1. CHEMICAL PERIODICITY

2. Development of Periodic Table
Dmitri Mendeleev, Russian Chemist in the mid-1800’s arranged the 70 known elements into a systematic way.
Put the names of elements on a card (Just like you arranged the pieces)
He also put atomic masses, physical and chemical properties.
When arranged in the order of atomic mass, he discovered a repeated, or periodic, pattern.
He left blank spaces where he didn’t know the element for the spot.
By doing this, he constructed the first periodic table for elements.

3. Mendeleev’s Problems
Why could most of the elements be arranged in the order of increasing atomic mass but a few could not?
What was the reason for chemical periodicity?

4. Henry Moseley’s Solution
Instead of arranging in atomic mass, in1913 Henry Moseley arranged in increasing order of nuclear charge. His work led to modern definition and recognition of atomic number.
He discovers that it is the number of electrons that affect an atoms’ reactivity, or chemical property.

5. Parts of the Periodic Table
Periods:
The horizontal rows
Periodic Law:
When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Groups: (a.k.a. Families)
The vertical columns

6. Parts of the Periodic Table, cont.
Representative Elements:
Group A elements. (The tall parts of the periodic table.)
These can be divided into 3 groups:
Metals
Nonmetals
Metalloids

7. Metals Malleable-Very easy to be hammered or beaten into thin sheets
Ductile-Very easy to produce a wire
High luster-De-excitation of electrons causes shiny appearance of metal surface
High electrical conductivity - Due to the free movement of electrons
High thermal conductivity
Most are solid at room temperature
Metallic Bond Strength
Directly proportional to the heat of vaporization
80% of all elements are metals

8. Nonmetals Upper right portion of the periodic table
Generally nonlustrous
Generally poor conductors of electricity
Some are gases at room temperature and some brittle solids
One is even a liquid at room temperature

10. Metalloids These border the stair step, except for aluminum
These have properties of both metals and nonmetals
Si, Ge, As, Sb, Te

11. Group 1 – Alkali Metal
Very reactive!Not found in nature as free elements
Combine vigorously with most nonmetals
React strongly with water to produce hydrogen gas and aqueous solutions of substances known as “alkalis.”
Usually stored in kerosene (paraffin oil)

12. Group 2 – Alkali Earth Metal
Harder, denser, and stronger than alkali metal
Higher melting points
Not so reactive compared to alkali metal; Still very reactive! Not found as free element

13. Group 3 – 12 Transition Metal
Good conductors of electricity and have a high luster
Less reactive than alkali and alkali-earth metal
Some (e.g. platinum and gold) are so unreactive that they do not form compounds easily.
Some are found as free element.

14. Group 17 - Halogens Most reactive nonmetal
React vigorously with most metals to form “salts”
Need one electron to obtain stable noble-gas configuration
F and Cl are gaseous.
Br is liquid.
Iodine is solid.

15. Group 18 – Noble Gases Lack chemical reactivity Very stable
Octet electronic configuration in the outermost energy level (or “shell”)

16. 16

17. 17

18. 18

19. Categories of electrons
Core (inner) electrons – all those shared by the previous noble gas
Outer electrons – those in the highest occupied energy level
Similar chemical properties of elements in groups is a result of similar outer electron configurations
In the main group, the group number equals the number of outer electrons
Valence electrons – those involved in bonding
In the main group, the outer electrons are valence
In the transition metals can include some d electrons 19

20. Trends in metallic behavior20

21. Metallic Behavior Metallic behavior increases left and down
Metals tend to lose electrons in reactions
Highly metallic elements, are likely to make positive ions
Least metallic elements are likely to make negative ions
In middle are more likely to make covalent bonds 21

22. Ion (Cation and Anion) # of electrons added/removed= # of charge
Neutral atom
Add electron
Negative ion (Anion) F
+ 1e-
F-1 (Anion)
Neutral atom
Remove electron
Positive ion (Cation) Ca
-2e-
Ca+2 (Cation)
# of electrons added/removed= # of charge

23. Main-group ions and the noble gas configurations
Octet (Duplet) Rule 23 23

24. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1
8.2

25. Periodic Law Be 1S2 2S2 2, 2 Mg 1S2 2S2 2p6 3S2 2, 8, 2
All the elements in a group have the same electron configuration in their outermost shells
Example: Group 2
Be 1S2 2S2
2, 2
Mg 1S2 2S2 2p6 3S2
2, ,
Ca 1S2 2S2 2p6 3S2 3p6 4S2
2, , ,
LecturePLUS Timberlake 25 25

