1. Periodicity AP Chemistry

2. Sodium and potassium react w /water to produce hydrogen gas. valence orbitals: outer-shell orbitals -- elements in the same group have the same valence-shell electron configuration -- since valence e – are involved in bonding, elements within a group have many of the same properties

3. Au Development of the Periodic Table -- few elements appear in elemental form in nature (Au, Ag, Hg, a few others) -- most are in combined forms with other elements -- In 19th century, advances in chemistry allowed more elements to be identified. AgHg

4. Dmitri Mendeleev 1869: Independently, Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) published classification schemes based on similarities in element properties. ** Mendeleev used his scheme to predict the existence of undiscovered elements, and so is given credit for inventing the first periodic table. ** D.M.’s guiding principle was… 1834–1907 Lothar Meyer 1830–1895 “atomic mass.”

5. -- 1913: Henry Moseley bombarded atoms with high-energy electrons and measured the frequency of the X rays given off. X ray frequency generally increased as atomic mass increased, but VERY nicely increased as ____________ increased. Henry Moseley 1887–1915 Graph of Moseley’s data atomic number

6. (shells are diffuse, i.e., “fuzzy”) Electron Shells Even before Bohr, the American Gilbert Lewis had suggested that e – are arranged in shells. -- Experiments show that e – density is a maximum at certain distances from nucleus. Gilbert Lewis 1875–1946 -- no clearly defined boundaries between shells e – density distance from nucleus Ar Ne He

7. 0.77 + 1.14 Approximate bonding atomic radii for the elements have been tabulated. The distance between bonded nuclei can be approximated by adding radii from both atoms. e.g., Bonding atomic radii are as follows: So the approximate distance between bonded C and Br nuclei = = 1.91 A C = 0.77 A, Br = 1.14 A rr d

8. more p +, but no new (i.e., farther away) energy levels Atomic Radius As we go down a group, atomic radius… -- principal quantum number increases (i.e., a new energy level is added) increases. As we go from left to right across the Table, atomic radius… decreases. -- effective nuclear charge increases, but principal quantum number is constant

9. Coulombic attraction depends on… 2– 2+ 2– 1– 2– 1+ 2+ amount of chargedistance between charges + + – – H He + – + – + – As we go, more coulombic attraction, no new energy level, more pull, smaller size

10. Sr < Ba < Cs Arrange the following atoms in order of increasing atomic radius: Sr, Ba, Cs

11. Ionization Energy: the minimum energy needed to remove an e – from an atom or ion M(g) + 1 st I.E.  M + (g) + e – M + (g) + 2 nd I.E.  M 2+ (g) + e – M 2+ (g) + 3 rd I.E.  M 3+ (g) + e – Successive ionization energies are larger than previous ones. -- (+) attractive force remains the same, but there is less e – /e – repulsion I.E. e–e–

12. The ionization energy increases sharply when we try to remove an inner-shell electron. e.g., For Mg, 1 st IE = 738 kJ/mol 2 nd IE = 1,450 kJ/mol 3 rd IE = 7,730 kJ/mol As we go down a group, 1 st IE…decreases. -- more e – /e – repulsion and more shielding (strong evidence that only valence e – are involved in bonding)

13. Generally, as we go from left to right, 1 st IE… Exceptions: e.g., B < Be Be:1s 2 2s 2 B:1s 2 2s 2 2p 1 B doesn’t like 2p (easier to remove B’s single 2p e – than one of Be’s two 2s e – s) N:1s 2 2s 2 2p 3 O:1s 2 2s 2 2p 4 More stable to have than to have 2p Subshells prefer to be either completely filled OR half-filled. This e – is easier to remove… …than any of these.

14. First across a period… down a group…

15. released. Electron affinity: the energy change that occurs when an e – is added to a gaseous atom For most atoms, adding an e – causes energy to be… Exceptions: noble gases: the added e – must go into a new, higher energy level group 2 metals: the added e – must go into a higher-energy p orbital group 15 elements: the added e – is the first one to double-up a p orbital eq. for e – affinity: A + e – A – I.E. e–e– e–e– energy

16. more (–) e – affinity The halogens have the most (–) electron affinities, meaning that they become very stable when they accept electrons. Electron affinities don’t vary much going down a group. more willing to accept an e – = –328 F Cl Br I O S Se Te Ne Ar Kr Xe He –349 –325 –295 –141 –200 –195 –190 + + + + +

17. metals: left side of Table; form cations properties: Regions of the Table ductile (can pull into wire) malleable (can hammer into shape) lustrous (shiny) good conductors (heat and electricity)

18. (i.e., they lose e – ) -- Because of their low ionization energies, they are often oxidized in reactions. -- Metallic character of the elements increases as we go down-and-to-the-left. increasing metallic character

19. nonmetals: right side of Table; form anions properties: good insulators; gases or brittle solids Br 2 I2I2 S8S8 Ne bromineiodinesulfurneon Regions of the Table (cont.) -- memorize the HOBrFINCl twins (or…Hans and Franz, the ClOBrHIFN twins) “Wer sind sie?” “Die ClOBrHIFN Zwillinge!”

