1. The Periodic Table and Trends Topics 3.1 - 3.3 and 12.1.1 - 12.1.2 SONG Please have a periodic table out.

2. IB prefers this one.

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 now atomic #) he even left empty spaces to be filled in later

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 all elements have 1,2,3,4,5,6,7, or 8 valence electrons

11. Lewis Dot-Diagrams/Structures valence electrons are represented as dots around the chemical symbol for the element Na Cl

12. 3.1.3 Electron arrangement and position on the table 3.1.4 # of electrons in the highest occupied energy level and position on the table

15. Look, they are following my rule!

16. 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?

17. 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

18. Definitions 3.2.1 atomic radii –the distance from the nucleus to the outermost electron ionic radii –same measurement as atomic radii but have lost or gained valence electrons first ionization energy (kJ mol -1 ) –the energy needed to remove the outermost/valence electrons/highest energy electron from a neutral atom in the gaseous phase, leaving behind a gaseous ion X(g)  X + (g) + e - Na(g)  Na + (g) + e -

19. don’t forget-- gaseous

20. electronegativity –measures the attraction for electrons in a bond Linus Pauling (1901 to 1994) came up with a scale where a value of 4.0 is arbitrarily given to the most electronegative element, fluorine, and the other electronegativities are scaled relative to this value. melting point chemical properties –how elements react with other elements

21. Trends in the table (3.2.2- 3.2.4) IB loves the alkali metals, the halogens, and period 3 (look at the assessment statements)

22. many trends are easier to understand if you comprehend the following the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces –the attraction between the electron and the nucleus –the repulsions between the electron in question and all the other electrons in the atom (often referred to the shielding effect) –the net resulting force of these two is referred to effective nuclear charge

25. A TOMIC RADII

26. –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 in the given level closer to it »added electrons are not contributing to the shielding effect H Li Na K Rb

30. –I ONIC RADII –across period 3 decreases at first when losing electrons (+ ion) then suddenly increases when gaining electrons (- ion) –then goes back to decreasing after that

31. DARK GREY IS THE SIZE OF THE ION

32. This is period 2 but same trend as period 3

33. –alkali metals cations (+ ions) are smaller than 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 Li+.078nm + Li forming a cation

34. –halogens (group 7) anions (- ions) are larger than parent atom –have gained an electron to achieve noble gas configuration –effective nuclear charge has decreased since same nuclues now holding on to more electrons 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

35. –I ONIZATION ENERGY recall, the energy to remove an electron –example: Na(g)  Na + (g) + e - 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 and therefore easier to remove

37. increases across a period –extra electrons are just filling up the same level (or shell, same thing) –the nucleus is becoming stronger and therefore the electrostatic force increases making the atom smaller this makes it harder to remove a valence electron since it is closer to the nucleus –or another way to look at it… a higher nuclear charge acting on more contracted orbitals

39. More on ionization energy-- Topic 12.1.1 and 12.1.2 Evidence for levels and sub-levels –using the F IRST IONIZATION energy for elements electrons are harder to remove… –when there are more protons to attract the valence electrons in a given level (higher nuclear charge) –when the atomic radius decreases and the electrons are closer to the nucleus –a sub-level (s,p,d,f) is completely filled –a sub-level (s,p,d,f) are half filled

40. only last sub- level

42. Evidence for levels and sub-levels –using SUCCESSIVE (1 st, 2 nd, 3 rd,…) ionization energy for an element as more electrons are removed, the effective nuclear charge increases, pulls the electrons in closer and holds them tighter –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 »less electrons are shielding every time you remove a level Topic 12.1.2

45. 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 K

47. –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 Back down to Topic 3 H Li Na K Rb

49. as you go across a period –electronegativity increases the atoms become smaller as the nuclear charge increases –attracts electrons as they can be closer to the nucleus moving from L to R on the table

50. next concept requires understanding of concepts covered in later topics –what a metal is –Van der Waals forces –bonding

51. –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

53. increases

54. 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

55. C HEMICAL PROPERTIES (3.3.1 AND 3.3.2) –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 attracts electrons better since they can be close to the nucleus 1+ charge 1- charge

56. most reactive least reactive

57. halogens (group 7) –diatomic molecules such as Cl 2, Br 2, I 2 »can react with halide ions (Cl -, Br -, and I - ) »the most reactive ends up as an ion (1 - charge) and is not visible (molecules 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

58. –periods from left to right in period 3 –metals…metaloids…nonmetals –oxides are »ionic…..and then covalent bonds –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

59. Just look for the on the next slide It is basically 3.3.2