1. Periodic Relationships Among the Elements
Chapter 8
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2. When the Elements Were Discovered

3. Development of the Periodic Table
Dimitri Mendeleev (1834 – 1907)
Arranged the elements by increasing ATOMIC MASS, he saw a periodic repetition of properties
Produced the first PERIODIC TABLE – 1871
The table placed elements with similar properties in the same column
Kept “holes” for undiscovered elements, and predicted the properties in advance

4. Development of the Periodic Table
Properties of elements predicted by Mendeleev

5. Development of the Periodic Table
H. G. Moseley ( )
Rearranged the elements by ATOMIC NUMBER
This has become the MODERN PERIODIC TABLE

6. Valence Electrons
valence electrons: electrons available to be lost, gained, or shared in the formation of chemical compounds
electrons in the outermost energy level
electrons that are responsible for reactions
Elements in a group have similar properties because they have the same valence electron configuration

7. All electrons under the highest energy level
Inner Core Electrons
All electrons under the highest energy level

8. Valence Electron Configuration
Group
e- configuration
Valence electrons 1
ns1 2
ns2 13
ns2np1 14
ns2np2 15
ns2np3 16
ns2np4 17
ns2np5 18
ns2np6

9. Charges Of Representative Elements+1 +2 +3 -3 -2 -1
8.2

10. What atoms are isoelectronic with Ne?
an ion that has the same electron configuration
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]

11. Electron Configurations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5

12. Periodic Table Groups

13. patterns on the periodic table are called periodic trends
PERIODIC LAW
When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic (repeating) pattern.
patterns on the periodic table are called periodic trends

14. half the distance from center-center of 2 like atoms
Atomic Radius
half the distance from center-center of 2 like atoms

15. Atomic Radii DOWN a Group
As you go down a group another energy level is added, the atom size gets larger
The number of occupied orbitals between the nucleus and the outermost energy level increases
Shielding Effect:
reduction of attraction between positive nucleus and outer electrons, outer electrons are not held tight and can move away

16. Atomic Radius: down groupX Na X X X X P X P X P P X P P P P P X P X P X X

17. Atomic Radius: down groupX K X X X X X X X P X X P X P P X X P P P P P X X P X P X X X X

18. Atomic Radii DOWN a Group
DOWN THE GROUP ATOMIC RADIUS
INCREASES
more energy levels,
the larger the size of the atom

19. Atomic Radii ACROSS a Period
Each atom gains one proton and one electron in the same energy level
Each added electron is the same distance from the nucleus
The positive charge increases and exerts a greater force on the electrons thereby pulling it closer to the nucleus

20. Effective nuclear charge (Zeff) :
“positive charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Shielding: effect of electrons in the same energy level
Zeff  Z – number of inner or core electrons
~Zeff
Core Z
Radius Na Mg Al Si 11 12 13 14 10 1 2 3 4
186
160
143
132
2s1 would have shielding from 1s2 but 1s2 does not have shielding form 2s1
Outer energy levels have shielding form inner, but its because of the repulsion of inner core electrons with the shielding constant
Shielding constant is very small and alters exact charge
* For our purposes, close enough to number of valence electrons
Within a Period
as Zeff increases
radius decreases

21. P e PROTONS REMEMBER! - + are bigger and stronger! electrons
are smaller
and weaker!

22. Atomic Radius: across periodX X X X X P X P P X P P P P P X P X P X X

23. Atomic Radii ACROSS a Period
ACROSS THE PERIOD ATOMIC RADIUS
DECREASES
stronger attraction of protons,
easier to hold on to the electrons

25. Atomic Radii

26. half the distance from center-center of 2 like ions
Ionic Radii
half the distance from center-center of 2 like ions

27. Ionic Radius ACROSS the Period
Cation: positive ion formed from losing an electron
A cation is always smaller than the original atom
The more electrons lost the more protons available to attract a smaller number of electrons.

28. Ionic Radius P X
Na + X X P P X P P P P P P P P P

29. Ionic Radius ACROSS the Period
ACROSS THE PERIOD IONIC RADIUS
DECREASES
greater pull on electrons,
the shorter the radius

30. Ionic Radius ACROSS the Period
Anion: negative ion formed from gaining an electron
A anion is always larger than the original atom
The more electrons gained, the less protons available to attract a larger number of electrons.

