1. 5.3 Electron Configuration & Periodic Properties
Chapter 5
The Periodic Law
5.3 Electron Configuration & Periodic Properties

2. Chapter 5 Notes The Periodic Law
F. Atomic Radius
G. Shielding Effect and Nuclear Charge
H. Ionization Energy
I. Ion formation and Ionic Radii
J. Electronegativity

3. Section 3 Electron Configuration and Periodic Properties
Chapter 5
Objectives
Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.
Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.

4. Determination of Atomic Radius:
Half of the distance between nuclei in
covalently bonded diatomic molecule
"covalent atomic radii"
F. Periodic Trends in Atomic Radius:
Period - size decreases across a period
due to increased nuclear charge
Group - size increases down a group
due to increased shielding effect

5. ALL PT Trends Influenced by 3 factors:
1. Shielding Effect & Energy Level
higher energy levels further from nucleus
outer electrons shielded from the nucleus
2. Nuclear Charge (# protons)
More + charge pulls e-’s in closer

6. Shielding
e-’s in outer energy level “looks through” other energy levels to see nucleus

7. G. Shielding effect and Nuclear Charge
increase in the number of occupied orbitals shields electrons in the highest occupied energy level (shielding effect) from the attraction of protons in the nucleus (nuclear charge)
shielding effect is greater than the effect of the increase in protons (nuclear charge)
shielding effect is reason atomic size increases when drop to next row (period) down a group
nuclear charge is reason atomic size decreases when move across a period

8. Shielding effect and Nuclear Charge
What do they influence?
Shielding Effect & Energy Level
have effect on GROUP ( )
Nuclear Charge has effect on PERIOD ( )

9. #1. Atomic Size - Group trends
Going down a group, each atom gains another energy level (floor)
atoms get….. b H i Li g Na g K e r Rb

10. #2. Atomic Size - Period Trends
L to R across period:
size gets….
e-’s occupy same energy level
more nuclear charge
Outer e-’s pulled closer Al Si P S Ar Cl Na Mg m S a l l e r

11. Table of Atomic Radii

12. How to Achieve an Octet…
Atoms can form ions by gaining or losing electrons to obtain a stable outer configuration
Cation- Positive ion (+) ion
Anion- Negative ion (-) ion
Ions attract (opposites attract)

13. Main Ideas:
removing electrons from atoms to form ions requires energy… ionization energy
adding electrons to atoms to form ions also requires energy … electron affinity

14. Predicting Ionization
Metals tend to lose electrons
They form cations.
Ex: Na, 1s22s22p63s1 becomes Na+1,1s22s22p6
Nonmetals tend to gain electrons.
They form anions.
Ex: O, 1s22s22p4 becomes O-2, 1s22s22p6

15. Trends in Ionization Energy
Ionization energy - energy required to completely remove e- (from gaseous atom)
energy required to remove only 1st e-called first ionization energy.

16. Ionization Energy
second ionization energy is E required to remove 2nd e-
Always greater than first IE.
third greater than 1st or 2nd IE.

17. Ionization Energy - Group trends
going down group
first IE decreases b/c...
e- further away from nucleus attraction
more shielding

18. Ionization Energy - Period trends
Atoms in same period:
same energy level
Same shielding
Increasing nuclear charge
So IE generally increases left - right
Exceptions…full & 1/2 full orbitals

19. Ionization Energy Symbol First Second Third
HHeLiBeBCNO F Ne
Why did these values increase so much?

20. Both w/ same shielding (e- in 1st level)He
He greater IE than H.
Both w/ same shielding (e- in 1st level)
He = greater nuclear charge H
First Ionization energy
Atomic number

21. These outweigh greater nuclear charge
Li lower IE than H
more shielding
further away
These outweigh greater nuclear charge H
First Ionization energy Li
Atomic number

22. same shielding (period) greater nuclear chargeHe
Be higher IE than Li
same shielding (period)
greater nuclear charge
First Ionization energy H Be Li
Atomic number

23. greater nuclear charge By removing e- s orbital half-filledHe
B has lower IE than Be
same shielding
greater nuclear charge
By removing e- s orbital half-filled
First Ionization energy H Be B Li
Atomic number

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

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

26. He
Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital N
First Ionization energy H C O Be B Li
Atomic number

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

28. Ne has a lower IE than He Ne more shielding First Ionization energy
Both full but…
Ne farther
Ne more shielding
b/c greater distance F N
First Ionization energy H C O Be B Li
Atomic number

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

30. First Ionization energy
Atomic number

31. Driving Forces Full Energy Levels require high E to remove e-
Noble Gases = full orbitals
Atoms want noble gas configuration

32. 2nd Ionization Energy
elements w/ filled or ½ filled orbital by removing 2 e-, 2nd IE lower than expected.
True for s2
Alkaline earth metals form 2+ ions.

33. 3rd IE
Using same logic s2p1 atoms have an low 3rd IE.
Atoms in Al family form 3+ ions
2nd IE and 3rd IE are always higher than 1st IE!!!

