1. The Periodic Table
Basic Concepts

2. What do the numbers mean?
This is the atomic number. It is the number of protons in a single atom of this element. It is also # of electrons. 11 Na
The symbol for this element.
This is the average atomic mass, it is the number of protons + neutrons, or the mass of the nucleus of an atom.
22.990
sodium
This is the name of the element.

3. Periods Horizontal rows on the table
Correspond to the outermost energy level of electrons being filled.

4. Groups Vertical columns on the table
All groups have number designations
Are also called families
Same/similar physical and chemical properties due to VALENCE ELECTRONS!
Some groups have special family names based upon characteristics of elements in that group

5. Regions of the Periodic Table
Metals
Nonmetals
Metalloids (semi-conductors)
Transition metals
Inner-transition metals
Lanthanide Series
Actinide Series
Transuranium elements

6. Metals

7. Properties of element types:
1) Metals- high luster, good conductors, ductile, malleable, most are solid at room temp (except Hg is liquid)
2)Nonmetals- low luster, poor conductors, very brittle, various states of matter at room temperature (ex: S is solid, O is gas, Br is liquid)
3) Semi-Conductors - sit on stair-step line metals and non-metals; have properties between metals and non-metals

8. Nonmetals

9. Semi Conductors

10. Transition Metals

11. Inner-Transition Metals

12. Regions of the Periodic Table Continued
Radioactive elements ( atomic # > 83)
Most active metals (Fr)
Most active nonmetals (F)

13. (A) Periods of the periodic table, and (B) groups of the periodic table.

14. Family Names Alkali Metals (Group 1) Alkaline Earth Metals (Group 2)
Halogens (Group 7)
Noble gases (Group 8)

15. Alkali Metals (Group 1)
Form metal hydroxides (strong bases) when reacting in water
2 Na HOH  2 NaOH H2
Are generally very reactive compared to other groups of metals
Have one valence electron
Form cations with a +1 charge

16. Alkali metals

17. Alkaline Earth Metals (Group 2)
Form metal hydroxides (strong bases) when reacting in water
Ca HOH  Ca(OH) H2
Are not as reactive as alkali metals but are generally more reactive than transition elements
Have two valence electrons
Form cations with a +2 charge

18. Alkaline earth metals

19. Halogens (Group 7) Form a multitude of salts
Are generally very reactive when compared to other nonmetals
Have seven valence electrons
Form anions with a -1 charge

20. Halogens

21. Noble Gases (Group 8) Are generally unreactive (inert)
Have eight valence electrons
Some compounds with xenon and krypton have been synthesized

22. Nobel Gases – Group 18

23. Periodic Table
Properties and Trends

24. Common Periodic Properties
Atomic Radius
Ionization Energy
Electron Affinity
Ionic Radius
Electronegativity
Metallic Character
Nonmetallic Character

25. Atomic Radii Trends on the Periodic Table
For the main group elements:
atomic radii increase going down a group
decrease going across a period.

26. Going down a group radii increases because:
Energy level is added for each successive period

27. Ionization Energy
Is defined as the energy required to remove an electron from an atom in the gas phase.
Each atom can have a series of ionization energies, since more than one electron can always be removed (except H).

28. First Ionization Energy Trends on the Periodic Table
First ionization energies generally increase across a period and decrease down a group.
Generally, the larger the atom the easier it is to remove an electron and the less ionization energy required.
WHY? The outermost electrons are found in higher energy levels as one goes down the group. Since the electrons are farther from the nucleus's pull the electrons are more easily removed.

29. Ability to lose an electron by type
Low ionization energies are typical of active metals.
High ionization energies are typical of active nonmetals.
Very high ionization energies are found with the Noble Gases

30. Plot of First Ionization Energies For Periods 1-4
Ionization Energy (kJ/mol)

31. Ionic Radii
Cations (+) have lost one or more electrons and are smaller than the atoms from which they were derived.
Anions (-) have gained one or more electrons and are larger than the atoms from which they were derived.

32. Anions (negative ions) are larger than their respective atoms.
WHY? Electron-electron repulsion forces spread them further apart. Electrons outnumber protons; the protons cannot pull the extra electrons as tightly toward the nucleus.

33. Cations (positive ions) are smaller than their respective atoms.
WHY? Protons outnumber electrons; the protons can pull the fewer electrons toward the nucleus more tightly. If the electron that is lost is the only valence electron so that the electron configuration of the cation is like that of a noble gas, then an entire energy level is lost.  In this case, the radius of the cation is much smaller than its respective atom.

34. Metallic Character Is associated with:
Larger atomic radii
Lower ionization energies
Lower electron affinities
Lower electronegativities
The most active metals (ones with the most metallic character) are located in the lower left corner of the table.

35. Nonmetallic Character
Is associated with:
Smaller atomic radii
Higher ionization energies
Higher electron affinities
Higher electronegativities
The most active nonmetals are found in the upper right hand corner of the periodic table (excluding the Noble gases).

36. What About the Noble Gases?
do not behave as metals or nonmetals.
very high ionization energies and positive electron affinity values.
Noble gases usually are not assigned electronegativities due to their tendency to not form chemical bonds.
Have full valence electron shells so they do not tend to form compounds at all.

37. Chemical Reactivity F Fr Metals increase in reactivity left and down.
Nonmetals become more reactive up and to the right.
Most reactive metal is?
Most reactive nonmetal is? Fr F