1. Periodic Table
Larry Scheffler
Lincoln High School

2. The Periodic Table-Key Questions
What is the periodic table ?
What information is obtained from the table ?
How can elemental properties be predicted based on the Periodic Table?

3. Periodic Table
The development of the periodic table brought a system of order to what was otherwise an collection of thousands of pieces of information
The periodic table is a milestone in the development of modern chemistry. It not only brought order to the elements but it also enabled scientists to predict the existence of yet undiscovered
elements.

4. Early Attempts to Classify Elements
Dobreiner’s Triads (1827)
Classified elements in sets of three having similar properties.
Found that the properties of the middle element were approximately an average of the other two elements in the triad.

5. Dobreiner’s Triads Cl Br I 35.5 79.9 126.9 81.2 1.56 3.12 4.95 3.25 Ca
Element
Atomic Mass
Average
Density Cl Br I
35.5
79.9
126.9
81.2
1.56
3.12
4.95
3.25 Ca Sr Ba
40.1
87.6
137.3
88.7
1.55
2.6
3.5
2.53
Note: In each case, the numerical values for the atomic mass and density of the middle element are close to the averages of the other two elements

6. Newland’s Octaves -1863
John Newland attempted to classify the then 62 known elements of his day.
He observed that when classified according to atomic mass, similar properties appeared to repeat for about every eighth element
His Attempt to correlate the properties of elements with musical scales subjected him to ridicule.
In the end his work was acknowledged and he was vindicated with the award of the Davy Medal in 1887 for his work.

7. Dmitri Mendeleev
Dmitri Mendeleev is credited with creating the modern periodic table of the elements.
He gets the credit because he not only arranged the atoms, but he made predictions based on his arrangement which were shown to be quite accurate.

8. Mendeleev’s Periodic Table
Mendeleev organized all of the elements into one comprehensive table.
Elements were arranged in order of increasing mass.
Elements with similar properties were placed in the same row.

9. Mendeleev’s Periodic Table

10. Mendeleev’s Periodic Table
Mendeleev left some blank spaces in his periodic table. At the time the elements gallium and germanium were not known. He predicted their discovery and estimated their properties

11. Periodic Table
The Periodic Table has undergone several modifications before it evolved in its present form. The current form is usually attributed to Glenn Seaborg in 1945

12. Periodic Table Expanded View
The Periodic Table can be arranged by energy sub levels The s-block is Group IA and & IIA, the p-block is Group IIIA - VIIIA. The d-block is the transition metals, and the f-block are the Lanthanides and Actinide metals
The way the periodic table usually shown is a compressed view. The Lanthanides and actinides (F block)are cut out and placed at the bottom of the table.

13. Periodic Table: Metallic Arrangement
Layout of the Periodic Table: Metals vs. nonmetals
Nonmetals
Metals

14. The Three Broad Classes Are Main, Transition, Rare Earth
Main (Representative), Transition metals,
lanthanides and actinides (rare earth)

15. Reading the Periodic Table: Classification
Nonmetals, Metals, Metalloids, Noble gases

17. Periodic Table: The electron configurations are inherent in the periodic tableLi
2s1 Be
2s2 B
2p1 C
2p2 N
2p3 O
2p4 F
2p5 Ne
2p6 Na
3s1 Mg
3s2 Al
3p1 Si
3p2 P
3p3 S
3p4 Cl
3p5 Ar
3p6 K
4s1 Ca
4s2 Sc
3d1 Ti
3d2 V
3d3 Cr
4s13d5 Mn
3d5 Fe
3d6 Co
3d7 Ni
3d8 Cu
4s13d10 Zn
3d10 Ga
4p1 Ge
4p2 As
4p3 Se
4p4 Be
4p5 Kr
4p6 Rb
5s1 Sr
5s2 Y
4d1 Zr
4d2 Nb
4d3 Mo
5s14d5 Tc
4d5 Ru
4d6 Rh
4d7 Ni
4d8 Ag
5s14d10 Cd
4d10 In
5p1 Sn
5p2 Sb
5p3 Te
5p4 I
5p5 Xe
5p6 Cs
6s1 Ba
6s2 La
5d1 Hf
5d2 Ta
5d3 W
6s15d5 Re
5d5 Os
5d6 Ir
5d7 Ni
5d8 Au
6s15d10 Hg
5d10 Tl
6p1 Pb
6p2 Bi
6p3 Po
6p4 At
6p5 Rn
6p6 Fr
7s1 Ra
7s2 Ac
6d1 Rf
6d2 Db
6d3 Sg
7s16d5 Bh
6d5 Hs
6d6 Mt
6d7

