Chapter 6 The Periodic Table p. 154 The Elements by Tom Lehrer The Elements by Tom Lehrer

Organizing the Elements Chemists used elements properties to sort into groups J. W. Dobereiner  triads – groups of 3 w/ similar properties One element in triad had properties intermediate of other 2 elements Cl, Br, and I look different…. similar chemically

Mendeleev’s Periodic Table 1800s, about 70 elements known Dmitri Mendeleev – Russian chemist & teacher Arranged elements by atomic mass

Mendeleev’s Periodic Table Blank spaces  undiscovered elements Predicted properties  predictions very accurate Problems w/ order  Te to I atomic mass decreases  I belongs w/ Br & Cl  Mendeleev broke rule – put Te before I

A better arrangement 1913, Henry Moseley  British physicist  Determined atomic #’s  Modern PT arranged by atomic #

The Elements by Tom Lehrer

Periodic Law Elements arranged by increasing atomic #, periodic repetition of properties present Horizontal rows = periods  7 periods Vertical column = group (family)  Similar properties  IUPAC labels (1-18) U.S. system (# & letter…i.e. IA, IIA)

Areas of periodic table 3 classes of elements:  1) Metals: electrical conductors, lustrous, ductile, malleable

 2) Nonmetals: generally brittle/non- lustrous, poor conductors of heat and electricity Some gases (O, N, Cl) some brittle solids (B, S) fuming red liquid (Br)

 3) Metalloids : border the line-2 sides  Properties are intermediate between metals and nonmetals

Section 6.2 Classifying the Elements p. 161

Groups of elements - family names Group IA – alkali metals  Forms “base” (or alkali) when reacting w/ H 2 O (not just dissolved!) Group 2A – alkaline earth metals  Also form bases with H 2 O; don’t dissolve well, hence “earth metals” Group 7A – halogens  Greek hals (salt) & genesis (to be born)

Electron Configurations in Groups sorted based on e- configs: 1) Noble gases 2) Representative elements 3) Transition metals 4) Inner transition metals Let’s now take a closer look at these.

Electron Configurations in Groups 1) Noble gases 1) Noble gases in Group 8A (also called Group 18)  very stable = don’t react  e- configuration full full outer s & p sublevels

Electron Configurations in Groups 2) Representative Elements 2) Representative Elements Groups 1A - 7A  Properties vary “Represent” all elements NOT filled s & p sublevels of highest PEL NOT filled Group # = valence e-’s

Electron Configurations in Groups 3) Transition metals 3) Transition metals in “B” columns s sublevel full  outer s sublevel full  Start filling “d” sublevel  “Transition” btwn metals & nonmetals

Electron Configurations in Groups 4) Inner Transition Metals 4) Inner Transition Metals below PT, 2 horizontal rows  outer s sublevel full  Start filling “f” sublevel

1A 2A3A4A5A6A 7A 8A Elements 1A-7A groups called representative elements outer s or p filling

The group B called transition elements u These are called the inner transition elements, and they belong here

Group 1A called alkali metals (but NOT H) Group 2A called alkaline earth metals H

Group 8A are noble gases Group 7A called halogens

Periodic table rap Let’s take a quick break……

1s11s1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6 s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in these configurations of the alkali metals?

He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s21s2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 Do you notice any similarity in the configurations of the noble gases?

Elements in the s - blocks Alkali metals end in s 1 Alkaline earth metals end in s 2  should include He, but… properties of noble gases  full outer EL  group 8A s2s2 s1s1 He

Transition Metals - d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10 Note the change in configuration.

The P-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

F - block Called “inner transition elements”

Each period # = energy level for s & p orbitals Period Number

“d” orbitals fill up in levels 1 less than period #  first “d” is 3d found in period d 4d 5d

f orbitals start filling at 4f….2 less than period # f5f

Demo p. 165

Trends in Atomic Size Atomic Radius - half distance btwn 2 nuclei of identical atoms  Increases top-bottom  Decreases L-R  picometers m… 1 trillionth Radius Section 6.3 Periodic Trends p. 170

ALL ALL PT Trends Influenced by 3 factors: Energy Level 1. Energy Level  Higher energy levels further from nucleus Charge on nucleus 2. Charge on nucleus (# protons)  More + charge pulls e -’ s in closer Shielding effect 3. Shielding effect

#1. Atomic Size - Group trends Going down a group, atoms gain another PEL (floor)  atoms get….. H Li Na K Rb b i g g e r

#1. Atomic Size - Period Trends L to R across period:  More p+ in nucleus  More e-’s occupy same energy level stronger nuclear charge Pulls e- cloud closer to nucleus  atoms get…. Na Mg Al Si P S Cl Ar Here is an animation to explain the trend. S m a l l e r

