1. Periodic Trends: Periodic Properties
Certain physical and chemical properties recur at regular intervals, and/or vary in regular fashion, when the elements are arranged according to increasing atomic number.
Melting point, boiling point, hardness, density, physical state, and chemical reactivity are periodic properties.
We will examine several periodic properties that are readily explained using electron configurations.

2. Periodic Trends Atomic Size – Group – increase in size down the group
(addition of energy levels)
Period – decrease in size across a period
(increase in the # protons in the nucleus pulling on the same energy level)

3. Periodic Properties: Atomic Radius
Half the distance between the nuclei of two atoms is the atomic radius.
Covalent radius: half the distance between the nuclei of two identical atoms joined in a molecule.
Metallic radius: half the distance between the nuclei of adjacent atoms in a solid metal.

4. Periodic Properties: Atomic Radius
Atomic radius increases from top to bottom within a group.
The value of n increases, moving down the periodic table.
The value of n relates to the distance of an electron from the nucleus.

5. Periodic Properties: Atomic Radius
Atomic radius decreases from left to right within a period.
Why? The effective nuclear charge increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.
Mg has a greater effective nuclear charge than Na, and is smaller than Na.

6. Multielectron and Shielding
Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.
Zeff = Zactual - Electron shielding

7. Atomic Radii of the Elements

8. Example
With reference only to a periodic table, arrange each set of elements in order of increasing atomic radius:
Na, Al, P
(b) Te, S, Se
(c) Ge, Sb, Sn

9. Ionic Radii
The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion.

10. Ionic Radii of Cations
Cations are smaller than the atoms from which they are formed; the value of n usually decreases. Also, there is less electron–electron repulsion.

11. Ionic Radii of Anions
Anions are larger than the atoms from which they are formed.
Effective nuclear charge is unchanged, but additional electron(s) increase electron–electron repulsion.
Isoelectronic species have the same electron configuration; size decreases with effective nuclear charge.

12. Atomic Radii and the Periodic Table

13. Example
Use the standard periodic table to arrange the following species in the expected order of increasing radius:
Sr2+, Ru3+, Rb+, I–, Br–
Isoelectronic atoms= Sr+2, Rb+1, Br-1
Atom
Protons
Electrons
Configuration
Rank
Sr+2 38 36
[Kr]5s2 
[Kr] 1
Ru+3 44 41
[Kr]5s24d6  [Kr]4d5 4
Rb+1 37
[Kr]5s1  2
I-1 53 54
[Kr]5s24d105p5  [Xe] 5
Br-1 35
[Ar]4s23d104p5 [Kr] 3

14. Periodic Trend Ionization Energy – Group – decreases down the group
(addition of energy levels makes it harder for the nucleus to hold on and easier to remove electrons)
Period – increases across the period
(addition of protons pulling on the electrons makes it harder to remove the electrons)

15. Ionization Energy
Ionization energy (I) is the energy required to remove an electron from a ground-state gaseous atom.
I is usually expressed in kJ per mole of atoms.
M(g)  M+(g) + e– ΔH = I1
M+(g)  M2+(g) + e– ΔH = I2
M2+(g)  M3+(g) + e– ΔH = I3

16. Ionization Energy Trends
I1 < I2 < I3
Removing an electron from a positive ion is more difficult than removing it from a neutral atom.
A large jump in I occurs after valence electrons are completely removed (why?).
I1 decreases from top to bottom on the periodic table.
n increases; valence electron is farther from nucleus.
I1 generally increases from left to right, with exceptions.
Greater effective nuclear charge from left to right holds electrons more tightly.

17. Selected Ionization Energies
Compare I2 to I1 for a 2A element, then for the corresponding 1A element.
Why is I2 for each 1A element so much greater than I1?
Why don’t we see the same trend for each 2A element? I2 > I1 … but only about twice as great …

18. Example
Why is lithium’s I2 14 times as much as the I1 while beryllium is only twice as much?
Li [He]2s1
Li [He] or 1s2
Li s1
Be [He]2s2
Be+1 [He]2s1
Be+2 [He]

19. Selected Ionization Energies
General trend in I1: An increase from left to right, but …
The electron being removed is now a p electron (higher energy, easier to remove than an s).
…I1 drops, moving from 2A to 3A.
Repulsion of the paired electron in 6A makes that electron easier to remove.
I1 drops again between 5A and 6A.

20. Example
Why is boron's I1 less than that of beryllium? B Be

21. Example
Why is oxygen’s I1 less than that of nitrogen? O N

22. First Ionization Energies

23. Second Ionization Energies
The amount of energy required to remove the second electron once the first is gone.
Always greater than THEIR first IE due to an INCREASED effective nuclear charge (a greater proton to electron ratio), making the atom smaller.
However, 2nd Ionization energies can be compared BETWEEN different atoms by evaluating the +1 ion of THAT atom.
If the +1 ion is a noble gas the 2nd IE will be LARGE. If the +1 ion is NOT a noble gas the 2nd IE will be SMALLER.