26. Periodic Table and Electron Configuration
Groups = s level
Groups 3-8 = p level
Transition = d level
Lantanides = f level
s1 s p1 p2 p3 p4 p5 p6 1 2
d1 - d10 4 5 6
f1 - f14 26

27. The relation between orbital filling and the periodic table
Figure 8.13
The relation between orbital filling and the periodic table 27

28. Figure 8.12 28

29. Periodic Patterns s p d (n-1) f (n-2) 1 2 3 4 5 6 7 6 7
© 1998 by Harcourt Brace & Company 29 29

30. Trends in Some Periodic Properties
The physical and chemical behavior of the elements is based on the electron configurations of their atoms.
e- configurations can be used to explain many of the repeating or “periodic” properties of the elements 30

31. Defining metallic and covalent radii Figure 8.1431

32. Atomic Radius - size of atom
8.3

33. Definition of Some Periodic Properties
Atomic Radius/Ionic Radius- size of atom/Ion
Electronegativity is a measure of an atom’s attraction for another atom’s electrons.
Ionization energy: the energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)
Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). 33

34. Atomic radii of the main-group and transition elements.
Vertically
The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family.
With each step down the family, we add an entirely new energy level, making the atoms larger with each step.
Horizontally
Each step adds a proton and an electron (and 1 or 2 neutrons).
Electrons are added to existing energy level.
The effect is that the more positive nucleus has a greater pull on the electron cloud.
The nucleus is more positive and the electron cloud is more negative.
The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period. 34

35. Electron Shells and Sizes of Atoms
Size in main group elements follow 2 general rules
1. AR increases going down in a group
2. AR decreases going left to right in a period
Why?
Moving down in a group:
Zeff is constant because Z and S increase equally
Electrons in a higher n and therefore larger
Moving across:
Z increases and S does not (electrons in same shell don't shield each other well) so e- attracted more strongly and are closer/smaller
In transition metals there is an initial decrease in size but then remains relatively constant 35

36. Periodicity of atomic radius
Figure 8.16
Periodicity of atomic radius 36

37. Ranking Elements by Atomic Size
SAMPLE PROBLEM 8.3
Ranking Elements by Atomic Size
PROBLEM:
Using only the periodic table (not Figure 8.15), rank each set of main group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
PLAN:
Elements in the same group increase in size and you go down; elements decrease in size as you go across a period.
SOLUTION:
(a) Sr > Ca > Mg
These elements are in Group 2A(2).
(b) K > Ca > Ga
These elements are in Period 4.
(c) Rb > Br > Kr
Rb has a higher energy level and is far to the left. Br is to the left of Kr.
(d) Rb > Sr > Ca
Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr. 37

38. Effective nuclear charge (Zeff) is the “positive charge” felt by an electron.
Zeff = Z - s
0 < s < Z (s = shielding constant)
Zeff  Z – number of inner or core electrons
Zeff
Core Z
Radius (pm) Na Mg Al Si 11 12 13 14 10 1 2 3 4
186
160
143
132

39. Effective Nuclear Charge (Zeff)
increasing Zeff
increasing Zeff

40. She’s unhappy and negative.
Ions
Here is a simple way to remember which is the cation and which the anion:
This is Ann Ion.
This is a cat-ion.
She’s unhappy and negative.
He’s a “plussy” cat!

41. Cation Formation
Effective nuclear charge on remaining electrons increases.
Na atom
1 valence electron
Remaining e- are pulled in closer to the nucleus. Ionic size decreases.
11p+
Valence e- lost in ion formation
Result: a smaller sodium cation, Na+

42. Anion Formation
A chloride ion is produced. It is larger than the original atom.
Chlorine atom with 7 valence e-
17p+
One e- is added to the outer shell.
Effective nuclear charge is reduced and the e- cloud expands.

43. Cation is always smaller than atom from which it is formed.
Anion is always larger than atom from which it is formed.