18. Periodic Table Organization------ Groups or Families
Vertical columns in the periodic table are known as groups or families The elements in a group have similar electron configurations

19. Periodic Table Organization ---- Periods
Horizontal Rows in the periodic table are known as Periods The Elements in a period undergo a gradual change in properties as one proceeds from left to right

20. Periodic Properties
Elements show gradual changes in certain physical properties as one moves across a period or down a group in the periodic table. These properties repeat after certain intervals. In other words they are PERIODIC
Periodic properties include:
-- Ionization Energy
-- Electronegativity
-- Electron Affinity
-- Atomic Radius
-- Ionic Radius

21. Trends in Ionization Energy
Ionization energy is the energy required to
Remove an electron from an atom
Ionization energy increases across a period because the positive charge increases.
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty (they like to GAIN electrons).

22. Trends in Ionization Energy
The ionization energy increases UP a group
Because size increases due to an effect known as the Shielding Effect

23. Ionization Energies

24. Ionization Energies are Periodic

25. Electronegativity
Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to itself.
This concept was first proposed by Linus Pauling ( ). He later won the Nobel Prize for his efforts

26. Periodic Trends: Electronegativity
In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.
In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

27. Trends in Electronegativity

28. Electronegativity

29. Electronegativity

30. Electron Affinities

31. Electron Affinities Are Periodic
Electron Affinity v Atomic Number

32. The Electron Shielding Effect
Electrons between the nucleus and the valence electrons repel each other making the atom larger.

33. Atomic Radius The radius increases on going down a group.
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
The radius decreases on going across a period.

34. Atomic Radius
The radius decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.
Large
Small

35. Atomic Radius

36. Atomic Radius

37. Trends in Ion Sizes
Radius in pm

38. Cations
Cations (positive ions) are smaller than their corresponding atoms

39. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

40. Ionic Radius CATIONS are SMALLER than the atoms from which they come.Li +
, 78 pm
2e and 3 p
Forming a cation.
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so the radius DECREASES.

41. Ionic Radius for Cations
Positve ions or cations are smaller than the corresponding atoms.
Cations like atoms increase as one moves from top to bottom in a group.

42. Anions
Anions (negative ions) are larger than their corresponding atoms

43. Ionic Radius-Anions Forming an anion.
, 133 pm
10 e- and 9 p+
F pm
9e- and 9p+
Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.

44. Ionic Radii for Anions Negative ions or anions are larger than the
corresponding atoms.
Anions like atoms increase as one moves from top to bottom in a group.

45. Ionic Radius for an Isoelectronic Group
Isoelectronic ions have the same number of electrons.
The more negative an ion is the larger it is and vice versa.

46. Summary of Periodic Trends

47. Properties of the Third Period Oxides

48. Properties of the Third Period Chlorides

49. The D Block Elements
The d block elements fall between the s block and the p block.
They share common characteristics since the orbitals of d sublevel of the atom are being filled.