34. H. Ionization Energy - the energy required to
H. Ionization Energy - the energy required to remove an electron from an atom
Increases for successive electrons taken from
the same atom
Period – IE increases across a period
Electrons in the same quantum level do not shield as effectively as electrons in inner levels
Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove
Group - IE decrease down a group
Outer electrons are farther from the nucleus

35. Ionization of Magnesium
Mg kJ  Mg+ + e-
Mg kJ  Mg e-
Mg kJ  Mg e-

36. Table of 1st Ionization Energies

37. Another Way to Look at Ionization Energy

38. Electron Affinity - the energy change associated
Electron Affinity - the energy change associated with the addition of an electron
Affinity tends to across a period
increase
Affinity tends to down a group
decrease
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals

39. Table of Electron Affinities

40. Electron Transfer: Anions
When an atom gains electrons it increases its negative charge so it becomes negatively charged. There are now more electrons than protons.
X + e- = X –
Ex: Nitrogen Atom + 7 protons
- 7 electrons
Nitrogen Ion +7 protons
- 10 electrons
Neutral
-3 charge

41. Electron Transfer: Cations
When an atom loses electrons, it loses negative charges so it becomes more positively charged. There are now more protons than electrons.
X - (e-) = X +
Ex: Potassium Atom protons
- 19 electrons
Potassium Ion +19 protons
-18 electrons
Neutral
+1 charge

42. Ionic Radii A positive ion is known as a cation
The formation of a cation by the loss of one or more electrons always leads to a in atomic radius.
cation
decrease
The electron cloud becomes smaller.
The remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.
A negative ion is known as an anion.
The formation of an anion by the addition of one or more electrons always leads to an in atomic radius.
increase

43. Trends in Ionic Size: Cations
Cations lose e-’s
metals
Cations smaller than atom they came from
lose e-’s
lose entire energy level.
Cations of representative elements have noble gas configuration before them

44. Trends in Ionic size: Anions
Anions gain e-‘s
nonmetals
Anions bigger than atom they came from
same energy level
greater area nuclear charge needs to cover

45. Configuration of Ions
Ions always have noble gas configurations (full outer level)
Na atom is: 1s22s22p63s1
Forms a 1+ sodium ion: 1s22s22p6
Same as Ne

46. Ionic Radii Cationic and anionic radii decrease across a period.
The electron cloud shrinks due to the increasing nuclear charge acting on the electrons in the same main energy level.
The outer electrons in both cations and anions are in higher energy levels as one reads down a group.
There is a gradual increase of ionic radii down a group.

47. Configuration of Ions
Non-metals form ions by gaining e-’s to achieve noble gas configuration
configuration of noble gas after them

48. Ion Group trends Each step down a group adds energy level
Li1+
Na1+
Each step down a group adds energy level
Ions - bigger going down
b/c extra energy level
K1+
Rb1+
Cs1+

49. Ion Period Trends Across period
nuclear charge increases
Ions get smaller
energy level changes btwn anions & cations
N3-
O2-
F1-
B3+
Li1+
Be2+
C4+

50. I. Ionic Radii 2. Anions 1. Cations positively charged ions
usually metal
smaller than the corresponding atom
2. Anions
negatively charged ions
usually nonmetal
larger than the corresponding atom

51. Table of Ion Sizes

52. Periodic Trends of Radii
Section 3 Electron Configuration and Periodic Properties
Chapter 5
Periodic Trends of Radii

53. Summation of Periodic Trends

54. OBJECTIVE:
Define valence electrons , and state how many are present in atoms of each main-group element.

55. Valence electrons
Chemical compounds form because electrons are lost, gained, or shared between atoms.
The electrons that interact in this manner are those in the highest energy levels.
The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons.
Valence electrons are often located in incompletely filled main-energy levels.
example: the electron lost from the 3s sublevel of Na to form Na+ is a valence electron.

56. Valence electrons are…?
The electrons responsible for the chemical properties of atoms, and are those in the outer energy level.
Valence electrons - The s and p electrons in the outer energy level
the highest occupied energy level
Core electrons – are those in the energy levels below.

57. valence electrons– electron in the highest occupied energy level of an atom.
*To find the number of valence e- in an atom of a representative element, simply look at its group number
s-block = group number
ex. Li is in group 1= 1 valence e-
p-block = group number – 10
ex. F is in group = 7 valence e-

58. Keeping Track of Electrons
Atoms in the same column...
Have the same outer electron configuration.
Have the same valence electrons.
The number of valence electrons are easily determined. It is the group number for a representative element
Group 2: Be, Mg, Ca, etc.
have 2 valence electrons

59. Electronegativity
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to one atom than to another.
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
Electronegativities tend to increase across periods, and decrease or remain about the same down a group.

60. Trends in Electronegativity
Electronegativity (EN)- tendency for atom to attract e-’s when atom in cmpd
Sharing e-, but how equally?
Element w/ big EN pulls e- towards itself strongly!

61. Electronegativity Group Trend
Further down group, farther e- away from nucleus
plus more e-’s atom has
more willing to share
Low EN

62. Electronegativity Period Trend
Metals let e-’s go easily
low EN
Nonmetals want more e-’s
take e-’s from others
High EN

63. J. Electronegativity
A measure of the ability of an atom in a chemical compound to attract electrons
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same

64. Periodic Table of Electronegativities

65. Trends of Atomic Radius

66. Trends in Ionization Energy (IE)

67. Trends in Electronegativity (EN)

68. Section 3 Electron Configuration and Periodic Properties
Chapter 5
Objectives
Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.
Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.
Define valence electrons, and state how many are present in atoms of each main-group element.
Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.

69. 5.3 Practice Problems Page 135 Sets E-G, Math Tutor
*answers on page R120

70. End of Chapter 5 Show