1. History Structure Trends
The Periodic Table
History
Structure
Trends

2. Mendeleev
Arranged elements in order of atomic mass with 8 long columns and several short columns
Arrangement reflected properties of the elements
Predicted existence of several elements to fill gaps in his table
These elements were later discovered and had the properties predicted

3. Mendeleev’s Periodic Law
Properties of the elements are a periodic function of their atomic masses.

4. Moseley
Performed experiments to determine an accurate mass for several elements which seemed out of place on the table
Noticed a pattern in the number of protons
Reorganized elements in order of atomic number rather than mass

5. Modern Periodic Law:
Properties of elements are periodic functions of their atomic numbers.

6. Periods 1 2
7 rows 3 4 5 6 7

7. Groups or Families 1 18 2 13 14 15 16 17 3 4 5 6 7 8 9 10 11 12

8. This Arrangement Reflects:
Properties
Increasing atomic number
Electron configuration

9. Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1
Electron configuration and properties by group, ex. Group 1 Li
Li – 1s22s1
Na – 1s22s22p63s1
K – 1s22s22p63s23p64s1 Na K
Same Group (family)
All have 1 e- in the outer level
All are similar in color and hardness
All react vigorously with water

10. Electron configuration by period, ex. Period 2
C - 1s22s22p2
N - 1s22s22p3
O - 1s22s22p4
Same period
All have a full 1st energy level with e- in s and p orbitals of the second level
Number of e- is increasing by 1

11. Blocks
Electron configuration for members of block end by filling designated orbital within period s p d f

12. General Properties of Metals
1 to 3 e- in outer energy level of most
Lose e- to form + ions (cations)
Shiny (lustrous)
Hard
Malleable and ductile
Good conductors of heat and electricity
Solid at room temp (298K), except Hg

13. General Properties of Nonmetals:
5 to 8 electrons in the outer level
Gain e- to form – ions (anions)
Brittle solids or gases
Poor conductors of heat
Poor conductors of electricity

14. Classification of Elements:
Some elements have properties similar to both metals and nonmetals.
These are found bordering the stair-step dividing line.
Exception: Al is a metal
These elements are called metalloids

15. Metals, Metalloids, and Nonmetals

16. Some of the Families Have Special Names.
Metals
Non-metals
17 – halogens
18 – noble gases
1 – alkali metals
2 – alkaline earth metals
3-12 – transition metals
Elements 58 – 71 lanthanides
Elements 90 – actinides
Rare Earth Metals

17. Trends: Atomic Radius
Radius of an atom without regard to surrounding atoms (size)
Radius depends on:
the number of energy levels
the strength of the nucleus r

18. As n increases, size of the e- cloud increases.
Each period represents a higher principal quantum number (n)
As n increases, size of the e- cloud increases. F
Atomic radius increases down a family. Cl Br

19. Across a period, nuclear charge increases by 1 for each element.
A stronger nucleus acts like a stronger magnet which attracts the e- cloud. C N O
Atomic radius decreases across a period.

20. General Trend for Atomic Radius:
Increases

21. Use position in the periodic table to determine which is larger?
Ag or Au?
Ni or Cu?
La or U?
H or He?
Na or Rb?
Cl or I?
Al or Si?
K or Ca?

22. Ionic Radius Ions are charged atoms formed when: Atoms lose e-
(+ Ion) = Cation
Atoms gain e-
(- Ion) = Anion

23. Cations are smaller than their respective neutral atoms.
Metals usually lose all valence electrons. So then, the atom has one less energy level.
- e- +
The nucleus pulls tighter on the remaining electrons.

24. Anions are larger than their respective neutral atoms.
Nonmetals usually gain electrons to complete the valence shell.
+ e- -
Electrons repel each other and spread out more. - -
more e- than p+

25. Which is larger?
Ca or Ca+2
F or F–1
K or K+1
O or O-2
F-, Ne or Na+

26. First Ionization Energy
Energy needed to remove the most loosely held electron from a gaseous atom.
Factors that affect ionization energy:
Radius
Nuclear charge
Shielding effect
Stability of sublevels

27. Radius:
The greater the distance between the nucleus and the valence electrons, the easier it is to lose an electron.

28. Nuclear charge: # of p+
Within a period, the higher the nuclear charge, the higher the ionization energy.

29. Shielding Effect:
Other e- block the pull of the nucleus on the outer e-.
Electrons repel each other.

30. General Trend for Ionization Energy:
Increases

31. Electron Affinity: Attraction of an atom for an additional electron.
Factors affecting EA:
Size
Nuclear charge
Shielding effect
Stability of sublevels

32. General Trend for Electron Affinity:
Increases

33. Electronegativity = ability of atom to become negative ion
Ability to “steal” electrons
Increases across a period and up a row

34. Valence electrons – see group #