1. The Periodic Table

2. History of Periodic Table
Developed by Dmitri Mendeleev 1869
Arranged 70 known elements by increasing atomic mass
Noticed a periodic recurrence of physical and chemical properties

3. Mendeleev’s Table Not completely correct
I and Te properties did not match elements in same row, switched
Predicted elements that had not been discovered
Ge, predicted mass, phy and chem properties based on location

4. Henry Moseley (1914)
Developed modern arrangement of the periodic table
According to # of protons
Periodic Law: the physical and chemical properties of the elements are periodic functions of their atomic number

5. General arrangement Periods Horizontal Rows
Elements in the same period have same valence energy level

6. General arrangement Groups (families) Vertical Columns
Elements in the same group have same # of valence electrons
Have similar chemical and physical properties

7. Groups of Elements

8. Group 1 – Alkali Metals Have 1 valence electron
Form ions with a +1 charge
Soft
Extremely reactive, do not occur as free elements in nature

9. Alkali Metals (cont’d)
React with halogens (group 17) to form compounds in a 1:1 ratio
General Formula MH
where, M = metal
H = Halogen
React with oxygen to form compounds with a 2:1 ratio
General Formula: M2O
Ex) Lithium oxide

10. Alkaline Earth Metals Have 2 valence electrons
Form ions with +2 charge
do not occur as free elements in nature

11. Alkaline Earth Metals React with halogens to form compounds
General Formula: MH2
Calcium chloride
Reacts with oxygen
1:1 ratio
Magnesium oxide

12. Halogens – Group 17 Have 7 valence electrons Most reactive nonmetals
Fluorine is most reactive

13. Transition Metals Groups 3 - 12
Have multiple charges
Lose electrons from both s and d sublevels
Form colored compounds

14. Noble Gases Have full valence shells Extremely nonreactive
Called inert gases

15. Periodic Trends

16. Atomic size Atomic radius (table S)
Def: half the distance between the nuclei of 2 like atoms

17. Atomic Radii Trend in a Group
looking from top to bottom in a group, the number of principal energy levels in each element increases
valence electrons are shielded by the inner electrons
 the valence electrons are held more loosely to the atom
As you go from top to bottom down a group, atomic size INCREASES

18. Atomic Radii Trend in a Period
As you go from left to right across a period atomic size DECREASES.
as # of protons increases, nuclear charge increases
the inner electrons (all but valence) are the same
Valence electrons are added to same energy level, but nuclear charge increases
 valence electrons are more closely attracted to the nucleus

19. Ionic Radii
as atoms become ions, the size of atom changes due to the change in # of electrons

20. Anion Formation when an atom gains electrons, ion has neg charge
radius of the atom increases
nuclear charge is spread across more electrons, less attraction to nucleus

21. Cation Formation when an atom loses electrons, ions have a + charge
radius of the atom decreases
atoms lose valence electrons, loss of outer shell

22. Ions with the same Electron Configurations (Isoelectronic)
When comparing ions that are isoelectronic, the radius of the ion with the greatest number of protons will be the smallest.
Ex) What is the Bohr configuration of the following ions? Which is largest ion?
Na+
Mg+2
O-2
F-1

23. Ionization Energy
Def: the amount of energy needed to remove the outermost electron from a neutral atom
Atoms may have more than one ionization energy
Second Ionization energy is the amt of energy needed to remove the 2nd outermost electron from an atom.
1ST ionization energies of elements listed on Table S

24. Ionization Energy Trend in a Group
the lower the IE the more loosely bound the electron
the larger the atom the more loosely bound the electron is to the atom
the ionization energy of atoms in a group will decrease looking from top to bottom

25. Ionization Energy Trend in a Period
looking from left to right in a period, atoms get smaller in size
Ionization energy increases

26. Homework Answers
Page 406
#9 Which element in each pair has the larger ionization energy?
sodium, potassium
Magnesium , phosphorous

27. Page 409
#19 Distinguish between the 1st and 2nd ionization energy of an atom.

28. Page 409
# 20 Indicate which element in each pair has the greater 1st ionization energy. a. Lithium, boron b. Magnesium, strontium c. Cesium, aluminum

29. Page 409
#21 Would you expect metals or nonmetals to have higher ionization energies?

30. Page 409
#22 Arrange the following elements in order of increasing ionization energy. a. Be, Mg, Sr b. Bi, Cs, Ba c. Na, Al, S

31. Page 409
#23 Why is there a large increase between the first and second ionization energies of the alkali metals?

32. Electronegativity
Chemical compounds are formed because atoms will lose, gain, or share valence electrons
Electronegativity – the ability of an atom to attract a pair of electrons when bonded to another atom

33. Electronegativity (cont’d)
The closer the pair of electrons is to the nucleus, the greater the attraction
as size of the atom decreases, its electronegativity increases
Electronegativity increases across a period, and decreases down a group

34. Metallic Character
Def: the measure of an atom’s ability to lose electrons and form positive ions
The stronger the metallic character, the easier it will lose electrons

35. Properties of elements with greatest metallic character:
Low ionization energies
low electronegativity

36. Trend in Metallic Character
Left to right in a Period: Metallic character decreases Top to bottom in a Group: Metallic character increases