The Periodic Table

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

16. Alkali metals

17. Alkaline Earth Metals (Group 2) • Form metal hydroxides (strong bases) when reacting in water Ca + 2 HOH  Ca(OH)2 + 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 • 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