20. computer chips metalloids (semimetals): “stair” between metals and nonmetals properties: in-between those of metals and nonmetals; “semiconductors” (B, Si, Ge, As, Sb, Te, Po, At) Si and Ge metals nonmetals Regions of the Table (cont.) computer chips Si and Ge

21. Reactivity Trends metal oxide + watermetal hydroxide MgO(s) + H 2 O(l) Mg(OH) 2 (aq) metal oxide + acid salt + water CaO(s) + 2 HNO 3 (aq) Ca(NO 3 ) 2 (aq) + H 2 O(l) metal + nonmetal salt 2 Al(s) + 3 Br 2 (l) 2 AlBr 3 (s) (i.e., a “basic” oxide)

22. Reactivity Trends (cont.) nonmetal oxide + water acid CO 2 (g) + H 2 O(l) H 2 CO 3 (aq) (i.e., an “acidic” oxide) nonmetal oxide + base salt + water CO 2 (g) + 2 KOH(aq) K 2 CO 3 (aq) + H 2 O(l)

23. Group Trends Alkali Metals -- the most reactive metals (one e – to lose) -- obtained by electrolysis of a molten salt e.g., chloride ion is oxidized and sodium ion is reduced 2 NaCl(l) 2 Na(l) + Cl 2 (g)

24. -- react with hydrogen to form metal hydrides: -- react with water to form metal hydroxides: -- react w /O 2 : Li yields Li 2 O, others yield (mostly) peroxides (M 2 O 2 ) 2 M(s) + H 2 (g) 2 MH(s) 2 M(s) + 2 H 2 O(l) 2 MOH(aq) + H 2 (g) 2 M(s) + O 2 (g) M 2 O 2 (s) Potassium in water, forming flammable hydrogen and soluble potassium hydroxide.

25. Alkaline-Earths -- not as reactive as alkalis (two e – to lose) -- Ca and heavier ones react w /H 2 O to form metal hydroxides -- MgO is a protective oxide coating around substrate Mg Ca(s) + 2 H 2 O(l) Ca(OH) 2 (aq) + H 2 (g) compared to alkalis: harder, denser, higher MPs Mg ribbonMgO

26. The Hindenburg Hydrogen -- a nonmetal, but belongs to no family -- reacts w /other nonmetals to form molecular (i.e., covalent) compounds (She was scuttled in June 1919, along with 71 other German ships.) (She burned up in May 1937, killing 36 passengers.)

27. Halogens -- At isn’t considered to be a halogen; little is known about it -- at 25 o C, F 2 and Cl 2 are gases, Br 2 is a liquid, I 2 is a solid -- their exo. reactivity is dominated by their tendency to gain e – -- Cl 2 is added to water; the HOCl produced acts as a disinfectant -- HF(aq) = weak acid; HCl(aq) HBr(aq) HI(aq) = strong acids A small amount of a halogen is mixed with a noble gas to fill halogen lamps. The halogen sets up an equilibrium with the tungsten filament to prevent the heated tungsten from being deposited on the inside of the bulb.

28. professional-grade Rn detector Noble Gases -- all are monatomic; have completely-filled s and p orbitals -- He, Ne, and Ar have no known compounds; Rn is radioactive -- Kr has one known compoud (KrF 2 ); Xe has a few (XeF 2, XeF 4, XeF 6 ) Fan for Rn mitigation

29. Ca atomCa 2+ ion Cl atom Cl – ion 20 p + 20 e – 20 p + 18 e – 17 p + 17 e – 17 p + 18 e – Ca Ca 2+ Cl Cl – Ionic Radius Cations are _______ than parent atoms; anions are ______ than parent atoms. EX. Compare the sizes of Fe, Fe 2+, and Fe 3+. Then compare Br with Br –. smaller larger Fe > Fe 2+ > Fe 3+ Br – > Br

30. electronegativity: the tendency for a bonded atom to attract e – to itself up and to-the-right. electronegativity increases Electronegativity Electronegativity increases going... Most electronegative element is...fluorine (F). Linus Pauling quantified the electronegativity scale.

31. “Oh, man… I forgot which ones the most electronegative elements are.” “Shee-oot… Ow teh ye… FO’ NCl.” F = 4.0 O = 3.5 N = Cl = 3.0 C = 2.5 H = 2.1 Others: