1. Chapter 11 The Periodic Table

2. I. History of the Periodic Table Johann Wolfgang Döbereiner and triads John Newlands and the Law of Octaves Dmitri Mendeleev and the 1 st periodic table

3. Mendeleev’s Periodic Table

4. Mendeleev’s Predictions Predicted properties for Mendeleev’s Eka-Silicon and properties of Germanium: ElementAtomic WeightDensity Oxide formulaChloride formula Eka-Silicon (predicted 1871) 72 5.5 g/cm 3 EsO 2 EsCl 4 Germanium (discovered 1886) 72.595.32 g/cm 3 GeO 2 GeCl 4

5. Periodic Law Basis: Element arranged according to their atomic masses present a clear periodicity of properties Modern: The properties of elements repeat periodically when the elements are arranged in increasing order by their atomic numbers

6. Regions of the Periodic table

7. Circular Periodic Table

8. Benfey’s Periodic Table

9. Physics Periodic Table

10. Representative Elements -EC Valence v. core electrons

11. Representative Elements - Ions Generalization of atom/ion stability – Usually means 8 valence = octet rule

12. Transition Elements - EC Remember the exceptions to filling d orbitals

13. Periodic Trends – Atomic Radii Worksheet: Atomic Size Why does atomic radius decrease across a period? – Higher # = more protons = higher core charge Increased attraction between p+ & e- – e- pulled closer to nucleus = ???? Why does atomic radius increase down a group? – Valence electron shell  higher n = higher probability of finding e- further from nucleus = ???? – Shielding by core e- = less pull on valence e- = ???? Smaller radius Larger radius

14. Atomic Radii

15. Periodic Trends – Ionic Radii Cation (+) radii are smaller than atomic radii – Why? Lose of valence e- Results in lower n, resulting in stronger nuclear pull Anion (-) radii are larger than atomic radii – Why? Gain of e- Results in increased repulsion between e-

16. Sizes of Anions (- ions)

17. Sizes of Cations (+ ions)

18. Graph of Atomic Radii

19. Definition of Ionization Energy (IE) Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first or initial ionization energy or E i of an atom or molecule is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions. You may think of ionization energy as a measure of the difficulty of removing electron or the strength by which an electron is bound. The higher the ionization energy, the more difficult it is to remove an electron. Therefore, ionization energy is an indicator of reactivity.

20. Periodic trends – 1 st Ionization Energy exceptions

22. Periodic Trends – 2 nd Ionization Energy 12345678 H1312 He23725250 Li520729711810 Be89917571484521000 B800242636592502032820 C10862352461962213782047260 N140228554576747394425325064340 O131433885296746710987133207132084070 F168033756045840811020151601786092010 Ne20803963613093611218015240 Na49645636913954113350166002011325666 Mg737145077311054513627179952170025662 *The teal colored cells represent ionization energies where the valence shell is now (n-1). (Why do you think there is such a large jump in the ionization energies when the n-1 shell is now valence?)

23. Periodic Trends - Electronegativity Definition Increases across a period (L to R), decreases down a group (top to bottom)

24. Electronegativity Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to Cesium and Francium which are the least electronegative at 0.7.

27. Electron Affinity!

28. Electron Affinity Definition the quantitative measure, usually given in electron-volts (eV), of the tendency of an atom or molecule to capture an electron and to form a negative ion.

30. Periodic Trend for electron affinity

31. Periodic Trends - All *Note: The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion

32. Wrap-up Periodic Table Activity