1. The Periodic Law History of the Periodic Table Electron Configuration & the Periodic Table Electron Configuration & Periodic Properties

2. History of the Periodic Table

3. Mendeleev  Dmitri Mendeleev (1869, Russian) Organized elements by increasing atomic mass. Elements with similar properties were grouped together. There were some discrepancies.

4. Mendeleev  Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered elements.

5. Moseley  Henry Moseley (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement. Periodic Law-the physical and chemical properties of the elements are periodic functions of the atomic numbers.

6. Organization of the Elements  Periodic table-an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

7. Additions to Mendeleev’s Periodic Table  Noble gases Group 18 Argon discovered in 1894 Took so long to discover because very unreactive  Lanthanides 14 elements with atomic numbers from 58-71 Placed below the periodic table to conserve space  Actinides 14 elements with atomic numbers 90-103 Also placed below periodic table

8. Electron Configuration & the Periodic Table

9. Periods & Blocks of the Periodic Table  Length of period (row) determined by how many electrons can occupy the sublevels being filled. 1 st period-1s sublevel being filled with 2 electrons  2 elements, H & He 3 rd period-3s & 3 p sublevels being filled with 2+6 electrons  8 elements  Periodic table is divided into “blocks” based on the filling of sublevels with electrons.

10. Blocks of the Periodic Table

11. Determining Period from Configuration  An element’s period can be determined by looking at its electron configuration  The highest occupied energy level corresponds to the element’s period As: [Ar]3d 10 4s 2 4p 3  4 in 4p 3 indicates that the highest energy level that electrons occupy is the 4 th. Therefore, As is located in the 4 th period of the periodic table.

12.  Metals  Nonmetals  Metalloids Metallic Character

13.  Main Group Elements  Transition Metals  Inner Transition Metals Areas of the Periodic Table

14. s-Block Elements: Groups 1 & 2  Chemically reactive metals  Include the alkali metals and the alkaline earth metals

15. Alkali metals  Group 1 metals  ns 1  Silvery appearance and very soft  Not found pure naturally because so reactive  Because of extreme reactivity with moisture, usually stored under kerosene  Video: Disposal of Surplus SodiumDisposal of Surplus Sodium  Video: Alkali Metals in WaterAlkali Metals in Water

16. Alkaline-Earth metals  Group 2 metals  ns 2  Harder, denser, & stronger than alkali metals  Also too reactive to be found free in nature (but less reactive than Gp. 1)  Video: Magnesium/silver nitrate mixture reacting with waterMagnesium/silver nitrate mixture reacting with water

17. d-Block Elements: Groups 3-12  Metals with typical metallic properties  Called “transition elements”  Typically less reactive than Gps. 1&2, & some are extremely unreactive  d sublevels first appears at the 3 rd energy level  Fills after 4s  Variations from expected in d-block, so elements in the same group do not necessarily have the same outer e- configuration

18. p-Block Elements: Groups 13-18  p and s-block elements together called “main-group elements”  Total number of electrons in highest energy level=group # - 10 Group 17 elements have 17-10=7 outer “valence” electrons  Properties of p-block elements vary greatly since metals, nonmetals, and metalloids are contained here

19. p-block Elements  Halogens Group 17 nonmetals Most reactive nonmetals  React with most metals to form salts  Metalloids Fall on both sides of a “stair-step” line separating metals and nonmetals Semi-conductors

20. f-Block Elements: Lanthanides & Actinides  Lanthanides Top row of f-block 14 elements Shiny metals similar in reactivity to the alkaline-earth metals  Actinides Bottom row of f-block 14 elements All radioactive 1 st 4 elements found naturally on Earth; remainder only lab-made elements

21. Electron Configuration & Periodic Properties

22. AP: 12/4  A blank piece of paper, the periodic table book, and the periodic trends paper from Wednesday.  Turn in your homework over the trends straw activity performed in class Wednesday  We will discuss trends today (Study over the weekend. Remember how much AP loves this topic)  Your test will be Tuesday over trends  Monday we will grade your free response question from Thanksgiving break based on point values assigned by AP

23. On the clean piece of paper that you picked up, draw the following… TrendDefinitionWhy trend for group Why trend for period/row Exceptions Explained Atomic Radius N/A Ionization Energy (I.E.) *Group 2 to 13 *Group 15 to 16 Electron Affinity To many to remember Electronegativity Noble gases tend to NOT follow this trend. Why? Ionic Radius: Cation/Anion NONE

24. Trends Tricks:  “A trend is an observation NOT an explanation”. Say it over and over until it sticks in your head.  Talk about BOTH atoms involved on the AP exam or you will lose credit.  Mention Coulombic attraction where it pertains to a trend.  Emphasize energy  Use effective nuclear charge (Z eff ) when in a period/row AND use size, distance, and SHIELDING when in a group as an explanation

25. Remember the Periodic Law  When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

26.  ½ the distance between the nuclei of identical atoms that are bonded together  Increases to the LEFT and DOWN Atomic Radius

