Evolution of the Periodic Table: Elemental Arrangements & Properties

1. Chapter 5 The Periodic Law

2. Sect. 5-1: History of the Periodic Table • Stanislao Cannizzaro (1860) proposed method for measuring atomic mass at First International Congress of Chemists • Dmitri Mendeleev (1869) arranged elements by atomic mass & similar chemical properties; left blanks for undiscovered elements

3. Henry Moseley (1911) arranged periodic table by atomic number • Periodic law – properties of elements are periodic functions of their atomic #’s • Noble gas group, lanthanide and actinide series added later

4. Sect. 5-2: Electron Configuration and the Periodic Table • S-block elements are highly reactive metals because they easily give up their 1 or 2 valence electrons • Group 1 – Alkali metals • Silvery, can be cut with knife • Group 2 – Alkaline Earth metals • Harder, denser, stronger, and slightly less reactive than group 1

5. Special cases: • Hydrogen grouped with 1 because of electron configuration, but doesn’t share their properties • Helium is grouped with 18 because it has similar properties since its outside energy level is full, even though it has the same electron configuration as group 2

6. D-block elements • Total # electrons in d plus electrons in highest s orbital = group # • Referred to as transition elements • Good conductors of heat/electricity • High luster • Not as reactive as s-block elements

7. P-block elements • Combined with s-block they are called main-group elements • Contains metals, non-metals, and metalloids, thus wide range of properties • Group 17 – halogens • Most reactive nonmetals • Group 18 – noble gases • nonreactive

8. F-block elements • Lanthanides • Shiny, similar in reactivity to group 2 • Actinides • All radioactive • First 4 have been found naturally, all others are man-made

9. Sect. 5-3: Electron Configuration and Periodic Properties • Atomic Radius – one half the distance between the nuclei of chemically bonded identical atoms • Decreases from left to right across a period due to higher positive charge on right pulling electrons closer • Increases going down a group because of adding energy levels

10. Ion – charged particle • Ionization energy (IE) – energy required to remove an electron from a neutral atom in the gas phase • Increases as you move to the right because those elements will less readily give up an electron • Decrease as you move down a group due to electrons being further away from nucleus and shielded by inside electrons

11. 2nd and 3rd ionization energies refer to removing additional electrons from positively charged ions • 2nd and 3rd Ionization energies have a drastic “jump” if the ion has the electron configuration of a noble gas

12. Electron Affinity – energy change when a neutral atom gains an electron • Reported as a negative # because of loss of energy • Generally decreases as you move down a group • Generally decreases as you move left on a period • Exceptions for half-filled or filled sublevels • Adding additional electrons will always have a positive value (requires energy)

13. Ionic Radii • Cation – positively charged ion (lost electron) • Will decrease radius because of loss of outer energy level • Anion – negatively charged ion (gained electron) • Will increase radius because protons “pulling in” are the same and with extra electrons they repel each other and spread out • Cation & anion radius increases from the right to the left across a period • Cation & anion radii increase down a group

14. Valence Electrons – electrons in outermost energy level (can be gained lost or shared) • For s-block, # valence electrons is equal to group number • For p-block, # valence electrons is equal to group number minus 10

15. Electronegativity – measure of ability of an atom in a compound to attract electrons • Generally decrease as you move to the left of a period • Generally decrease or stay the same moving down a group • Nitrogen, Oxygen, and Halogens are most electronegative

16. Trends for d- and f-blocks • atomic radius trend is same a main group, but with smaller changes • Ionization energy trend is same for period, but increases going down a group • Ion formation – electrons are removed from the s orbital 1st, then the d • Electronegativity trends are same