1. Periodic Chart

2. Diderot's Alchemical Chart of Affinities (1778)

3. Where to begin
Dobereiner observes similarities in elements. Proposes Law of Triads-Middle element in the triad had atomic weight that was the average of the other two members.

4. First Periodic Table 1862
Alexandre Beguyer de Chancourtois ( ), professor of geology at the School of Mines in Paris, publishes a periodic table constructed as a helical graph.
Ignored

5. Helix Periodic Table

6. John Newlands proposes Octet rule
having arranged the 62 known elements in order of increasing atomic weights, noted that after interval of eight elements similar physical/chemical properties reappeared.
Because these properties seemed to repeat every eight elements, Newlands called this pattern the law of octaves.

8. More Periodic History Mendeleev
Then in 1869, Russian chemist Dimitri Mendeleev ( ) proposed arranging elements by atomic weights and properties (Lothar Meyer independently reached similar conclusion but published results after Mendeleev).

9. Mendeleev noticed that the chemical properties of the elements repeated each time he started a new row.
Mendeleev made two interesting observations
1. Mendeleev’s table contains gaps that elements with particular properties should fill.
2. The elements do not always fit neatly in order of atomic mass.
Mendeleev predicted the properties of the missing elements.

10. Mendeleev’s Periodic Chart

11. More Mendeleev
Mendeleev's periodic table of 1869 contained 17 columns with two partial periods of seven elements each (Li-F & Na-Cl) followed by two nearly complete periods (K-Br & Rb-I).

12. Mendeleev Periodic Table

13. Mendeleev’s Periodic Law
Properties of the elements is a periodic function of their atomic mass

14. Why Mendeleev’s basic version of the periodic chart?
Mendeleev/Meyer periodic chart is the basis of today’s chart because it:
Predicts Atomic Properties
Indicates Trends
Indicates groups that will react/not react together

15. Mosley Periodic Table After Rutherford’s experiments
Henry Moseley ( ) subjected known elements to x-rays. He was able to derive the relationship between x-ray frequency and number of protons.

16. Modern Periodic Law
Mosley’s Periodic Law- Properties of an element are a periodic function of their atomic number.

17. Mosley’s Periodic Table

18. When the elements were arranged by increasing
atomic number, the discrepancies in
Mendeleev’s table disappeared.
Moseley’s work led to both the modern definition of atomic number, and showed that atomic number, not atomic mass, is the basis for the organization of the periodic table.

19. Last Major Change Gene Seaborg
Starting with plutonium in 1940, Seaborg discovered transuranium elements 94 to 102 and reconfigured the periodic table by placing the lanthanide/actinide series at the bottom of the table. In 1951 Seaborg was awarded the Nobel Prize in chemistry and element 106 was later named seaborgium (Sg) in his honor.

20. Modern Periodic Law
Mendeleev’s principle of chemical periodicity is known as the periodic law, which states:
The elements are arranged according to their atomic numbers, elements with similar properties appear at regular intervals.

22. A Little More about Mendeleev and the Periodic Table

24. F Block Placed in the Periodic Table

25. Circular Periodic Table

26. Kansas Board Approved Periodic Table

27. 4 Main Blocks

28. More Periodic Table
Period - rows of elements with the same energy level
Groups - elements in the same column with the same electron configuration
Metals- left of metalloids
Nonmetals- rt of metalloids

29. Majority of elements are metals

30. Special names for Some Main Groups
Four groups within the main-group elements have special names. These groups are:
alkali metals (Group 1)
alkaline-earth metals (Group 2)
halogens (Group 17)
noble gases (Group 18)

31. Elements in Group 1 are called alkali metals
Elements in Group 1 are called alkali metals. lithium, sodium, potassium, rubidium, cesium, and francium
Alkali metals are so named because they are metals that react with water to make alkaline solutions.
Because the alkali metals have a single valence electron, they are very reactive.
In losing its one valence electron, potassium achieves a stable electron configuration.
Alkali metals are never found in nature as pure elements but are found as compounds.

32. Group 2 elements are called alkaline-earth metals.
The alkaline-earth metals are slightly less reactive than the alkali metals.
They are usually found as compounds.
The alkaline-earth metals have two valence electrons and must lose both their valence electrons to get to a stable electron configuration.
It takes more energy to lose two electrons than it takes to lose just the one electron that the alkali metals must give up to become stable.

33. Elements in Group 17 of the periodic table are called the halogens.
The halogens are the most reactive group of nonmetal elements.
When halogens react, they often gain the one electron needed to have eight valence electrons, a filled outer energy level.
Because the alkali metals have one valence electron, they are ideally suited to react with the halogens.
The halogens react with most metals to produce salts.

34. Group 18 elements are called the noble gases.
The noble gas atoms have a full set of electrons in their outermost energy level.
The low reactivity of noble gases leads to some special uses.
The noble gases were once called inert gases because they were thought to be completely unreactive.
In 1962, chemists were able to get xenon to react, making the compound XePtF6.
In 1979, chemists were able to form the first xenon-carbon bonds.

35. Hydrogen Hydrogen is in a class by itself in the periodic table
It is placed in group 1 because it has a 1+ charge and a 1s1 electron configuration
It can also be placed in Group 17 because of its behavior as well

36. Periodicity
With increasing atomic number, the electron configuration of the atoms display a periodic variation.
Which leads us to Trends

37. Trends
Atomic radii
One-half the diameter of the distance between two like nuclei
Increases down a group
Decreases across a period

38. Different Atomic Radii

43. Ionic Radii
Radius of the ion after it has lost or gained an electron
- increases across the chart
- increases down the chart

46. Ionization Energy
Amount of energy required to remove an electron from an atom
Increases across the chart
Decrease down the chart

48. Electron shielding
is the reduction of the attractive force between a positively charged nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charges of the inner electrons.

49. Electron Affinity
The energy emitted upon the addition of an electron to an atom or group of atoms in a gas phase
Generally become more negative as you go across the chart\
No clear trend

51. Other trends Increase in Atomic Number and Mass
Melting and Boiling Points
-metals increase L-R then decrease to nonmetals
Electronegtivity
Energy required by an element in a compound to attract an electron
- increase from L-R
- decrease top to bottom

52. Trends in Melting Point