31. Ionic Radius P X F - X X X X P X P P P X P P P P P X P P X X

32. Ionic Radius ACROSS the Period
ACROSS THE PERIOD IONIC RADIUS
DECREASES
more electrons added,
more difficult to keep track of,
increasing the size of the atom

33. Ionic Radius DOWN a Group
As you go down a group another energy level is added, increasing the size of the atom.
(just like the atomic radius)

34. Ionic Radius DOWN the Group
DOWN THE GROUP IONIC RADIUS
INCREASES
more energy levels,
increase in atom size

35. 8.3

37. Ionization Energy
minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state
amount of energy needed to remove an electron from an atom
Ion: an atom that has gained or lost electrons
if you lose an electron…
if you gain an electron…
BECOMES POSITIVE!
BECOMES NEGATIVE!

38. Multiple Ionization Energies
I1 + X (g) X+(g) + e-
I1 first ionization energy
I2 + X (g) X2+(g) + e-
I2 second ionization energy
I3 + X (g) X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3

39. Ionization Energy DOWN a Group
As you go down a group atoms become larger, electrons are farther from the nucleus and more easily removed
The more electrons in an atom between the nucleus and valence shell, the greater the shielding effect

40. Ionization Energy DOWN a Group
DOWN THE GROUP IONIZATION ENERGY
DECREASES
greater distance from the nucleus,
the easier to lose an electron
(less energy needed)

41. Ionization Energy ACROSS a Period
As atomic radius decreases there is a greater attraction between protons and electrons.
The stronger the attraction, the more energy needed to remove an electron.
The more electrons present, the more energy required to remove them all to become STABLE

42. Ionization Energy ACROSS a Period
ACROSS THE PERIOD IONIZATION ENERGY
INCREASES
more electrons on an energy level,
more energy required to remove them all

43. Filled n=1 shell
Filled n=2 shell
Filled n=3 shell
Filled n=4 shell
Filled n=5 shell

44. : . H F . : : Electronegativity
tendency for an atom to attract electrons
It is a “tug of war” between the two atoms of a bond . : .
H F . : :
Which is the more electronegative element?

45. Electronegativity ACROSS the Period
As you go across a period atomic radius decreases because there is a greater attraction between protons and electrons
Metals do not attract electrons.
Non-metals do attract electrons.

46. Electronegativity ACROSS the Period
ACROSS THE PERIOD ELECTRONEGATIVITY
INCREASES
stronger the attraction,
the easier to add more electrons

47. Electronegativity DOWN the Group
The farther away from the nucleus, the greater the shielding effect
The larger the atom, the less likely it is to accept more electrons.

48. Electronegativity DOWN the Group
DOWN THE GROUP ELECTRONEGATIVITY
DECREASES
farther the distance from the nucleus,
more difficult to attract electrons

49. Increases with ability to attract and hold an electron
Electron Affinity
the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
Increases with ability to attract and hold an electron

50. Electron Affinity
A large positive value means the anion is very stable
X (g) + e X-(g)
F (g) + e X-(g)
DH = -328 kJ/mol
EA = +328 kJ/mol
O (g) + e O-(g)
DH = -141 kJ/mol
EA = +141 kJ/mol

51. Electron Affinity ACROSS the Period
Across the period the energy released becomes more negative, making electron affinity more positive
As the atomic radius decreases and shielding is constant, it is easier to attract an electron.
Non-metals want to attract electrons to become stable

52. Electron Affinity ACROSS the Period
ACROSS THE PERIOD ELECTRON AFFINITY
INCREASES
more electrons and smaller,
easily forms anions

53. Electron Affinity DOWN the Group
Down a group, the energy released becomes less negative so electron affinity is smaller
The larger the atom, the more difficult to accept electrons
Metals want to lose electrons.

54. Electron Affinity DOWN the Group
DOWN THE GROUP ELECTRON AFFINITY
DECREASES
farther the distance from the nucleus,
forms poor anions

56. Reactivity of Metals Video 1 Reactivity of Metals Video 2
Other Trends
Reactivity of Metals Video 1
Reactivity of Metals Video 2

57. Increasing reactivity
METAL REACTIVITY
Increasing reactivity

58. Increasing reactivity
NONMETAL REACTIVITY
Increasing reactivity