Periodic Table.1

The Periodic Table-Key Questions What is the periodic table ? What information does the table provide ? ? How can one use the periodic table to predict the properties of the elements ?.2

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 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. 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 elements that had not yet been discovered..3

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 also made predictions based on his arrangements His predictions were later shown to be quite accurate..4

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..5

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..6

The Modern 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

The Three Broad Classes are the Representative, Transition, & Rare Earth Main (Representative), Transition metals, lanthanides and actinides (rare earth).8

Periodic Table: The electron configurations are inherent in the periodic table B 2p 1 H 1s 1 Li 2s 1 Na 3s 1 K 4s 1 Rb 5s 1 Cs 6s 1 Fr 7s 1 Be 2s 2 Mg 3s 2 Ca 4s 2 Sr 5s 2 Ba 6s 2 Ra 7s 2 Sc 3d 1 Ti 3d 2 V 3d 3 Cr 4s 1 3d 5 Mn 3d 5 Fe 3d 6 Co 3d 7 Ni 3d 8 Zn 3d 10 Cu 4s 1 3d 10 B 2p 1 C 2p 2 N 2p 3 O 2p 4 F 2p 5 Ne 2p 6 He 1s 2 Al 3p 1 Ga 4p 1 In 5p 1 Tl 6p 1 Si 3p 2 Ge 4p 2 Sn 5p 2 Pb 6p 2 P 3p 3 As 4p 3 Sb 5p 3 Bi 6p 3 S 3p 4 Se 4p 4 Te 5p 4 Po 6p 4 Cl 3p 5 Be 4p 5 I 5p 5 At 6p 5 Ar 3p 6 Kr 4p 6 Xe 5p 6 Rn 6p 6 Y 4d 1 La 5d 1 Ac 6d 1 Cd 4d 10 Hg 5d 10 Ag 5s 1 4d 10 Au 6s 1 5d 10 Zr 4d 2 Hf 5d 2 Rf 6d 2 Nb 4d3 Ta 5d 3 Db 6d 3 Mo 5s 1 4d 5 W 6s 1 5d 5 Sg 7s 1 6d 5 Tc 4d 5 Re 5d 5 Bh 6d 5 Ru 4d 6 Os 5d 6 Hs 6d 6 Rh 4d 7 Ir 5d 7 Mt 6d 7 Ni 4d 8 Ni 5d 8.9

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.10

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.11

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 -- Atomic Radius -- Ionic Radius.12

Metals lose electrons more easily than nonmetals. Nonmetals lose electrons with difficulty. (They like to GAIN electrons). Ionization energy increases across a period because the positive charge increases. Ionization energy is the energy required to remove an electron from an atom Trends in Ionization Energy.13

The ionization energy is highest at the top of a group. Ionization energy decreases as the atom size increases. This results from an effect known as the Shielding Effect Trends in Ionization Energy.14

Ionization Energies of the Representative Groups.15

Ionization Energies are Periodic.16

Effective Nuclear Charge Many properties depend on: –Electron configuration –How strong outer electrons are attracted to nucleus Effective nuclear charge: net electrical charge acting on an electron –Increases with nuclear charge (number of protons) –Decreases with distance from nucleus

Effective Nuclear Charge Shielding: repulsion of inner electrons on outer electrons, reducing the effect of the nuclear charge on electrons Z eff = Z – S –Z = number of protons –S = number of non valence electrons

The Electron Shielding Effect Electrons between the nucleus and the valence electrons repel each other making the atom larger..19

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..20

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..21

Trends in Electronegativity.22 Electronegativity increases across a period and up a group

Electronegativity.23

Electronegativity.24

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. 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. Atomic Radius.25

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, whereas the electrons are scattered. Large Small All values are in nanometers.26

Atomic Radius.27

Atomic Radius.28

Trends in Ion Sizes Radius in pm.29

Cations Cations (positive ions) are smaller than their corresponding atoms Cations (positive ions) are smaller than their corresponding atoms.30

CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so the radius DECREASES. Li nm 3e and 3p Li +, nm 2e and 3 p + Ionic Radius Forming a cation..31

Ionic Radius for Cations Positive ions or cations are smaller than the corresponding atoms. Cations like atoms increase as one moves from top to bottom in a group..32

Anions Anions (negative ions) are larger than their corresponding atoms Anions (negative ions) are larger than their corresponding atoms.33

Ionic Radius-Anions 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. Forming an anion. F nm 9e - and 9p + F nm 10 e - and 9 p

Does the size go up or down when gaining an electron to form an anion? Does the size go up or down when gaining an electron to form an anion? Ion Sizes.35

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..36

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..37

Summary of Periodic Trends

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..39

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.40