1. Periodic Trends OBJECTIVES:
Interpret periodic trends in atomic radii, ionic radii, ionization energies, and electronegativities.

2. Trends in Atomic Size
First problem: Where do you start measuring from?
The electron cloud doesn’t have a definite edge.
We get around this by measuring more than 1 atom at a time.

3. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule. Diatomic means “two atoms” or a two atom molecule

4. Trends in Atomic Size Influenced by three factors: 1. Energy Level
Higher energy level is further away.
2. Charge on Nucleus
More charge pulls electrons in closer.
3. Shielding Effect
Electrons between nucleus and valance electrons

5. Group Trends As we go down a group...H
As we go down a group...
each atom adds another energy level,
so the atoms get bigger. Li Na K Rb

6. Periodic Trends As you go across a period, the radius gets smaller.
Electrons are in same energy level.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

7. Atomic Radius Trends

8. Atomic radii 76

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

10. Periodic Trend in Atomic Radius77

11. Ion Formation
An atom can easily lose or gain electrons. The resulting ion is an atom that has an imbalance of charge or carries either a positive charge (a loss of electrons) or a negative charge (a gain of electrons).
A positive ion is called a cation.
A negative ion is called an anion.

12. Trends in Ionization Energy
The amount of energy required to completely remove an electron from a gaseous atom.
Removing one electron makes a 1+ ion.
The energy required to remove the first electron is called the first ionization energy.

13. 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 or 2nd IE.

14. Symbol First Second Third
HHeLiBeBCNO F Ne

15. Symbol First Second Third
HHeLiBeBCNO F Ne

16. Ionization Energy
How many valance electrons does Magnesium have? How many is it willing to give up? 85

17. First Ionization Energy86

18. Second Ionization Energy87

19. Third ionization Energy
Mg3+
1s22s22p5 87

20. What Determines IE The greater the nuclear charge, the greater IE.
Greater distance from nucleus decreases IE
Filled and half-filled sublevel have lower energy, so achieving them is easier, lower IE.
Shielding effect on valence electrons

21. Shielding
The electron on the outermost energy level (valence electron) has to look through all the other inner electrons to see the nucleus.
Shielding electrons hinder the nucleus form holding on to the valence electrons

22. Group Trends As you go down a group, first IE decreases because...
the valence electron is further away
and there is more shielding.

23. Periodic Trends As you go across a period,
All the atoms in the same period have the same highest energy level.
They have similar shielding.
But, increasing nuclear charge reduces atomic radius (bringing valence closer)
so IE generally increases from left to right.
Exceptions are at full and 1/2 full sublevels.

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

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

26. greater nuclear chargeHe
Be has higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

27. greater nuclear charge He
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron we leave the s orbital full
First Ionization energy H Be B Li
Atomic number

28. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

29. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

30. First Ionization energyHe N
Breaks the pattern, because removing an electron leaves 1/2 filled p orbital
First Ionization energy H C O Be B Li
Atomic number

31. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

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

33. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

34. First Ionization energy
Atomic number

35. Driving Force
Full Energy Levels require lots of energy to remove their electrons.
Noble Gases have full energy level.
Atoms behave in ways to achieve noble gas configuration.

36. Periodic trend in first Ionization Energy88

37. 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 s2
Alkaline earth metals form 2+ ions.

38. 3rd Ionization Energy
Using the same logic s2p1 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!!!

39. Trends in Electron Affinity
The atoms ability to acquire an additional electron
Cl + 1e Cl1-
The energy change associated with adding an electron to a gaseous atom.
Easiest to add to group 7A.
Gets them to a full energy level.
Increase from left to right: atoms become smaller, with greater nuclear charge.
Decrease as we go down a group.

40. Trends in Ionic Size Cations form by losing electrons.
Cations are smaller than the atom they come from.
Metals form cations.
Cations of representative elements have noble gas configuration. Na
Na+1

41. Ionic Size Anions form by gaining electrons.
Anions are bigger than the atom they come from.
Nonmetals form anions.
Anions of representative elements have noble gas configuration. Cl
Cl-1

42. Configuration of Ions Ions always have noble gas configuration.
Na is: 1s22s22p63s1
Forms a 1+ ion: 1s22s22p6
Same configuration as neon.
Metals form ions with the configuration of the noble gas before them - they lose electrons.

43. Configuration of Ions
Non-metals form ions by gaining electrons to achieve noble gas configuration.
They end up with the configuration of the noble gas after them.

44. Ion Orbital Notation 90

45. Ion Group Trends Li1+ Na1+ K1+ Rb1+ Cs1+ Adding energy level
Ions get bigger as you go down.
Li1+
Na1+
K1+
Rb1+
Cs1+

46. Ion Periodic Trends
Across the period, nuclear charge increases so they get smaller.
Energy level changes between anions and cations.
N3-
O2-
F1-
B3+
Li1+
C4+
Be2+

47. Ionic Radius Trends

48. Size of Isoelectronic Ions
Iso- means the same
Iso electronic ions have the same # of electrons
Al3+ Mg2+ Na1+ Ne F1- O2- and N3-
all have 10 electrons
all have the configuration: 1s22s22p6

49. Size of Isoelectronic Ions
Positive ions that have more protons would be smaller.
N3-
O2-
F1- Ne
Na1+
Al3+
Mg2+

50. Electronegativity is a measure of the ability of an
atom in a molecule to attract
electrons to itself.
Concept proposed by
Linus Pauling

51. Electronegativity
The ability of an atom that is bonded to another atom or atoms to attract electrons to itself.
It is related to ionization energy and electron affinity.
It cannot be directly measured.
The values are unitless since they are relative to each other.
The values vary slightly from compound to compound but still provide useful qualitative predictions.

52. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.
How fair is the sharing?
Large electronegativity means the atom pulls the electrons toward it.
Atoms with large electron affinity have larger electronegativity.

53. Group Trend
The further down a group, the farther the valence electrons are away from the nucleus and the more electrons an atom has shielding.
The larger the atom, the more willing it is to share electrons.
Low electronegativity.

54. Electronegativities Electronegativity is a periodic property.
Atomic number

55. Periodic Trend Metals are at the left of the table.
They let their electrons go easily
Low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away from others
High electronegativity.

56. Electronegativity
Relative ability of atoms to attract electrons of bond. At
1.9 I
2.2 Br
2.7 Cl
2.8 Po
1.8 Te
2.0 Se
2.5 S
2.4 Bi
1.7 Sb As P
2.1 Pb
1.5 Sn Ge Si F
4.1 O
3.5 N
3.1 Tl
1.4 Na
1.0 Cs
0.9 Rb K Ba Mg
1.2 Sr Ca In Ga Al H Li Be B C

57. Electronegativity

58. Ionization Energy, Electronegativity, and Electron Affinity INCREASE

59. Atomic Size Increases
Ionic size increases