1. Periodic Trends

2. Atomic Size First problem where do you start measuring.
The electron cloud doesn’t have a definite edge.
They 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.

4. Trends in Atomic Size Influenced by two factors. 1) Energy Level
Higher energy level is further away.
2) Charge on nucleus (#protons)
More charge pulls electrons in closer.

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

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

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

8. 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 is called the first ionization energy.

9. 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 & 2nd IE.

10. Symbol First Second Third
HHeLiBeBCNO F Ne

11. Symbol First Second Third
HHeLiBeBCNO F Ne

12. What determines IE The greater the nuclear charge the greater IE.
The greater the distance from nucleus, the lower IE
Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE
Shielding - lowers the IE

13. Shielding
This outermost electron is shielded from the attractive nuclear forces by the inner electrons
When a new orbital is started, every orbital of lower energy shields these electrons from feeling the full nuclear charge.

14. Group trends
As you go down a group first IE decreases because the electron is further away (more energy levels) and there is more shielding.

15. Periodic Trends
All the atoms in the same period have the same energy level and same shielding, BUT there is increasing nuclear charge SO IE generally increases from left to right.
IE decreases to make full and half filled orbitals.

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

17. First Ionization energyHe
Li has lower IE than H  more shielding & further away which outweighs greater nuclear charge H
First Ionization energy Li
Atomic number

18. Be has higher IE than Li  same shielding BUT greater nuclear charge
First Ionization energy H Be Li
Atomic number

19. B has lower IE than Be  same shielding, greater nuclear chargeHe
B has lower IE than Be  same shielding, greater nuclear charge
BUT the 2p electron in B is easier to remove than the 2s electron in Be
First Ionization energy H Be B Li
Atomic number

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

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

22. First Ionization energyHe
Oxygen breaks the pattern because removing an electron gets to 1/2 filled p orbital (1s22s22p3) N
First Ionization energy H C O Be B Li
Atomic number

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

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

25. 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

26. First Ionization energy
Atomic number

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

28. 2nd Ionization Energy
For elements that reach a filled or half filled orbital by removing 2 electrons, the 2nd IE is lower than expected.
True for s2
Alkali earth metals form +2 ions.

29. 3rd IE Using the same logic s2p1 atoms have a low 3rd IE.
Atoms in the aluminum family form + 3 ions.
2nd IE and 3rd IE are always higher than 1st IE!!!

30. X(g) + e- → X-(g) + energy 
Electron Affinity
The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions.
X(g) + e- → X-(g) + energy 
In other words, the neutral atom's likelihood of gaining an electron.

31. Electron Affinity
Easiest to add to group 7A….gets them to full energy level.
EA -Increase from left to right B/C atoms become smaller, with greater nuclear charge.
EA-Decrease as we go down a group.

32. Ionic Size: Cations 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.

33. Configuration of Ions 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.

34. Ionic size: Anions Anions form by gaining electrons
Anions are bigger that the atom they come from.
Non-metals form anions
Anions of representative elements have configuration of noble gas after them.

36. Sizes of Ions: Explanation
1) Cations are smaller than their parent ions.
Electrons have been removed from the most spatially extended orbital.
The effective nuclear charge has increased.
Therefore, the cation is smaller than the parent.
2) Anions are larger than their parent ions.
Electrons have been added to the most spatially extended orbital. This means total e--e- repulsion has increased.
The nuclear charge has remained the same, but the number of screening electrons has increased.
Therefore, anions are larger than their parents.

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

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

39. Size of Isoelectronic ions
Isoelectronic ions have the same # of electrons (same electron config)
Eg. Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
all have 10 electrons
all have the configuration 1s22s22p6

40. Size of Isoelectronic ions
Positive ions have more protons so they are smaller.
N-3
O-2
F-1 Ne
Na+1
Al+3
Mg+2

41. O2- > F- > Na+ > Mg2+ > Al3+
Sizes of Ions: Summary
For ions of the same charge, ion size increases down a group.
All the members of an isoelectronic series have the same number of electrons, BUT as nuclear charge increases in an isoelectronic series the ions become smaller:
O2- > F- > Na+ > Mg2+ > Al3+

42. Electronegativity
A chemical property that describes the ability of an atom to attract electrons towards itself in a covalent bond.

43. Electronegativity
The tendency for an atom to attract electrons in a bond to itself
How fairly it shares electrons.
Big electronegativity means it pulls the electron toward it.

44. Group Trend
The further down a group the farther the electron is away and the more electrons an atom has.
More willing to share.
Low electronegativity.

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

46. Ionization energy, electronegativity
Electron affinity INCREASE

47. Atomic size increases,
Ionic size increases - see p Heath