1. Periodicity
Glencoe
Chapter 6

3. Development of the Modern Periodic Table
1790s – 23 known elements
By 1870s – 70 known elements
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)

4. 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

5. 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, Click here for a translation

6. 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

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

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

9. 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

11. 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

13. 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”

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

15. 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-)

16. 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!....

17. 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

19. Summary of Trends