1. The Periodic Table & Periodic Law
Unit 6
The Periodic Table & Periodic Law

2. 6.1 Development of the Modern Periodic Table
The periodic table developed over time as scientists discovered more useful ways to compare and organize the elements

3. Development of the Periodic Table
Antoine Lavoisier – late 1700s
Compiled first list of elements known at the time (33 of them)
Organized elements into four categories
Gases
Metals
Nonmetals
Earths

4. By 1860 60 elements had been discovered but scientists had no way of organizing them

5. J.W. Dobereiner classified elements that had similar properties into triads, organizing them by atomic mass

6. Triads were useful because they grouped elements with similar properties and revealed an orderly pattern in some of their physical and chemical properties which are related to related to atomic mass.

7. John Newlands – 1864
Noticed when elements were arranged by increasing atomic mass their properties repeated every 8th element
Pattern is periodic (repeats in a specific manner)
Named this the law of octaves
Law did not work for all known elements

8. Dmitri Mendeleev – 1869
Arranged elements in order of increasing atomic mass into columns with similar properties
Predicted the existence and properties of undiscovered elements that were later found

9. Mendeleev’s Periodic Table

10. When elements are arranged according to mass, some elements end up in groups with properties different from their own
Henry Moseley – 1913
Arranged table by atomic number

11. Periodic law – when elements are arranged according to increasing atomic number, there is a periodic repetition of chemical and physical properties

12. The Modern Periodic Table
Columns = groups
Numbered 1 – 18
Correspond to the number of outermost electrons
Have similar properties
Some have special names
Rows = periods
Numbered 1 – 7
Correspond to outermost energy level

13. Representative elements – in s & p block
Transition elements – in d block

14. What element is in group 4 & period 5?
Is this a representative element?
What element is in group 18 & period 6?

15. 6.2 Classification of the Elements

16. Physical States and Classes of Elements
Most elements are solid at room temperature
Br & Hg are liquid
N, O, F, Cl, and Noble gases are gas

17. Metals Shiny Solid at room temperature
Good conductors of heat & electricity
Malleable
Ductile
Found to the left of the staircase
1 exception =

18. Alkali metals – group 1 elements
Except for
Very reactive
# of valence electrons =

19. Alkaline earth metals – group 2 elements
Also highly reactive
# of valence electrons =

20. Transition metals – group 3 – 12 (d block)
Number of valence electrons varies

21. Inner transition metals – f block
Number of valence electrons varies
Lanthanide & actinide series

22. Nonmetals Gases or dull looking solids
1 exception =
Poor conductors of heat and electricity
Found to the right of the staircase

23. Halogens – group 17
Highly reactive
# of valence electrons =

24. Nobel gases – group 18
Extremely unreactive
# of valence electrons =

26. Metalloids Border the staircase
Have properties of both metals and nonmetals
Semi-conductors = conduct electricity only under certain conditions

28. Valence electrons – electrons in the highest energy level of an atom
determine an atoms properties
All atoms in same group have same number of valence electrons and therefore…

29. Valence electrons and period
Period an element is in = energy level of its valence electrons
Li =
Ga=

30. Valence electrons of representative elements

31. 1 = s1 (1 valence electron)
2 = s2 (2 valence electrons)
13 = s2p1 (3valence electrons)
14 = s2p2 (4 valence electrons)
15 = s2p3 (5 valence electrons)
16 = s2p4 (6 valence electrons)
17 = s2p5 (7 valence electrons)
18 = s2p6 (8 valence electrons)

33. Transition Metals Outermost s and nearby d sublevels contain electrons
All transition metals usually have 1, 2, or 3 valence electrons

34. Inner Transition Metals
Outermost s and nearby f sublevels contain electrons

35. Valence electrons & ion formation
Representative elements lose or gain electrons in order to obtain the same electron configuration as a noble gas
How many valence electrons does a nobel gas have?

36. Octet rule = an atom tends to gain, lose, or share electrons in order to acquire a full set of eight valence electrons

37. Neutral atom Na
1s22s22p63s1 B
1s22s22p1 P
1s22s22p63s23p5 F
1s22s22p5

38. Ion
Na1+
1s22s22p6 (Ne)
B3+
1s2 (He)
P3-
1s22s22p63s23p6 (Ar)
F1-

39. 1 = s1 (1 ve, easily lost, ion formed = )
13 = s2p1
14 = s2p2
15 = s2p3
16 = s2p4
17 = s2p5
18 = s2p6

40. 6.3 Periodic Trends

41. Nuclear Charge
Proton number
Increases as atomic number increases

42. Shielding Electrons
Electrons between the nucleus and valence electrons
Shield the valence electrons from the force of the nucleus
Total electrons – valence electrons Na S

43. Shielding Effect
The further an energy level is from the nucleus, the more it is shielded from the attraction to the nucleus
The outer energy level is shielded by the inner energy levels

44. Effective Nuclear Charge
The nuclear charge that the valence electrons actually feel
ENC = total electrons – shielding electrons

45. Potassium Arsenic Total Valence Shielding ENC Total Valence Shielding

46. ENC and Group Group 1 Group 2 Group 13 Group 14 Group 15 Group 16Na
Group 2 Mg
Group 13 Al
Group 14 Si
Group 15 P
Group 16 S
Group 17 Cl
Group 18 Ar

47. Atomic Radius Within period – decreases from left to right
More valence electrons = more attracted to nucleus
Higher ENC shrinks outer energy level
Na vs Cl

48. Within group – increases from top to bottom
More energy levels = bigger atom
ENC for all atoms in a group is the same
Smallest atom =
Largest atom =

50. Ionic Radius ALWAYS SMALLER THAN NEUTRAL ATOM
Cation –
ALWAYS SMALLER THAN NEUTRAL ATOM
When Cations form the outer most energy level is lost

51. ALWAYS LARGER THAN THE NEUTRAL ATOM
Anions
ALWAYS LARGER THAN THE NEUTRAL ATOM
When atoms gain electrons they repel each other more

52. Within period – increases from left to right
Anions are larger, where are anions located?
Within group – increases from top to bottom
More energy levels = bigger atom

53. Ionization energy
Energy required to remove an electron from a gaseous atom
Determined by how strongly the valence electrons are attracted to the nucleus
Hard to remove electrons from atoms with outer energy level close to nucleus
Hard to remove electrons when ENC is large

54. Within period – increases from left to right
More valence electrons = more attraction to nucleus (higher ENC)
Within group – decreases from top to bottom
More energy levels = valence electrons further from nucleus = less attraction

56. Electronegativity
The ability of an atom to attract electrons in a chemical bond

57. Within period – increases from left to right
Higher ENC attracts electrons more easily
Within group – decreases from top to bottom
Closer outer energy levels attract electrons more

58. Highest electronegativity
Lowest electronegativity