27.  Why bigger across a period/row: Effective nuclear charge decrease attraction of nucleus and therefore the pull of the nucleus results in larger radius  Why larger down a group: Increase number of energy levels down a group, increase distance over which nucleus must pull electrons and therefore because of shielding of inner core, reduces attraction for electrons

28. Atomic Radius Li Ar Ne K Na

29.  Why larger going down? Higher energy levels have larger orbitals Shielding - core e - block the attraction between the nucleus and the valence e -  Why smaller to the right? Increased nuclear charge without additional shielding pulls e - in tighter Atomic Radius

30.  Why bigger across a period/row: Effective nuclear charge decrease attraction of nucleus and therefore the pull of the nucleus results in larger radius  Why larger down a group: Increase number of energy levels down a group, increase distance over which nucleus must pull electrons and therefore because of shielding of inner core, reduces attraction for electrons

31.  First Ionization Energy-energy required to remove one electron from a neutral atom  Increases UP and to the RIGHT Ionization Energy

32.  First Ionization Energy Ionization Energy K Na Li Ar Ne He

33.  Why more energy across a period? Effective nuclear charge increases the attraction of the nucleus therefore holds electrons more tightly (WITH EXCEPTIONS from GROUP 2 to 13 AND from GROUP 15 to 16)  Why more energy up a group? Decreased number of energy levels, decreases the distance over which nucleus must pull and therefore reduces attraction for electrons Full energy levels provide some shielding between nucleus and valence electrons so closer to the top of group, less energy levels and less shielding Ionization Energy

34. Ionization Energy~EXCEPTIONS  Exceptions to these general trends occur at group 3 (shielding of p electron by s electrons) and group 6 (paired electron repulsion).

35. Activity: Ionization Energy  1. Draw the orbital notation for an element in group 3 and group 16  2. Highlight the electrons that cause the exception to this trend  3. Explain the exception to the trend from 2 to 13  4. Explain the exception to the trend from 15 to 16  5. Why does rubidium have a smaller I.E. than Na  6. Justify chlorine having a higher I.E. than potassium  7. Which group: Ist I.E. 352 kJ/mol 2 nd I.E. 658 kJ/mol 3 rd I.E. 8005 kJ/mol

36.  Successive Ionization Energies Mg1st I.E.736 kJ 2nd I.E.1,445 kJ Core e - 3rd I.E.7,730 kJ Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

37. Al1st I.E.577 kJ 2nd I.E.1,815 kJ 3rd I.E.2,740 kJ Core e - 4th I.E.11,600 kJ  Successive Ionization Energies Large jump in I.E. occurs when a CORE e - is removed. Ionization Energy

38.  Energy change that occurs when an electron is acquired by a neutral atom (addition of electron in gaseous atom or ion).  Tends to become less negative (less energy released) DOWN and to the LEFT Electron Affinity

39. Electron Affinity Explained:  Why down a group: Change little moving down a group.  Why more negative across a row/period towards noble gases: Become increasingly negative from left to right. More positive, the less attractive to electrons around them. Careful: addition or subtraction can be exothermic (-) or endothermic (+). As you more toward the noble gases, the affinities become more negative. Trend explained because of octet rule. (atoms close to full valence will tend to gain electrons and have very negative affinities (give off great deal energy when gaining electrons) NOBLE GASES DO NOT CONFORM TO THIS. They have very positive values.

40.  Ionic Radius (WHY) Cations (+)  lose e -  smaller © 2002 Prentice-Hall, Inc. Anions (–)  gain e -  larger Ionic Radius

41. ISOELECTRONIC SERIES  An atom or ion containing the same amount of electrons  Ex: Ar, Ca 2+, S 2-, K +, Cl -  Practice questions: List all atoms and common ions of representative (main group) elements that are isoelectronic with the nitrogen ion. Which is the smallest?

42. Electronegativity  A measure of the ability of an atom in a chemical compound to attract electrons  Most electronegative element is fluorine Given arbitrary value of 4; all others relative

43. Electronegativity Explained  Why higher up a group: Decrease number of energy levels, distance over which the nucleus must pull and therefore increases attraction for electron. Full energy levels AT THE bottom of a group provides shielding for valence electrons  Why higher up a group: Effect nuclear charge increases attraction of the nucleus increases and therefore strengthens the attraction of the electrons

44.  Which atom has the larger radius? BeorBa CaorBr Ba Ca Examples

45.  Which atom has the higher 1st I.E.? NorBi BaorNe N Ne Examples

46.  Which has the greater electonegativity? KorLi AlorCl Li Cl Examples

47.  Which particle has the larger radius? SorS 2- AlorAl 3+ S 2- Al Examples

49. Add to review:  Define isoelectronic and give a set of examples.  Place the following in order of increasing first ionization energy: Na, Al, Sn, Ca  Place the following in order of increasing atomic radii: Na, Mg, Rb, K  Why does the He atom have a smaller radius than the H atom? Why is the He atom smaller than the Ne atom. EXPLAIN.  List all atoms and common ions of representative elements that are isoelectronic with the aluminum ion. Which is the smallest?  Why does Mg atom require a larger amount of ionization energy than the Ba atom? Why does the Cl atom require a larger amount of ionization energy than the Mg atom. EXPLAIN.