1. The Periodic Table History Arrangement of Elements
Electron Configuration Trends
Periodic Trends
Reactivity

2. A. Johann Dobreiner’s Law of Triads in 1817

3. B. John Newlands – Law of Octaves

4. Lothar Meyer (1835-1895 - German) repetitive pattern when they are
Properties of elements show a
repetitive pattern when they are
arranged by atomic mass
D.Dimitri Mendeleev ( Russian)
(father of modern periodic table)
Published system used today (1869)
2. Elements arranged by increasing mass
3. Left spaces for elements not yet
discovered - predicted properties
(scandium, gallium, germanium)

5. Dimitri Mendelev

6. Mendeleev’s Table His table re-organized

7. Mendeleev’s Periodic Table

8. E. Henry Mosley (1887-1915) English
1.Arrange elements
by increasing atomic number – this led to the periodic law
2. Periodic Law - properties
of elements are periodic
functions of their atomic #
periodic repetition of
physical and chemical
properties

9. II. Arrangement of Elements
Periodic Table – arrangement of elements in order of increasing atomic number so that elements with similar properties are in the same column
period – horizontal row (7)
group(family)- vertical columns (1-18)
periodicity – reoccurrence of similar
properties of elements in groups

11. C. Special Groups on the Periodic Table

12. D. Periodic Table Showing s,p,d,f Blocks

13. E. Metals – Metalloids - Nonmetals
1. Metals are on the left side – all are solids
except mercury (Hg)
a. elements near the left of a period are more metallic than those near the right
b. elements near the top of a group are more metallic than those near the bottom
2. Metalloids – group of elements between
metals and nonmetals(B,Si,Ge,As,Sb,Te)
3. Nonmetals are on the right side – all are
solids or gases except bromine(Br) liquid

14. Metals – Metalloids - Nonmetals

15. PROPERTY METAL NONMETAL
Luster high low
Deformability malleable brittle
and ductile
Conductivity good poor
Electron gain/lose lose gain
Ion formed cation (+) anion(-)
Ionization energy low high
Electronegativity low high

16. IV.Periodic Trends(Main Group Elements)
Atomic Radii
1. atomic radius is ½ the distance
between nuclei of identical atoms
joined in a molecule
2. decreases across periods (left-right)
a. caused by increasing attraction
between protons and electrons
3. increases from top to bottom
a. caused by adding electrons to
new shells

17. What is the atomic radius?
Atomic radii include the region in which electrons are found 90% of the time

18. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

19. Periodic Trends in Atomic Radii

20. Periodic Trends in Atomic Radii

21. Trends in Main Groups

22. Atomic Radii Period Trends

23. A. Periodic Trends in Atomic Radii

24. B. Ionization Energy
1. Energy required to remove an electron
from an atom of an element (KJ/mol)
2. Increases across periods (left to right)
a. result of increased nuclear attraction
3. Decreases down groups (families)
a. electrons added to higher energy levels
b. shielding effect of inner shell electrons
c. repulsion of inner shell electrons
4. Energy to remove second and third electron is greater

25. B. Trend in Ionization Energy

26. B. Periodic Trends in Ionization Energy

27. Symbol First Second Third
HHeLiBeBCNO F Ne

28. C. Electronegativity
Measures how strongly one atom attracts the electrons of another atom when they form a compound
Increases across periods (left to right)
a. Fluorine has greatest value of 4
3. Decreases down groups
a. electrons far from the nucleus in
larger atoms have less attraction
b. Cesium and Francium with large radii
have the smallest electronegativity

29. Periodic Trends in Electronegativity

30. C. Periodic Trends in Electronegativity

31. C. Periodic Trends in Electronegativity

32. D. Ionic Radii
Ion – atom that acquires a charge by
gaining or losing electrons
a. cation (+) ion anion (-) ion
2. Period trends
a. cation radii decrease across periods
b. anion radii increase across periods
3. Group trends
a. increase in cation and anion radii
down groups

33. Formation of an Anion (- ion)

35. D. Comparison of Atomic and Ionic Radii

36. D. Periodic Trends in Ionic Radii

37. E. Electron Affinity
1. Energy change that occurs when an electron is added to a neutral atom
2. If it is easy to add an electron to an
atom the energy value is negative
a. halogens have large negative values
3. If it is difficult to add an electron to an
atom the energy value is positive
a. atoms in groups 2 and 18 have high
positive values (due to filled subshells)
b. usually higher values in larger atoms

39. Electron Affinity for Chlorine

40. Periodic Trends in Electron Affinity

41. Periodic Trends in Electron Affinity

42. PERIODIC TRENDS

43. Periodic Trends in Melting Point

44. Periodic Trends in Density

45. V. Reactivity
Reactivity – measure of the tendency of an element to engage in chemical reactions by losing, gaining or sharing electrons
1. atoms of reactive elements are very
likely to gain, lose or share electrons
2. atoms of reactive elements are likely to form chemical bonds with other elements

46. B. Reactivity and the Periodic Table
1. alkali metals (group 1) most reactive
metals
2. alkaline earth metals (group 2)
second most reactive group of metals
3. halogens (group 17) most reactive
nonmetals
4. noble gases (group 18) least reactive
C.Ionization Energy and Electronegativity
1. elements with very high and very
low values are very reactive

47. Electron Arrangement and Reactivity

48. Electron Configuration
S block [groups 1 and 2]
P block [groups 13,14,15,16,17,18]
D block [groups 3,4,5,6,7,8,9,10,11,12]
F block (lanthanide and actinide series)

49. 1s group 1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87

50. S- block s1 s2 Alkali metals all end in s1
Alkaline earth metals all end in s2
Should include He but helium has the properties of the noble gases.
- its outer shell is filled with the
maximum number of electrons
allowed for the first shell (2)

51. He
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

52. The P-block p1 p2 p3 p4 p6 p5

53. Transition Metals -d block

54. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
inner transition elements- hold a maximum of 14 therefore there are 14 elements in both the actinides and lanthanides f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

55. Group Ion Formed Electron Changes
Group 1 X+
Group 2 X2+
Group 13 X3+
Group 14 Xvaries
Group 15 X3-
Group 16 X2-
Group 17 X-
Group 18 X0
Loses 1 electron
Loses 2 electrons
Loses 3 electrons
Varies
Gains 3 electrons
Gains 2 electrons
Gains 1 electron
Does not gain or lose electrons