1. A Chemist’s Most Important Tool
The Periodic Table
A Chemist’s Most Important Tool

2. The modern periodic table
Periods – Horizontal Rows
Groups or Families – Vertical columns
Elements in groups have similar physical and chemical properties.

3. Across the Periodic Table
Periods: Are arranged horizontally across the periodic table (rows 1-7)
These elements have the same number of valence shells.
2nd Period
6th Period

4. Periodic Table: Metallic arrangement
Layout of the Periodic Table: Metals vs. nonmetals
Nonmetals
Metals

5. Down the Periodic Table
Family: Are arranged vertically down the periodic table (columns or group, or 1-8 A,B)
These elements have the same number electrons in the outer most shells, the valence shell.
Alkali Family:
1 e- in the valence shell
Halogen Family:
7 e- in the valence shell

6. Infamous Families of the Periodic Table
Notable families of the Periodic Table and some important members:
Alkali
Alkaline (earth)
Transition Metals
Noble Gas
Halogen
Chalcogens

7. Main group & transition elements

8. “Blocks” of the Periodic Table

9. The Development of the Periodic Table
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.

10. Early Classification Dobereiner (1829) Newlands (1863) Model of triads
Groups of 3 ie: Ca, Sr, Ba OR Cl, Br, I
Newlands (1863)
Law of Octaves
Worked well for first 20 elements

11. The First Periodic Table
Dmitri Mendeleev (1869)
Arranged elements by increasing atomic mass
Noticed a regular pattern in chemical and physical properties
Left blank spaces – Predicted the existence of elements not yet discovered based on properties

12. Support for Mendeleev’s Table
He was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.
After discovery of Sc, Ga, & Ge his predictions were shown to be correct, lending support to his table.

13. A Problem
The combination of the discovery of new elements & better ability to measure mass led to some problems with Mendeleev’s table. Ar K Co Ni Te I Th Pa
Atomic # 18 19 27 28 52 53 90 91
Atomic mass
39.9
39.09
58.9
58.7
127.6
126.9
232.03
231.03

14. A new arrangment
In 1913 Henry Moseley was able to determine the nuclear charges of the elements
He used this information to rearrange the elements in order of atomic number.
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 Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

15. A “final” adjustment
After co-discovering 10 new elements, in Glenn Seaborg moved 14 elements from the body of the periodic table to the Actinide Series.
He is the only person to have an element named after him while he was still alive.

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

17. General Explanations
Nuclear charge is a measure of the strength of a nucleus’ pull – (# of protons)
Increases left to right

18. Shielding
Shielding effect plays a major role in how strongly the nucleus can pull on its outermost electrons – This mainly effects group trends.

19. Atomic Radius Bonding atomic radius tends to…
…decrease from left to right across a row
due to increasing p+.
…increase from top to bottom of a column
due to increasing e- energy levels.
Atomic Radius

20. Sizes of Ions Ionic size depends upon:
Nuclear charge.
Number of electrons.
Orbitals in which electrons reside.
Cations are smaller than parent atom
Anions are larger

21. Sizes of Ions Ions increase in size as you go down a column.
Due to increasing value of n.

22. Sizes of Ions
In an isoelectronic series, ions have the same number of electrons.
Ionic size decreases with an increasing nuclear charge.

23. Ionization Energy
Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
First ionization energy is that energy required to remove first electron.
Second ionization energy is that energy required to remove second electron, etc.

24. Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
Generally, as one goes across a row, it gets harder to remove an electron.

25. Electron Affinity
Energy change accompanying addition of electron to gaseous atom:
Cl + e−  Cl−
What group of elements tend to add electrons?

26. Trends in Electron Affinity
Group: Decreases top to bottom
Period: Increases from left to right

27. Electronegativity
Electronegativity: tendency for an atom to attract electrons to itself (similar to Electron Affinity)
Group Trends: decreases from top to bottom on the periodic table
Period Trends: increases from left to right

28. Summary of Trend Periodic Table and Periodic Trends
1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT

29. Metals versus Nonmetals
Differences between metals and nonmetals tend to revolve around these properties.

30. Metals versus Nonmetals
Metals tend to form cations.
Nonmetals tend to form anions.
Metals get more reactive down a column (easier to lose electrons)
Nonmetals get LESS reactive down a column (easier to gain electrons)