1. History of the Periodic Table
PeriodicTable.com

2. Who developed the first periodic table?
Dobereiner (Germany, 1817) noticed that atomic mass of Sr was closely related to the atomic mass of Ca and Ba.
He called this grouping a triad.
Several other triads were created based on his system but eventually there were not enough triads to make the system useful.

3. Who developed the modern periodic table?
Dmetri Mendeleev (Russia) and Lothar Meyer (Germany) both published almost identical systems of classifying elements.
Arranged periodic table based on their physical and chemical properties and atomic mass.

4. Why was Mendeleev credited with the periodic table?
He insisted that elements be grouped together by properties and thus by family.
The gaps in his periodic table were actually elements yet to be discovered!!
He was able to predict the properties of these undiscovered elements.

8. Prediction of Germanium’s Properties
Property
Mendeleev’s Prediction
Observed Properties
Atomic Weight 72
72.59
Density
5.5
5.35
Specific Heat
0.305
0.309
Melting point
High
947C
Color
Dark gray
Grayish white
Formula of oxide
XO2
GeO2
Density of oxide
4.7
4.70
Formula of Chloride
XCl4
GeCl4
Boiling point of chloride
A little under 100C
84C

9. Periodic Law
Mendeleev stated that the physical and chemical properties of elements vary periodically with increasing atomic mass.

10. How were the table and atomic numbers proven?
Two years after Rutherford created his atomic model, Robert Moseley developed the concept of atomic numbers.
He bombarded elements with high energy electrons and measured the frequency of x-rays given off.
Then he arranged the elements in order according to the frequency given off. It was then concluded that as the atomic mass of the elements increased the frequency also increased.
Moseley then assigned a unique whole number to each element and called it the atomic number.

11. How did this change the Mendeleev’s Periodic Table?
This allowed elements such as Argon, which has a greater atomic mass, to be placed before Potassium.
It also allowed others to determine fully the number of “holes” in the periodic table.

12. Glenn Seaborg and Actinoid Series
After discovering U, Th, Pa, Seaborg was advised to revise the periodic table.
After examining the properties of the elements, Seaborg was convinced that the elements were part of the inner transition elements because they behaved similarly to other transition metals.
Instead of creating a new system, he integrated the elements into a new section called the actinoid series, which would be inserted into the transition elements.

14. Modern Periodic Law
Elements are placed on the table according to increasing atomic number.
Trends and properties are a result of the similar atomic structure in a family.

15. Metals
And
Non-metals

16. Metals and Non Metals
For the following, write ‘M’ if it applies to metals, and ‘NM’ if it applies to nonmetals:
____ have luster or shine
____ can be gases at room temperature
____ good conductors of heat and electricity
____ high melting points
____ ductile and malleable
____ are found on the right side of the periodic table
What about metalloids?

17. Classification of Elements
Chapter 14.1

18. Classification of Elements
Animated PT
Interactive PT

19. Alkali Metals ____ column ___-block # of valence electrons ____
How many valence electrons would it like to have ______
What it will do to follow the Octet Rule ______
Charge of ion formed______

20. Alkaline Earth Metals ____ column ___-block
# of valence electrons ____
How many valence electrons would it like to have ______
What it will do to follow the Octet Rule ______
Charge of ion formed______

21. Transition and Inner Transition Metals
___, ___-blocks
# of valence electrons ____
With all the d and f electrons, can these metals really ever be happy?____
What can they do?
___________________

22. Halogens 7th column (or 17th) ___-block # of valence electrons ____
How many valence electrons would it like to have ______
What it will do to follow the Octet Rule ______
Charge of ion formed______

23. Noble Gases 8th column (or 18th) ___-block # of valence electrons ____
How many valence electrons would it like to have ______
What it will do to follow the Octet Rule ______
Charge of ion formed______

24. Periodic Trends
Chapter 14.2

25. Electron Shielding
Electron Shielding reduces the force of the nucleus’
positive charge and its outermost electron due to cancellation of charge by inner (core) electrons

27. Effective Nuclear Charge
In many atoms, there is repulsion between electrons.
Each electron is also attracted to the positive charge on the nucleus.
“Effective” or actual nuclear charge (“pull”) experienced by an valence electron is lessened as a result of repulsion (“push”) from core electrons.

28. Effective Nuclear Charge
Zeff = Z – S where Z is the nuclear charge and S is the shielding constant.
Electrons in the same type of orbital are not effective in shielding one another.
Core electrons closer to the nucleus provide effective shielding (“push”) for outer electrons.
Effective nuclear charge increases from left to right in a periodic table.

29. Atomic radii

30. Atomic radii

31. Periodic Properties
Atomic Radii:

32. Trend: atomic radii gets ________as ones goes down the column
Reason: as one moves down the column of the PT you are adding entire _____________ ___________ of electrons, increasing the # of ______ ________________ and therefore increasing the _______________ _______________ (________), and decreasing the _____________ ____________ ______________.
Trend: atomic radii gets ________ as one moves left to right
Reason: as one moves left to right, you are keeping the same number of __________ electrons (________), and increasing the number of ________________ ________________ (pull), therefore increasing the _______________ _____________ ___________.

35. Ions and Ionic Radii

36. Periodic Trends – Ionization Energy
Ionization energy is defined as the energy required to remove the electron from an atom.
As you move down a column the ionization energy _________.
R: Atoms are getting ______; ____ to lose electrons due to ____ effective nuclear charge (pull).
As you move left-to-right across a period, the ionization energy __________.
R: Atoms are getting ______; ____ to lose electrons due to ____ effective nuclear charge (pull).

37. Ionization Energy

38. Ionization Energy
The smallest elements are really good at holding onto electrons, thus the trend is opposite of Atomic Radius

39. Periodic Trends-Electronegativity
Electronegativity is defined as the relative ability of an atom to attract electrons in a chemical bond.
T: As you move down a column the electronegativity _______.
R: Atoms are getting ______; ____ able to attract electrons due to ____ exposure to outside world.
As you move left-to-right across a row, the electronegativity _________.
R: Atoms are getting ______; ____ able to attract electrons due to ____ exposure to outside world.

40. Electronegativity

41. Electronegativity

42. Reactivity is defined as the ability for an atom to react
Periodic Trends - Reactivity
Reactivity is defined as the ability for an atom to react
With reactivity, we must look at the 2 separate groups
metals and non-metals.

43. WHY??? Periodic Trends Atomic Radius Ionization Energy Reactivity
Electronegativity