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Periodic Table Larry Scheffler Lincoln High School.

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Presentation on theme: "Periodic Table Larry Scheffler Lincoln High School."— Presentation transcript:

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 table
Li 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

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