24. Second Ionization Energies
Example:
1st IE
2nd IE Na
[Ne]3s2
Small
(wants to lose 1e-)
Big
(Doesn’t want to lose 2e-) Mg
Bigger than Na (doesn’t want to lose any)
Smaller than Na
(once 1e- is gone, it is easy to lose 2e-)

25. Periodic Trends: Periodic Properties
Certain __________________________recur at regular intervals, and/or vary in regular fashion, when the elements are arranged according to ______________________.
Melting point, boiling point, hardness, density, physical state, and chemical reactivity are periodic properties.
We will examine several periodic properties that are readily explained using electron configurations.

26. Periodic Trends
Atomic Size –
Group –
Period –

27. Periodic Properties: Atomic Radius
Half the distance between the nuclei of two atoms is the _____________________.
_____________________: half the distance between the nuclei of two identical atoms joined in a molecule.
_____________________: half the distance between the nuclei of adjacent atoms in a solid metal.

28. Periodic Properties: Atomic Radius
Atomic radius __________ from top to bottom within a group.
The value of n increases, moving down the periodic table.
The value of n relates to the __________ of an electron from the nucleus.

29. Periodic Properties: Atomic Radius
Atomic radius __________ from left to right within a period.
Why? The ________________________increases from left to right, increasing the attraction of the nucleus for the valence electrons, and making the atom smaller.
Mg has a greater effective nuclear charge than Na, and is smaller than Na.

30. Multielectron and Shielding
Effective Nuclear Charge (Zeff): The nuclear charge actually felt by an electron.
Zeff = Zactual - Electron shielding

31. Atomic Radii of the Elements

32. Example
With reference only to a periodic table, arrange each set of elements in order of increasing atomic radius:
Na, Al, P
(b) Te, S, Se
(c) Ge, Sb, Sn

33. Ionic Radii
The ___________ of each ion is the portion of the distance between the nuclei occupied by that ion.

34. Ionic Radii of Cations
__________ are _________ than the atoms from which they are formed; the value of n usually decreases. Also, there is less electron–electron repulsion.

35. Ionic Radii of Anions
________ are _______than the atoms from which they are formed.
Effective nuclear charge is unchanged, but additional electron(s) increase electron–electron repulsion.
_____________ species have the same electron configuration; size decreases with effective nuclear charge.

36. Atomic Radii and the Periodic Table

37. Example
Use the standard periodic table to arrange the following species in the expected order of increasing radius:
Sr2+, Ru3+, Rb+, I–, Br–

38. Periodic Trend
Ionization Energy –
Group –
Period –

39. Ionization Energy
____________________(I) is the energy required to remove an electron from a ground-state gaseous atom.
I is usually expressed in kJ per mole of atoms.
M(g)  M+(g) + e– ΔH = I1
M+(g)  M2+(g) + e– ΔH = I2
M2+(g)  M3+(g) + e– ΔH = I3

40. Ionization Energy Trends
I1 < I2 < I3
Removing an electron from a __________ is more difficult than removing it from a ____________.
A large jump in I occurs after valence electrons are completely removed (why?).
I1 ____________ from top to bottom on the periodic table.
n increases; valence electron is farther from nucleus.
I1 generally _________ from left to right, with exceptions.
Greater effective nuclear charge from left to right holds electrons more tightly.

41. Selected Ionization Energies
Compare I2 to I1 for a 2A element, then for the corresponding 1A element.
Why is I2 for each 1A element so much greater than I1?
Why don’t we see the same trend for each 2A element? I2 > I1 … but only about twice as great …

42. Example
Why is lithium’s I2 14 times as much as the I1 while beryllium is only twice as much?

43. Selected Ionization Energies
General trend in I1: An increase from left to right, but …
The electron being removed is now a p electron (higher energy, easier to remove than an s).
…I1 drops, moving from 2A to 3A.
Repulsion of the paired electron in 6A makes that electron easier to remove.
I1 drops again between 5A and 6A.

44. Example
Why is boron's I1 less than that of beryllium?

45. Example
Why is oxygen’s I1 less than that of nitrogen?

46. First Ionization Energies

47. Second Ionization Energies
The amount of energy required to remove the second electron once the first is gone.
Always greater than _____ first IE due to an ____________ effective nuclear charge (a greater proton to electron ratio), making the atom smaller.
However, 2nd Ionization energies can be compared ________ different atoms by evaluating the +1 ion of ____ atom.
If the +1 ion is a noble gas the 2nd IE will be _____. If the +1 ion is NOT a noble gas the 2nd IE will be _________.

48. Second Ionization Energies
Example:
1st IE
2nd IE Na
Small (wants to lose 1e-)
Big (Doesn’t want to lose 2e-) Mg
Bigger than Na (doesn’t want to lose any)
Smaller than Na (once 1e- is gone, it is easy to lose 2e-)