50. The D Block Elements
The D block elements include the transition metals. The transition metals are those d block elements with a partially filled d sublevel in one of its oxidation states.
Since the s and d sublevels are very close in energy, the d block elements show certain special characteristics including:
Multiple oxidation states
The ability to form complex ions
Colored compounds
Catalytic behavior
Magnetic properties

51. Some common D block oxidation states
The D Block Elements
The d electrons are close in energy to the s electrons.
D block elements may lose 1 or more d electrons as well as s electrons. Hence they often have multiple oxidation states
Some common D block oxidation states

52. Multiple Oxidation States
There is no sudden sharp increase in ionization energy as one proceed through the d electrons as there would be with the s block.
D block elements can lose or share d electrons as well as s electrons, allowing for multiple oxidation states.
Most d Block elements have a +2 oxidation State which corresponds to the loss of the two s electrons.
This is especially true on the right side of the d block, but less true on the left.
---- For example Sc+2 does not exist, and
Ti+2 is unstable, oxidizing
in the presence of any
water to the +4 state.

53. Complex Ions
The ions of the d block and the lower p block have unfilled d or p orbitals.
These orbitals can accept electrons either an ion or polar molecule, to form a dative bond. This attraction results in the formation of a complex ion.
A complex ion is made up of two or more ions or polar molecules joined together.
The molecules or ions that surround the metal ion donating the electrons to form the complex ion are called ligands.

54. Complex Ions K3Fe(CN)6 Cu(NH3)42+
Compounds that are formed with complex ions are called coordination compounds
Common ligands
Complex ions usually have either 4 or 6 ligands.
K3Fe(CN) Cu(NH3)42+

55. Complex Ions
The formation of complex ions stabilizes the oxidations states of the metal ion and they also affect the solubility of the complex ion.
The formation of a
complex ion often has
a major effect on the
color of the solution of
a metal ion.

56. The D Block Colored Compounds
In an isolated atom all of the d sublevel electrons have the same energy.
When an atom is surrounded by charged ions or polar molecules, the electric field from these ions or molecules has a unequal effect on the energies of the various d orbitals and d electrons.
The colors of the ions and complex ions of d block elements depends on a variety of factors including:
The particular element
The oxidation state
The kind of ligands bound to the element
Various oxidation
states of Nickel (II)

57. Colors in the D Block
The presence of a partially filled d sublevels in a transition element results in colored compounds.
Elements with completely full or completely empty subshells are colorless,
For example Zinc which has a full d subshell. Its compounds are white
A transition metal ion is colored, if it absorbs light in the visible range (
nanometers).
If the compound absorbs a
particular wavelength of light its
color will be the composite of those
wavelengths that it does not absorb.
In other words it shows its
complimentary color.

58. Colors and d Electron Transitions
When ligands are attached to transition metal ions, the d orbitals may split into two groups. Some of the orbitals are at a lower energy than the others
The difference in energy of these orbitals varies slightly with the nature of the ligand or ion surrounding the metal ion
The energy of the transition: ∆E =hn may occur in the visible region.
When white light passes through a compound of a transition metal, light of a particular frequency is absorbed as an electron is promoted from a lower energy d orbital to a higher one.
The result is a colored compound

59. Magnetic Properties
Paramagnetism --- Molecules with one or more unpaired electrons are attracted to a magnetic field. The more unpaired electrons in the molecule the stronger the attraction. This type of behavior is called
Diamagnetism --- Substances with no unpaired electrons are weakly repelled by a magnetic field.
Transition metal complexes with unpaired electrons exhibit simple paramagnetism.
The degree of paramagnetism depends on the number of unpaired electrons

60. Catalytic Behavior
Many D block elements are catalysts for various chemical reactions
Catalysts speed up the rate of a reaction with out being consumed.
The transition metals form complex ions with ligands that can donate lone pairs of electrons.
This results in close contact between the metal ion and the ligand.
Transition metals also have a wide variety of oxidation states so they gain and lose electrons in oxidation- reduction reactions

61. Some Common D Block Catalysts
Examples of D block elements that are used as catalysts:
Platnium or
rhodium in a
catalytic converter
MnO2 decomposition
of hydrogen peroxide
V2O5 in the contact
process
Fe in Haber process
Ni in conversion of
alkenes to alkanes

62. The Periodic Table--Summary
The periodic table is a classification system. Although we are most familiar with the periodic table that Seaborg proposed more than 60 years ago, several alternate designs have been proposed.

63. Alternate Periodic Tables

64. Alternate Periodic Tables II