1. The Periodic Table and Trends SONG Please have a periodic table out.

2. Use this one.

5. Dmitri Mendeleev 1834 – 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.

7. Henry Moseley 1887 – 1915 arranged the elements in increasing atomic numbers (Z) –properties now recurred periodically

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

13. What two blocks will always be the highest occupied level?

16. Look, they are following my rule!

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

18. 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

19. Trends in the table IB loves the alkali metals and the halogens

20. 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

21. This is a simple, yet very good picture. Do you understand it?

22. –the distance from the nucleus to the outermost electron –cannot measure the same way as a simple circle due to electrons are not in a fixed location –therefore measure distance between two nuclei and divide by two A TOMIC RADII

24. –groups increases downwards as more levels are added more shielding –periods across the periodic table 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 because they are still in the same level H Li Na K Rb McGraw Hill video

28. –cations (+ ions) are smaller than the parent atom have lost an electron (actually, lost an entire level!) therefore have fewer electrons than protons Li 0.152 nm Li+.078nm + Li forming a cation

29. –anions (- ions) are larger than parent atom have gained an electron(s) to achieve noble gas configuration effective nuclear charge has decreased since same nucleus now holding on to more electrons plus, the added electron repels the existing electrons farther apart (kind of “puffs it out”) F 0.064 nm 9e - and 9p + F - 0.133 nm 10 e - and 9 p + -

30. –trends across a period –decreases at first when losing electrons (+ ion) –then suddenly increases when gaining electrons (- ion) –then goes back to decreasing after just like neutral atoms because of more protons pulling in the outer level down a group (same as neutral atoms) –increases as new levels are added –more levels shielding

31. DARK GREY IS THE SIZE OF THE ION

34. Looking at ions compared to their parent atoms meaning does an atom become smaller or larger as it gains or loses electrons? I ONIC RADII

35. –I ONIZATION ENERGY the minimum energy (kJ mol -1 ) needed to remove an electron from a neutral gaseous atom in its ground state, leaving behind a gaseous ion –X(g)  X + (g) + e - first ionization energy- energy to remove first electron second ionization energy- energy to remove second electron third ionization energy- and so on…

36. don’t forget-- gaseous

37. 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

39. increases across a period –the nucleus is becoming stronger (effective nuclear charge) and therefore the valence electrons are pulled closer atomic radii is decreasing this makes it harder to remove a valence electron since it is closer to the nucleus –or another way to look at it… a stronger nuclear charge acting on more contracted orbitals

41. E LECTRONEGATIVITY –measures the attraction for a shared pair of 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.

43. trends (same as ionization energy and for the same reasons) 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 bonding pair of electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons H Li Na K Rb

45. as you go across a period –electronegativity increases the atoms become smaller as the effective nuclear charge increases –easier to attract the shared pair of electrons as they will be in an orbital closer to the nucleus moving from L to R on the table

46. –M ELTING POINT down group 1 (alkali metals) –decreases down group 17 (halogens) –decreases Element Melting Point (K) Li453 Na370 K336 Rb312 Cs301 Fr295

48. increases

49. across the table (period 3) –from left to right increases until group 14 (think diamonds) then decreases starting at group 15 (gases)