1. The Periodic Table and Physical Properties SONG Topics 3.1 - 3.3 and 12.1.1 - 12.1.2 Get a periodic table out.

5. Dmitri Mendeleev 8 February 1834 – 2 February 1907 Russian chemist and teacher given the elements he knew about, he organized a “Periodic Table” based on increasing atomic mass (it’s atomic # now) he even left empty spaces to be filled in later (TOK– he was a “scientist” and “risk taker”!)

6. At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic table. He predicted their discovery and estimated their properties.

9. Design of the Table Groups are the vertical columns. –elements have similar, but not identical, properties most important property is that they have the same # of valence electrons

10. valence electrons- electrons in the highest occupied energy level

11. http://images.google.com/imgres?imgurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/imgper/econfig.gif&imgrefurl=http://hyperphysics.phy- astr.gsu.edu/hbase/pertab/perlewis.html&h=267&w=512&sz=22&tbnid=__EXctBwlG0J:&tbnh=66&tbnw=128&hl=en&start=1&prev=/images%3Fq %3DElectron%2BDot%2BDiagrams%26svnum%3D10%26hl%3Den%26lr%3D Electron arrangement (SL level – 3.1.3)

14. B is 1s 2 2s 2 2p 1 ; –2 is the outermost energy level –it contains 3 valence electrons, 2 in the s and 1 in the p Br is [Ar] 4s 2 3d 10 4p 5 How many valence electrons are present?

15. Periods are the horizontal rows –do NOT have similar properties –however, there is a pattern to their properties as you move across the table that is visible when they react with other elements

16. Definitions atomic radii –the distance from the nucleus to the outermost electron ionic radii –same distance, but for ions (atoms that have lost or gained valence electrons) first ionization energy (kJ mol -1 ) –the energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gaseous phase

17. electronegativity –measures the attraction for a shared pair of electrons melting point chemical properties –how elements react with other elements

18. Trends in the table

19. But first, the electron shielding effect electrons between the nucleus and the valence electrons repel each other

20. A TOMIC RADII –McGraw Hill videoMcGraw Hill video –groups (alkali metals and halogens) increases downwards as more levels are added –periods across the periodic table (period 3) radii decreases –the number of protons in the nucleus increases »increases the strength of the positive nucleus and pulls electrons closer to it H Li Na K Rb

24. I ONIC RADII –decreases across periods for same reason as atomic radii (nucleus becomes stronger) –alkali metals cations are smaller that the parent atom –have lost an electron (actually, lost an entire level) –therefore have fewer electrons than protons radii still increases downwards as more levels are added on Li 0.152 nm 3e and 3p Li +, 0.078 nm 2e and 3 p + forming a cation

25. –halogens anions are larger than parent atom –have gained an electron to achieve noble gas configuration radii still increases downwards as more levels are added on F 0.064 nm 9e - and 9p + F - 0.133 nm 10 e - and 9 p + - forming an anion

28. –I ONIZATION ENERGY decreases down a group –outer electrons are farther from the nucleus and therefore easier to remove –inner core electrons “shield” the valence electrons from the pull of the positive nucleus

30. increases across a period –extra electrons are just filling up the same level –the nucleus is becoming more powerful and therefore the electrostatic force increases making it harder to remove an electron

31. 12.1.1 Evidence for levels and sub-levels –First ionization energy electrons are harder to remove… –when there are more protons to attract them –a sub-level (s,p,d,f) is completely filled –a sub-level (s,p,d,f) are half filled

33. 12.1.1 Evidence for levels and sub-levels –successive (1 st, 2 nd, 3 rd ) ionization energy as more electrons are removed, the electrostatic pull of the protons holds the remaining electrons closer therefore, more energy is required to remove them (even have to use a logarithmic scale to show this) large “jumps” are when the electrons are being removed from the next, lower level that are much closer to the nucleus

35. 4s 1 removed starting to remove the 3p sub-level starting to remove the 2p sub-level starting to remove the 3s sub-level starting to remove the 2s sub-level starting to remove the 1s sub-level

36. –E LECTRONEGATIVITY as you go down a group electronegativity decreases –the size of the atom increases »the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+) »the valence electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons

37. as you go across a period –electronegativity increases the atoms become smaller so the positive nucleus can hold onto the electrons better

39. –M ELTING POINT group 1 (alkali metals) –decreases as “sea of negative electrons” are farther away from the positive metal ions group 7 (halogens) –increases downwards as the van der Waals’ forces increase »larger molecules have more electrons which increases the chance that one side of the molecule could be negative Element Melting Point (K) Li453 Na370 K336 Rb312 Cs301 Fr295

41. increases decreases

42. across the table (period 3) –from left to right bonding goes from strong metallic to very strong macromolecules (network covalent) to weak van der Waals’ attraction

43. C HEMICAL PROPERTIES –groups alkali metals –react vigorously with water and air »2Na (s) + H 2 O (l)  2Na (aq) + 2OH- (aq) + H 2 (g) »(Li, Na, K… all the same equation) »reactivity increases downwards »because the outer (valence) electron is in higher energy levels (farther from the nucleus) and easier to remove –react with the halogens »halogens’ reactivity increases upwards »smaller size can attract electrons better »(see next slide) 1+ charge 1- charge

44. most reactive least reactive

45. halogens –diatomic molecules such as F 2, Cl 2, Br 2, I 2 »can react with halide ions (Cl -, Br -, and I - ) »the single bond is broken and each atom can gain one electron to form halide ions (F 1-, Cl 1-, Br 1-, I 1- ) »the most reactive ends up as an ion (1 - charge) and is not visible (molecules F 2, Cl 2, Br 2, I 2 are a visible gas) »Cl > Br > I Cl - (aq)Br - (aq)I - (aq) Cl 2 Colorless- no reaction turns red due to formation of Br 2 turns brown due to formation of I 2 Br 2 no reaction turns brown due to formation of I 2 I2I2 no reaction

46. –periods from left to right in period 3 –metals…metaloids…nonmetals –when oxides react with water »basic…amphoteric (either basic or acidic)…acidic »Na 2 O(s) + H 2 O (l)  2 NaOH (aq) strong base »MgO (s) +H 2 O (l)  Mg(OH) 2 (aq) weaker base »P 4 O 10 (s) + 6H 2 O (l)  4 H 3 PO 4 (aq) weak/strong acid »SO 3 (g) + H 2 O (l)  H 2 SO 4 (aq) strong acid