1. The History of the Modern Periodic Table

2. During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table.

3. Johann Dobereiner Model of triads 1780 - 1849
In 1829, he classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical properties and orderly physical properties.
(ex. Cl, Br, I and Ca, Sr, Ba)
Model of triads

4. John Newlands Law of Octaves 1838 - 1898
In 1863, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element.
Law of Octaves

5. John Newlands 1838 - 1898 Law of Octaves
Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society.
Law of Octaves

6. John Newlands 1838 - 1898 Law of Octaves
His law of octaves failed beyond the element calcium.
WHY?
Would his law of octaves work today with the first 20 elements?
Law of Octaves

7. Dmitri Mendeleev
In 1869 he published a table of the elements organized by increasing atomic mass.

8. Lothar Meyer
At the same time, he published his own table of the elements organized by increasing atomic mass.

9. Elements known at this time

10. Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass.
Both left vacant spaces where unknown elements should fit.
So why is Mendeleev called the “father of the modern periodic table” and not Meyer, or both?

11. Mendeleev...
stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. (He corrected the atomic masses of Be, In, and U)
was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.

12. After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.

13. However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined. By looking at our modern periodic table, can you identify what problems might have caused chemists a headache?
Ar and K
Co and Ni
Te and I
Th and Pa

14. Henry Moseley
In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.
*“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.”

15. Henry Moseley
His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

16. Glenn T. Seaborg
After co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became known as the Actinide series.

17. Glenn T. Seaborg
He is the only person to have an element named after him while still alive.
"This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize."

18. Periodic Table Geography

19. The horizontal rows of the periodic table are called PERIODS.

20. The elements in any group of the periodic table have similar physical and chemical properties!
The vertical columns of the periodic table are called GROUPS, or FAMILIES.

21. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

22. Alkali Metals

23. Alkaline Earth Metals

24. Transition Metals

25. InnerTransition Metals
These elements are also called the rare-earth elements.
InnerTransition Metals

26. Halogens

27. Noble Gases

28. The s and p block elements are called REPRESENTATIVE ELEMENTS.

29. The periodic table is the most important tool in the chemist’s toolbox!

30. S block
P block
D block
F block

31. Electron Configuration and Periodic Trends
Atomic Radii – The size of an atom – one half the distance between the nuclei of two identical atoms bonded together

32. Atomic Radii
Decreases as you go across a period due to the added positive charge to the nucleus.
Increases down a group due to the “shielding effect” caused by the addition of new energy levels. The inner energy levels act in a way to shield the attractive charges of the nucleus for the outer electrons.

33. Ionization Energy – the energy required to strip away an electron from an atom A + energy  A+ + e-
Ion – atom or group of atoms that have a positive or negative charge.
Ionization – the process that results in the formation of an ion
Ionization energy generally increases as you go across a period. Alkali Metals have a very low ionization energy….. Why?????? Halogens have a very high IE…why????
Ionization energy generally decreases as you move down a group
First Ionization Energy IE1 – is the amount of energy needed to remove a first electron. Second Ionization Energy IE2 – is the amount of energy needed to remove a second electron,

34. Electron Affinity – the energy change that occurs when an electron is acquired by a neutral atom. A + e-  A- + energy Period Trends – generally decreases as you move across a period. Group trends – generally increases as you go down a group. Ionic Radii – The size of the resulting ion. Cation – positively charged ion resulting from the loss of one or more electrons (metals) Anions – negatively charged ion resulting from the gain of one or more electrons (nonmetals) Period Trend – generally decreases from groups Large jump in size in group 15, then continues to decrease to group 18. Group trend – increases down a group due to the “shielding effect”

35. Valence Electrons – the electrons available to be lost, gained, or shared in the formation of chemical compounds. Electrons in the outer energy level.
Group 1 – 1 valence electron
Group 2 – 2 valence electrons
Group 13 – 3 valence electrons
Group 14 – 4 valence electrons
Group 15 – 5 valence electrons
Group 16 – 6 valence electrons
Group 17 – 7 valence electrons
Group 18 – 8 valence electrons (except helium)

36. Electronegativity – measure of the ability of the atom in a chemical compound to attract electrons. “Electron Hunger”
Period Trend – increases across a period
Group Trend – decreases down a group