Periodicity Glencoe Chapter 6

http://www.lenntech.com/periodic/name/alphabetic.htm

Development of the Modern Periodic Table 1790s – 23 known elements By 1870s – 70 known elements 1864 - John Newlands proposed arrangements by mass and properties by octaves 1864 – Lothar Meyer proposes arrangements by mass and columns of properties—but doesn’t publish! 1869 – Dmitri Mendeleev also proposed arrangements by mass and columns of properties—and announces! 1913 – Henry Moseley proposed arrangements by atomic number. (periodic law)

Newlands’ “octaves” H 1 F 8 Cl 15 Co/Ni 22 Br 29 Pd 36 I 42 Pt/Ir 50 Li 2 Na 9 K 16 Cu 23 Rb 30 Ag 37 Cs 44 Tl 53 Gl 3 Mg 10 Ca 17 Zn 25 Sr 31 Cd 34 Ba/V 45 Pb 54 Bo 4 Al 11 Cr 18 Y 24 Ce/La 33 U 40 Ta 46 Th 56 C 5 Si 12 Ti 19 In 26 Zr 32 Sn 39 W 47 Hg 52 N 6 P 13 Mn 20 As 27 Di/Mo 34 Sb 41 Nb 48 Bi 55 O 7 S 14 Fe 21 Se 28 Ro/Ru 35 Te 43 Au 49 Os 51

Shortly after, his ideas were presented to the Russian Physico-chemical Society. They were read by Professor Menschutkin because Mendeleev was ill. His ideas were then published in the main German chemistry periodical of the time, Zeitschrift fϋr Chemie.                                                                                                                                                                                                                                                  The world’s first view of Mendeleev’s Periodic Table – an extract from Zeitschrift fϋr Chemie, 1869. Click here for a translation

Key “landmarks” of the modern periodic table Periods (horizontal) Groups/families (vertical) “Representative” elements s & p block Groups 1A – 8A Groups 1,2,13,14,15,16,17,18 “Transition” elements d block (f block = “inner transition elements”) Groups 1B – 8B Groups 3 - 12

Other notable classifications: Metals Alkali (group 1) Alkaline (group 2) Metalloids Nonmetals Halogens (group 17) Noble gases (group 18)

Periodic trends Vary systemically across a period (horizontally) down a group (vertically)

Atomic radii Based on probability of electron cloud, therefore, defined by how closely an atom lies to a neighboring atom DECREASES to the right across a period Due to larger nuclear attraction INCREASES down a group Due to more “layers” of electrons

Ionic Radii Ions (charged atoms) form when electrons are gained or lost….(the number of protons and electrons don’t match!) DECREASES to the right across a period (in two phases) INCREASES down a group

Ionization Energy Defined as “amount of energy required to remove an electron from a gaseous atom” 1st ionization energy 2nd ionization energy Etc. Think of this as the atom’s ability to hold onto its valence electron! INCREASES across a period Harder to remove e- Positive energy means “harder” DECREASES down a group Easier to remove e- More negative energy means “easier” or “more stable”

Element I1 I2 I3 I4 I5 I6 I7 Na 490 4560 Mg 735 1445 7730 Al 580 1815   Mg 735 1445 7730 Al 580 1815 2740 11,600 Si 780 1575 3220 4350 16,100 P 1060 1890 2905 4950 6270 21,200

Octet Rule Atoms tend to gain, lose, or share electrons in order to acquire a full set of eight valence electrons Note chemical stability of noble gases Predicts ionic charge of main block elements CATIONS—positively charged ions (lost e-) ANIONS—negatively charged ions (gained e-)

Electron Affinity Energy associated with adding an electron to an atom’s electron cloud---think of the opposite of Ionization Energy…but same effect! INCREASES (but the energy gets more negative = means “more stable”) across period DECREASES (but the energy value gets more positive = means “more difficult”) down group Therefore, a great idea!....

Electronegativity Indication of the relative ability of the atom to attract electrons in a chemical bond Think of this quantity as how strongly an atom might want to gain an electron. Arbitrary rating scaled to 4.0….. Most electronegative element is fluorine with 3.98 INCREASES across a period DECREASES down a group

Summary of Trends