1. The Periodic Table and Trends Topics 2 and 3 SONG Please have a periodic table out.

2. IB prefers 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. –I ONIC RADII –atoms tend to gain or loose electrons in order to have the electron configuration of a noble gas –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

29. DARK GREY IS THE SIZE OF THE ION

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

33. –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 and shielding increases Li 0.152 nm Li+.078nm + Li forming a cation

34. –anions (- ions) are larger than parent atom have gained an electron 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 + -

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. the increase in ionization energy is not continuous across the table electrons are also harder to remove… –a sub-level (s,p,d,f) is completely filled –a sub-level (s,p,d,f) are half filled

42. only last sub- level shown notice discontinuity as move across period 3

43. E LECTRON A FFINITY (L OVE ) –the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion –in other words, the neutral atom's likelihood of gaining an electron –example F(g) + e -  F - (g) will release 328 kJ/mole of energy –trends are irregular and IB should not be requiring this

44. 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.

46. 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 valence electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons H Li Na K Rb

48. as you go across a period –electronegativity increases the atoms become smaller as the effective nuclear charge increases –easier to attracts electrons as they will be in a level closer to the nucleus moving from L to R on the table

49. next concept requires understanding of concepts covered in later topics only need to know the trends, not the reason why until later –metals do what they do –Van der Waals forces –bonding

50. –M ELTING POINT down group 1 (alkali metals) –decreases as “sea of negative electrons” are farther away from the positive metal ions down group 7 (halogens) –increases 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

52. increases

53. 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

54. 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 attracts electrons better since they can be close to the nucleus 1+ charge 1- charge

55. most reactive least reactive

56. 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

57. –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