1. The Periodic Table Chapter 6

2. Organizing the Elements  Demitri Mendeleeve (Russian – 1869)  Published the 1 st periodic table  Based on atomic mass  He left blanks for yet undiscovered elements  he had made good predictions  But, there were problems:  Such as Co and Ni; Ar and K; Te and I  Henry Moseley – British physicist  Arranged elements according to increasing atomic number  Symbol, atomic number & mass

3. Periodic Table Periodic Law: Elements are arranged in the order if increasing atomic number, there is a periodic repetition of their physical and chemical properties. Periods = rows 1-7 ; equal to the principal energy level Groups = columns 1-18 ; elements have similar properties Period → Group ↓

4. Metallic Character Approx. 2/3 of the elements are metals Metallic Properties  Luster  Conductivity  Malleability  Ductility  Tend to lose electrons Nonmetallic Properties  Dull luster  Poor conductors  Brittle Tend to gain electrons Metalloids (semimetals)  Some properties characteristic of metals and other properties characteristic of nonmetals

5. Metals and Nonmetals Metallic character decreases → Metallic character increases ↓

6. Element on the periodic table

7. What can you learn about each element from the periodic table? You can learn an element’s name, its symbol, its atomic number, its atomic mass, and the number of electrons in each energy level (its electron configuration). Element on the periodic table

8. Classifying the Elements Alkali Metals – Alkaline Earth Metals – Halogens – Noble Gases – Representative Elements – Transition Elements – Group 1 (1A) Group 2 (2A) Group 17 (7A) Group 18 (8A) Inert gases Groups 1A – 7A All Group B

9. Electron Configurations in Groups The Representative Elements This figure shows a portion of the periodic table containing Groups 1A through 7A. Elements in Groups 1A through 7A are often referred to as representative elements because they display a wide range of physical and chemical properties.

10. The Representative Elements Some elements in these groups are metals, some are nonmetals, and some are metalloids. Most of them are solids, but a few are gases at room temperature, and one, bromine, is a liquid.

11. In atoms of representative elements, the s and p sublevels of the highest occupied energy level are not filled. The Representative Elements Lithium (Li)1s22s11s22s1 Sodium(Na)1s22s22p63s11s22s22p63s1 Potassium(K)1s22s22p63s23p64s11s22s22p63s23p64s1 In atoms of these Group 1 A elements, there is only one electron in the highest occupied energy level. –The electron is in an s sublevel.

12. The Representative Elements Carbon (C)1s22s22p21s22s22p2 Silicon (Si)1s22s22p63s23p21s22s22p63s23p2 Germanium(Ge)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 For any representative element, its group number equals the number of electrons in the highest occupied energy level. In atoms of the Group 4A elements carbon, silicon, and germanium, there are four electrons in the highest occupied energy level.

13. The s and p sublevels are completely filled with electrons — two electrons in the s sublevel and six electrons in the p sublevel. The Noble Gases Look at the description of the highest occupied energy level for each element, which is highlighted in yellow. Helium (He)1s21s2 Neon (Ne)1s22s22p61s22s22p6 Argon (Ar)1s22s22p63s23p61s22s22p63s23p6 Krypton (Kr)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

14. Electron Configurations in Groups Transition Elements In the periodic table, the B group elements separate the A groups on the left side of the table from the A groups on the right side. Elements in the B groups are referred to as transition elements. There are two types of transition elements—transition metals and inner transition metals. –They are classified based on their electron configurations.

15. Transition Elements The transition metals are the Group B elements that are usually displayed in the main body of a periodic table. Copper, silver, gold, and iron are transition metals. In atoms of a transition metal, the highest occupied s sublevel and a nearby d sublevel contain electrons. These elements are characterized by the presence of electrons in d orbitals.

16. Transition Elements The inner transition metals are the elements that appear below the main body of the periodic table. In atoms of these elements, the highest occupied s sublevel and a nearby f sublevel generally contain electrons. The inner transition metals are characterized by the presence of electrons in f orbitals.