Trends of Atomic Radius increases decreases increases

p. 172 Ions p. 172 Some compounds composed of “ions”  Ion - atom (or group of atoms) w/ + or - charge formed when e- transferred btwn atoms  Cation (loses e-’s…+ ion)  Anion (gains e-’s… - ion)

Cation Formation 11p+ Na atom 1 valence e- Valence e- lost in ion formation Effective nuclear charge on remaining e-’s increases. Remaining e- pulled closer to nucleus. Ionic size decreases. Result: a smaller sodium ion, Na +

Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell (from Na for example). Effective nuclear charge is reduced and the e- cloud expands. ion A chloride ion is produced. It is larger than the original atom.

p.173 #2. Trends in Ionization Energy p.173 Ionization energy - energy required to completely remove e- (from gaseous atom) energy required to remove only 1st e- called first ionization energy.

Ionization Energy second IE is E required to remove 2nd e-  Always greater than first IE third IE greater than 1st or 2nd IE IE helps predict what ions elements form Li 1+ Mg 2+ Al 3+

SymbolFirstSecond Third H He Li Be B C N O F Ne Why did these values increase so much ? Table 6.1, p. 173

Cation Formation 11p+

Anion Formation 17p+

What factors determine IE? greater nuclear charge = greater IE Greater distance from nucleus decreases IE Filled & half-filled orbitals have lower energy  Easier to achieve (lower IE) Shielding effect

Shielding Effect e-’s in outer PEL “look thru” other PEL’s to “see” nucleus Stays same thru blocks Greater influence on IE than nuclear charge

Shielding Trends remains constant increases

Ionization Energy - Group trends p. 174 going down group  first IE decreases b/c... e- further from p+ attraction  more “shielding”

Ionization Energy - Period trends p. 174 Same period atoms have same: # energy levels “shielding” (within a block – slight decrease btwn “s” and “p”)  Increasing nuclear charge  IE generally increases left - right Exceptions…full & 1/2 full orbitals

First Ionization energy Atomic number He He greater IE than H. Both w/ same shielding (e- in 1st level)  He - greater nuclear charge H

First Ionization energy Atomic number H He Li lower IE than H more shielding further away l These outweigh greater nuclear charge Li

First Ionization energy Atomic number H He Be higher IE than Li = shielding (period) l greater nuclear charge Li Be

First Ionization energy Atomic number H He B has lower IE than Be l greater nuclear charge l shielding has greater influence on IE l Slight decrease (“p” e-) l “p” e- removed l “s” orbital ½ filled Li Be B

First Ionization energy Atomic number H He Li Be B C

First Ionization energy Atomic number H He Li Be B C N

First Ionization energy Atomic number H He Li Be B C N O Oxygen breaks the pattern, b/c removing e- leaves it w/ 1/2 filled p orbital

First Ionization energy Atomic number H He Li Be B C N O F

First Ionization energy Atomic number H He Li Be B C N O F Ne Ne has a lower IE than He  Both full but… Ne more shielding  b/c greater distance

First Ionization energy Atomic number H He Li Be B C N O F Ne Na has a lower IE than Li Both are s 1 Na - more shielding l Greater distance Na

Trends in Ionization Energy (IE) increases decreases

Trends in Ionic Size: Cations Cations lose e-’s  metals Cations smaller than atom they came from  lose e-’s  lose entire energy level. Cations of representative elements have noble gas config before them

Trends in Ionic size: Anions Anions gain e-‘s  nonmetals Anions bigger than atom they came from  same energy level  greater area nuclear charge needs to cover

Configuration of Ions Ions always have noble gas configurations (full outer level) Na atom is: 1s 2 2s 2 2p 6 3s 1  Forms 1+ Na ion: 1s 2 2s 2 2p 6  Same as Ne

Configuration of Ions Non-metals form ions by gaining e-’s to achieve noble gas configuration  configuration of noble gas after them

Ion Group trends Each step down a group adds energy level Ions - bigger going down  more energy levels Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

Ion Period Trends Across period  nuclear charge increases  Ions get smaller energy level changes btwn anions & cations Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

#3. Trends in Electronegativity Electronegativity (EN)- tendency for atom to attract e-’s when atom in cmpd Sharing e-, but how equally? Element w/ big EN pulls e- towards itself strongly!

Electronegativity Group Trend Further down group, farther e- away from nucleus  plus more e-’s atom has more willing to share Low EN

Electronegativity Period Trend Metals let e-’s go easily  low EN Nonmetals want more e-’s  take e-’s from others  High EN

Trends in Electronegativity 0 decreases increases decreases

Chemistry Song "Elemental Funkiness" - Mark Rosengarten The Elements – Tom Lehrer