1. Ch. 14: Chemical Periodicity Standard: Matter consists of atoms that have internal structures that dictate their chemical and physical behavior. Targets: Describe the arrangement of elements in the periodic table in order of increasing atomic number. Distinguish between the terms group and period. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Discuss the similarities and differences in the chemical properties of elements in the same group. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities.

2. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Development of the Periodic Table Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829). John Newlands Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863). Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

3. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Distinguish between the terms group and period. Development of the Periodic Table Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913) Glen Seaborg Discovered the transuranium elements (93-102) and added the actinide and lanthanide series (1945)

4. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Distinguish between the terms group and period. Development of the Periodic Table Elements are arranged by increasing atomic number into periods (rows) and groups or families (columns)

5. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Metals –Left side of the periodic table (except hydrogen). –High electrical conductivity, high luster, ductile, malleable –Alkali metals: Group 1 (1A) –Alkaline earth metals: Group 2 (2A) –Transition metals: Group B, lanthanide & actinide series

6. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Nonmetals –Right side of the periodic table –Poor conductors and nonlustrous –Halogens: Group 17 (7A) –Noble gases: Group 18 (0)

7. Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Metalloids –Between metals and nonmetals –Properties intermediate between metals and nonmetals

8. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg 392-393 Noble Gases: Outermost s and p sublevels are filled. –Ending configuration is s 2 p 6 (except He) –Eight valence electrons (except He) –Row number equals highest energy level

9. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg 392-393 Representative Elements: Outermost s and p sublevels are partially filled. –Group 1,2, 13-17 – 1 (s 1 ); 2 (s 2 ); 13 (s 2 p 1 ); 14 (s 2 p 2 )… –Group number equals valence electrons (1-2, 13-17 subtract 10) –Row number equals highest energy level

10. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg 392-393 Transition Metals –Filling the d & f sublevels

11. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Shortcut Electron Configuration Based on the electron configuration of the noble gases. He ends in 1s 2 ; Ne ends in 2p 6 ; Ar ends in 3p 6 ; Kr ends in 4p 6 ; etc. Write the electron configuration and orbital filling diagram for Se –Se has 34 electrons –Go back to the previous noble gas: Ar (18 electrons). Begin the configuration with [Ar] which accounts for 18 electrons and then begin with 4s 2. Continue until you reach 34 electrons –[Ar]4s 2 3d 10 4p 4 –[Ar] __ __ __ __ __ __ __ __ __ 4s 3d 4p

12. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Shortcut Electron Configuration Write the electron configuration and orbital filling diagram for Au –Au has 79 electrons –Go back to the previous noble gas: Xe (54 electrons). Begin the configuration with [Xe] which accounts for 54 electrons and then begin with 6s 2. Continue until you reach 79 electrons –[Xe]6s 2 4f 14 5d 9 –[Xe] __ __ __ __ __ __ __ __ __ __ __ __ __ 6s 4f 5d

13. Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Shortcut Electron Configuration Electron dot diagrams Group 1A: 1 dotXGroup 5A: 5 dotsX Group 2A: 2 dotsXGroup 6A: 6 dots X Group 3A: 3 dotsXGroup 7A: 7 dotsX Group 4A: 4 dotsXGroup 0: 8 dots (except He)X

14. Discuss the similarities and differences in the chemical properties of elements in the same group. Group 1: Alkali Metals Have 1 valence electron Shiny, silvery, soft metals React with water & halogens Oxidize easily (lose electrons) Reactivity increases down the group

15. Discuss the similarities and differences in the chemical properties of elements in the same group. Group 17: Halogens Have 7 valence electrons Colored gas (F 2, Cl 2 ); liquid (Br 2 ); Solid (I 2 ) Oxidizer (gain electrons) Reactivity decreases down the group

16. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. Atomic Radii Size of the atom Group trend –Atomic size increases as you move down a group of the periodic table. –Adding higher energy levels.

17. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. Atomic Radii Size of the atom Periodic Trend –Atomic size decreases as you move from left to right across a period. –Effect of increasing nuclear charge pulls the electrons closer to the nucleus..

18. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. First Ionization Energies The energy required to remove the first electron from a gaseous atom. Second ionization removes the second electron and so on. Can be used to predict ionic charges. Group trend –Generally decreases as you move down a group in the periodic table –Since size increases down a group, the outermost electron is farther away from the nucleus and is easier to remove.

19. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. First Ionization Energies The energy required to remove the first electron from a gaseous atom. Second ionization removes the second electron and so on. Can be used to predict ionic charges. Periodic Trend –Increases as you move from left to right across a period. –Effect of increasing nuclear charge makes it harder to remove an electron.

20. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. Electronegativity Tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Helps predict the type of bonding (ionic/covalent). Group trend –Generally decreases as you move down a group in the periodic table. –For metals, the lower the number the more reactive. –For nonmetals, the higher the number the more reactive.

21. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities. Electronegativity Tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Helps predict the type of bonding (ionic/covalent). Periodic Trend –Increases as you move from left to right across a period. –Nonmetals have a greater attraction for electrons than metals.