The Evolution of the Periodic Table: From Mendeleev to Modern Trends

1. ChemistryChapter 5 The Periodic Table

3. Sept 1860, group of chemists met in Germany to review scientific matters & coming to consensus about: • measurement of atomic masses • determining composition of compounds using atomic masses • Cannizzaro presented method for measuring mass & scientists agreed upon the std values for atomic masses

4. in 1869, Dmitri Mendeleev used Cannizzaro’s method for measuring relative masses of atoms in textbook he wrote • published the arrangement of elements in – periodic table

5. only 60 elements known at this time • organized elements acc to properties & new atomic masses on cards • “game of patience” • Mendeleev grouped elements according to incr atomic mass & noticed certain properties appeared at regular intervals- periodic

7. in 1871, Mendeleev predicted properties of elements that weren’t even discovered at that time! • not all elements fit in according to increasing atomic mass: • I & Te Ar & K Co & Ni • Mendeleev couldn’t explain why • other scientist accepted the periodic table & considered him the Father of the Periodic Law

9. in 1913, Henry Moseley discovered patterns w/ x-ray tubes that led to atomic number (ch 3 notes) • he noticed that when he reordered elements on table acc to incr atomic number they fit into their patterns in better way • led to the Modern periodic table • periodic law- phy & chem properties of the elements are periodic functions of the atomic #

10. Getting Acquainted With the Periodic Table

11. Group Properties • valence e- same for all elements in group • group 1: Alkali metals • e- conf: ns1 • group 2: Alkaline Earth metals • e- conf: ns2 • groups 3-12: Transition metals • e- conf: (n-1)d1ns2

12. groups 4-11 deviations occur • sum of outer s & d e- equal to group # • group 13: ns2np1 • group 14: ns2np2 • group 15: ns2np3 • group 16: ns2np4 • group 17: Halogens ns2np5 • group 18: Noble gases ns2np6

13. noble gases have 8 valence e- • have stable octet very stable & unreactive • f-block elements • lanthanides- rare earth metals 1st row • actinides- all radioactive; most synthetic

14. Periodic Properties • phy & chem properties vary in periodic fashion • properties arise from e- configuration • 5 properties:

15. 1. Atomic Radii • ½ distance betw nuclei of identical atoms joined in a molecule • e- occupy large region around nucleus & size atom varies • periodic trends- gradual decrease in radii across periods • due to increasing pos chrg of nucleus (pulled tighter by nucleus)

16. group trends: as go down group, atomic radii increases due to addition of e- to larger orbitals in higher energy levels

18. 2. Ionization Energy • minimum amount of energy required to remove the most loosely bound e- from an isolated gaseous atom to form an ion w/ a +1 charge • if enough energy is supplied, e- can be removed from atoms • ex: 1st IE for Ca is 590 kJ/mol • Ca + 590kJ/mol  Ca+ + e-

19. ionization- process that results in the formation of ion • 2nd IE is 1145kJ • IE2 > IE1 • ALWAYS more difficult to remove additional e- from positive ion • IE measures how tightly e- are bound to atoms

20. low IE indicates ease of e- removal & cation formation • group trends: as atomic radii increases in a group, 1st IE decreases • b/c the valence e- are further from nucleus “shielding effect” • period trends: IE incr from L to R due to increasing nuclear charge which holds e- tighter

21. nonmetals tend to have higher IE than metals

22. 3. E- affinity • amount of energy involved in the process in which an e- is added to an isolated gaseous atom to produce an ion w/ a -1 charge • many atoms readily add e- & release energy • ex: Cl + e-  Cl- + energy (exothermic) • Why?

23. some have to be forced to gain e- by the addition of energy • Be + e- + energy  Be- • period trends: group 17 elements gain e- most easily ( large neg values) reason for the reactivity of these halogens

25. exceptions are seen betw groups 14 & 15 b/c ½ filled sublevels are a little more stable than ones not ½ full • group trends: generally more difficult to add e- to larger atoms than to smaller atoms • elements w/ very negative EA gain e- readily to form anions (ions w/ negchrg)

26. more difficult to add e- to an anion so 2nd EA are all positive • cation- positive ion • anion- negative ion

27. 4. Ionic radii • ½ the diameter of an ion in a chemical compound • formation of a cation leads to a decrease in radius due to the e- cloud being drawn inward as valence e- are removed • formation of anion leads to an increase in radius as additional e- repel one another

29. periodic trends- metals form cations • nonmetals form anions • group trends- IR increases down group Why? • as you add higher energy levels, radius of ion incr

30. chemical compounds form b/c e- are lost, gained, or shared to bring an atom to a stable octet

31. 5. Electronegativity EN • measure of the power of an atom in a chemical compound to attract e- • valence e- hold atoms in compound together & properties of compound are influenced by conc of negchrg closer to one atom than another • ex: NaCl

33. numerical values assigned to indicate the tendency of atom to attract e- • Fluorine – most EN element & assigned value of 4 • periodic trends- gradual incr in EN from L to R across period • nonmetals tend to be more EN than metals

34. groups 1 & 2 least EN elements • halogens are most EN elements • group trends- EN either decreases down group or remains similar