44. The Radii (in pm) of Ions of Familiar Elements

45. Ionic Size Cations are smaller than the parent atom
Anions are larger than the parent atom.
Ionic size increase down in a group
In an isoelectronic series, the most negative ion is largest, the most positive is smallest.
Atoms making more than one ion, most positive is smallest. 45 45

46. Ionization Energy This is the second important periodic trend.
The atom has been “ionized” or charged.
The number of protons and electrons is no longer equal.
The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ)
The larger the atom is, the easier its electrons are to remove.
Ionization energy and atomic radius are inversely proportional.
Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

47. Ionization Energy (Potential)
Draw arrows on your help sheet like this:

48. Ionization energy Trends: But Why?
Generally as size decreases, IE1 increases.
1. IE1 generally increases from left to right, some exceptions
2. IE1 generally decreases going down in a group
3. transition and f-block elements have much smaller variances in IE1
But Why?
Across - increasing Zeff and smaller size(e- closer to nucleus)
exceptions - Be to B because s shields p and lowers Zeff
N to O because repulsions in first paired electron
Down - Zeff constant and larger so easier to ionize 48

49. Periodicity of first ionization energy (IE1)
Figure 8.17
Periodicity of first ionization energy (IE1) 49

50. First ionization energies of the main-group elements
Figure 8.18
First ionization energies of the main-group elements 50

51. Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.
I1 + X (g) X+(g) + e-
I1 first ionization energy
I2 + X+(g) X2+(g) + e-
I2 second ionization energy
I3 + X2+(g) X3+(g) + e-
I3 third ionization energy
I1 < I2 < I3

52. Ranking Elements by First Ionization Energy
SAMPLE PROBLEM 8.4
Ranking Elements by First Ionization Energy
PROBLEM:
Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:
(a) Kr, He, Ar
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
PLAN:
IE decreases as you proceed down in a group; IE increases as you go across a period.
SOLUTION:
(a) He > Ar > Kr
Group 8A(18) - IE decreases down a group.
(b) Te > Sb > Sn
Period 5 elements - IE increases across a period.
(c) Ca > K > Rb
Ca is to the right of K; Rb is below K.
(d) Xe > I > Cs
I is to the left of Xe; Cs is furtther to the left and down one period. 52

53. Ionization energy Result - Rx only involve outer e-
Second ionization energy IE2 – to remove 2nd e-
Always larger than IE1
as e- are removed, Z remains constant and remaining e- are harder to remove
A huge increases occurs in ionization energies when core electrons are reached because of the lower shielding/higher Zeff.
Result - Rx only involve outer e- 53

54. The first three ionization energies of beryllium (in MJ/mol)
Figure 8.19
The first three ionization energies of beryllium (in MJ/mol) 54

55. 55

56. Identifying an Element from Successive Ionization Energies
SAMPLE PROBLEM 8.5
Identifying an Element from Successive
Ionization Energies
PROBLEM:
Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:
IE1
IE2
IE3
IE4
IE5
IE6
1012
1903
2910
4956
6278
22,230
PLAN:
Look for a large increase in energy which indicates that all of the valence electrons have been removed.
SOLUTION:
The largest increase occurs after IE5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s23p3 and the element must be phosphorous, P (Z = 15).
The complete electron configuration is 1s22s22p63s23p3. 56

57. Electronegativity
Electronegativity is a measure of an atom’s attraction for another atom’s electrons.
It is an arbitrary scale that ranges from 0 to 4.
The units of electronegativity are Paulings.
Generally, metals are electron givers and have low electronegativities.
Nonmetals are are electron takers and have high electronegativities.
What about the noble gases?

58. Electronegativity
Your help sheet should look like this:

59. Overall Reactivity
This ties all the previous trends together in one package.
However, we must treat metals and nonmetals separately.
The most reactive metals are the largest since they are the best electron givers.
The most reactive nonmetals are the smallest ones, the best electron takers.

60. Overall Reactivity
Your help sheet will look like this:

61. Electron Affinity What does the word ‘affinity’ mean?
Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ).
Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

62. Electron Affinity
Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy.
If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic.
This is true for the alkaline earth metals and the noble gases.

63. Electron Affinity
Your help sheet should look like this: + +

64. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
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

65. electron affinities
energy required to put an e- on an atom (making it negative or more negative)
EA1 usually negative
EA2 always positive
Irregular trend: increases going right and up
Remember: Full shells best, full subshells good, and half full is kinda cool
Cl and other halogens need only 1 e- so most exothermic reaction
noble gases have endo because new electron in higher n
Mg, Be endo because new e in higher l
N endo because new e paired 65

66. Electron affinities of the main-group elements
Figure 8.20
Electron affinities of the main-group elements 66

67. Trends in three atomic properties
Figure 8.21
Trends in three atomic properties 67

68. Atomic radii increases Atomic radii decreases
Electronegativity Increases
Atomic radii decreases 68 68