Periodic Table History and Trends

1. X Unit 10:The Periodic Table

2. History of the Periodic Table Antoine Lavoisier (1743 – 1794) • Published Elements of Chemistry in 1789 • Included a list of “simple substances” (which we now know to be elements) • Formed the basis for the modern list of elements • Only classified substances as metals or nonmetals

3. History of the Periodic Table Johann Döbereiner (1780 – 1849) • Classified elements into “triads” • Groups of three elements with related properties and weights • Began in 1817 when he realized Sr was halfway between the weights of Ca and Ba and they all possessed similar traits • Döbereiner’s triads: • Cl, Br, I  S, Se, Te • Ca, Sr, Ba  Li, Na, K

4. History of the Periodic Table John Newlands (1837 – 1898) • Law of Octaves (1863) • Stated that elements repeated their chemical properties every eighth element • Similar to the idea of octaves in music

5. History of the Periodic Table Dmitri Mendeleev (1834 – 1907) • Russian chemist (“The father of the periodic table”) • Arranged elements based on accepted atomic masses and properties that he observed • Listed elements with similar characteristics in the same family/group • Left blank spots for predicted elements (Ted-Ed Video)

6. Dmitri Mendeleev (1834 – 1907)

7. Dmitri Mendeleev (1834 – 1907)

8. History of the Periodic Table Henry Moseley (1887 – 1915) • English physicist • Arranged elements based on increasing atomic number • Remember: atomic number = # of p+ in nucleus • Periodic table looked similar to Mendeleev’s design since as atomic number increases, so does the atomic mass

9. Periodic Law • Periodic – occurring at regular intervals • Relates to trends on the periodic table of elements • Modern Periodic Law • When elements are arranged in order of increasing atomic number, there is a periodic repetition of their properties • Just like Mendeleev suspected!!

10. Reading the Periodic Table • Periods - “Horizontal Rows” • Groups (or Families) - “Vertical Columns”

11. Reading the Periodic Table • Valence electrons are periodic! • Notice the similarities • Ex.) Write the noble gas configurations for: • F [He]2s22p5 7 valence electrons • Cl [Ne]3s23p5 7 valence electrons • Br [Ar]4s23d104p5 7 valence electrons • I [Kr]5s24d105p5 7 valence electrons • GROUPS have similar valence electron configurations!

12. Groups of Elements • Group 1 = Alkali Metals • Located in Group 1 (except Hydrogen) • Extremely reactive • Want to lose 1 e- to become “noble gas-like” • Group 2 = Alkaline Earth Metals • Also very reactive • Both Group 1 & 2 occur naturally as compounds not elements

13. Groups of Elements • Group 17 = Halogens • Very active nonmetals • Want to gain 1 e- to become like a noble gas

14. Groups of Elements • Group 18 = Noble Gases • Sometimes called “inert gases” since they generally don’t react • Mainly true, but not always (Kr, Xe will react sometimes) • Have a full valence shell (8 e-) Mythbusters Noble Gas Demo

15. Groups of Elements • Transition Metals • Located in the center of the Periodic Table • 10 elements wide (“d” orbitals) • Semi-reactive, valuable, crucial to many life processes • Lanthanides and Actinides • Located at the bottom of the Periodic Table • 14 elements wide (“f” orbitals) • Some are radioactive, though not all • Lanthanides = Period 6 (4f) • Actinides = Period 7 (5f)

16. Alkali Metals = Alkaline Earth Metals = Transition metals = Metalloids = Lanthanides = Halogens = Actinides = Noble Gases =

17. Periodic Properties & Trends • Electronegativity • Ability of an atom to pull e- towards itself • Increases going up and to the right • Across a period  more protons in nucleus = more positive charge to pull electrons closer • Down a group more electrons to hold onto = element can’t pull e-as closely

18. Periodic Properties & Trends • Atomic Radius • Distance between the nucleus and the furthest electron in the valence shell • Increases going down and to the left • Down a group more e- = larger radius • Across a period  elements on the right can pull e- closer to the nucleus (more electronegative) = smaller radius • *Remember* • LLLL  Lower, Left, Large, Loose

19. Periodic Properties & Trends • Ionization Energy • Energy required to remove an e- from the ground state • 1st I.E. = removing 1 e-, easiest • 2nd I.E. = removing 2 e-, more difficult • 3rd I.E. = removing 3 e-, even more difficult • Ex.) B --> B+ + e- I.E. = 801 kJ/mol • Ex.) B+ --> B+2 + e- I.E.2 = 2427 kJ/mol • Ex.) B+2 --> B+3 + e- I.E.3 = 3660 kJ/mol

20. Periodic Properties & Trends Ionization Energy • Increases going up and to the right • Down a group more e- for the nucleus to keep track of = easier to rip an e- off • Across a period elements on the right can hold electrons closer (more electronegative) = harder to rip an e- off

21. Periodic Properties & Trends • Metallic Character • How “metal-like” an element is • Metals lose e- • Most Metallic: Cs, Fr • Least: F, O • Increases going down and to the left Think about where the metals & nonmetals are located on the periodic table to help you remember!

22. Periodic Properties & Trends • Ionic Radius • Radius of an atom when e- are lost or gained different from atomic radius • Ionic Radius of Cations • Decreases when e- are removed • Ionic Radius of Anions • Increases when e- are added

23. + + Li , 78 pm 2e and 3 p Sizes of Ions • CATIONS are SMALLER than the atoms from which they are formed. • Size decreases due to increasing he electron/proton attraction. Li,152 pm 3e and 3p

24. - - F, 71 pm F , 133 pm 9e and 9p 10 e and 9 p Sizes of Ions • ANIONS are LARGER than the atoms from which they are formed. • Size increases due to more electrons in shell.

25. Overall Periodic Trends

26. Practice: Rank the elements from lowest to highest… Electronegativity - C, F, Mg Atomic Radius - Ir, Re, Bi Metallic Character - Rb, Mn, P Ionization Energy - B, Ga, In

27. Summary of Periodic Trends