atoms become smaller, with greater nuclear charge. Decrease as we go down a group.

44. Electron Affinity The greater the attraction between a given atom and an added electron, the more negative the atom’s electron affinity The more negative the E.A., the greater the attraction of the atom for the electron The trends in E.A. are not very evident.

46. Difference between I.E. & E.A. Ionization Energy measures the ease with which an atom loses an electron Electron Affinity measures the ease with which an atom gains an electron

47. Trends in Ionic Size Atoms of metallic elements have low ionization energies. They form positive ions easily Atoms of nonmetallic elements readily form negative ions. How does the lose or gain of electrons affect the size of the ion formed?

48. Group Trends Positive Ions are always smaller than the neutral atoms from which they form. The loss of outer-shell electrons results in increased attraction by the nucleus for the fewer remaining electrons Negative Ions are always larger than the neutral atoms from which they form The effective nuclear attraction is less for an increased number of electrons

49. Group trends Adding energy level Ions get bigger as you go down. Li +1 Na +1 K +1 Rb +1 Cs +1

50. The Sodium Atom is larger than the Sodium Cation. Why is this true?

51. The Chlorine Atom is smaller then the Chlorine Anion. Why is this true?

52. Sodium Cation is smaller than the Sodium Atom Chlorine Anion is larger than the Chlorine Atom

55. Periodic Trends Going from left to right across a row, there is a gradual decrease in the size of the positive ions. Beginning with group 5A, the negative ions, which are much larger, gradually decrease in size an you continue to move right.

56. Periodic Trends Across the period nuclear charge increases so they get smaller. Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

59. Trends in Electronegativity Electronegativity- is the tendency for the atoms of the element to attract electrons when chemically combined with atoms of another element. Electronegativities have been calculated for elements and are expressed in arbitrary units on the Pauling electronegativity scale The scale is based on a number of factors

61. Group Trends Electronegativity generally decreases as you move down a group The metallic elements have a low electronegativity meaning they don’t want to want attract electrons

63. Periodic Trends As you go across a period from left to right, the electronegativity of representative elements increases The non-metallic elements (excluding Noble Gases) have high electronegativities The trends in electronegativities among transitional metals are not so regular

64. Electronegativity values help predict the type of ionic or covalent bonding that can exist between atoms in compounds