17. Periodic Table e - configurations are inherent in the periodic table B 2p 1 H 1s 1 Li 2s 1 Na 3s 1 K 4s 1 Rb 5s 1 Cs 6s 1 Fr 7s 1 Be 2s 2 Mg 3s 2 Ca 4s 2 Sr 5s 2 Ba 6s 2 Ra 7s 2 Sc 3d 1 Ti 3d 2 V 3d 3 Cr 4s 1 3d 5 Mn 3d 5 Fe 3d 6 Co 3d 7 Ni 3d 8 Zn 3d 10 Cu 4s 1 3d 10 B 2p 1 C 2p 2 N 2p 3 O 2p 4 F 2p 5 Ne 2p 6 He 1s 2 Al 3p 1 Ga 4p 1 In 5p 1 Tl 6p 1 Si 3p 2 Ge 4p 2 Sn 5p 2 Pb 6p 2 P 3p 3 As 4p 3 Sb 5p 3 Bi 6p 3 S 3p 4 Se 4p 4 Te 5p 4 Po 6p 4 Cl 3p 5 Be 4p 5 I 5p 5 At 6p 5 Ar 3p 6 Kr 4p 6 Xe 5p 6 Rn 6p 6 Y 4d 1 La 5d 1 Ac 6d 1 Cd 4d 10 Hg 5d 10 Ag 5s 1 4d 10 Au 6s 1 5d 10 Zr 4d 2 Hf 5d 2 Rf 6d 2 Nb 4d 3 Ta 5d 3 Db 6d 3 Mo 5s 1 4d 5 W 6s 1 5d 5 Sg 7s 1 6d 5 Tc 4d 5 Re 5d 5 Bh 6d 5 Ru 4d 6 Os 5d 6 Hs 6d 6 Rh 4d 7 Ir 5d 7 Mt 6d 7 Ni 4d 8 Ni 5d 8

18. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.. Electron Configurations in Groups Blocks of Elements If you consider both the electron configurations and the positions of the elements in the periodic table, another pattern emerges.

19. What does the period an element is located in tell you about that element’s electron configuration? What does the group tell you? The period tells you the highest occupied principal energy level. The group number is equal to the number of electrons in the highest occupied energy level.

20. Some compounds are composed of particles called ions. An ion is an atom or group of atoms that has a positive or negative charge. Ions

21. An atom is electronically neutral because it has equal numbers of protons and electrons.  For example, an atom of sodium (Na) has 11 positively charged protons and 11 negatively charged electrons.  The net charge on a sodium atom is zero [(+11) + (-11) = 0]. Ions

22.  Ion is an electrically charged atom  Atoms may either gain or lose electrons  Loss of electrons = (+) charge → CATION (Metals)  Gain of electrons = (-) charge → ANION (Nonmetals)

23. AnionCation

24. What type of element tends to form anions? What type tends to form cations? Nonmetals tend to form anions. Metals tend to form cations.

25. Atomic Radius Atoms are roughly spherical Sphere size can be determined using the radius Problem: Edges of orbitals are fuzzy and difficult to measuer.28 Solution: Atomic Radius – half the distance between the nuclei of identical atoms that are bonded together.

26. Periodic trend of Atomic Radius

27. Ionization energy Ionization energy: is the energy required to remove an electron from an atom (in a gaseous state). Trends: Across a period, the ionization energy → increases Why? : Electrons are held closer to the nucleus and Electrons are more difficult to remove Down a group, the ionization energy → decreases Why? : Electrons are further away from the nucleus and Electrons are easier to remove A + energy  A + + e -

28. Periodic trend of ionization energy Increases  decreases ↓

29. Trends in Ionic Size Cations form when atoms loose e -. The cation is smaller. Anions form when atoms gain e -. The anion is larger.  Trends: 1. Across a period, ionic size → decreases Why? : When e - are lost, the attraction between the remaining e - increases, and e - are drawn closer to the nucleus 2. Down a group, ionic size → increase Why? : Increased atomic radius, and e - are further away from the nucleus

31. Electronegativity Electronegativity is a measure of the ability of an atom to attract a pair of electrons when bonded to another atom. Ability to ‘pull’ electrons from another atom  Trends: 1. Across a period, the electronegativity → increases Why? : Atoms are very close to a stable octet, and ‘pull’ electrons from other atoms 2. Down a group, the electronegativity → decreases Why? : Top atoms have few electrons, and hold them tightly. Bottom atoms have numerous electrons and little desire to acquire more.

32. Periodic trend of Electronegativity