1. Periodicity
And the Periodic Table

2. The Big Idea
Periodic trends in the properties of atoms allow us to predict physical and chemical properties

3. Development of the Periodic Table
The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements

4. The Periodic Table Mendeleev (1869) – first periodic table
Based on increasing atomic mass and repeating properties of elements
Had spaces for “missing” elements he predicted

6. Henry Moseley
Discovered in 1914 that elements’ properties more closely followed their atomic number
Modern Periodic Table based on this discovery

7. Periodic Law
Physical and chemical properties of elements are periodic functions of atomic numbers

8. Classification of the Elements
Elements are organized into different blocks in the periodic table according to their electron configurations

9. Organizing the Elements by Electron Configuration
Electrons in the highest principal energy level are called valence electrons.
All group 1 elements have one valence electron.

10. Organizing the Elements by Electron Configuration
The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found.
The number of valence electrons for elements in groups is ten less than their group number.

11. Organizing the Elements by Electron Configuration

12. What is periodicity?
Properties of the elements change in a predictable way as you move through the periodic table
These properties include
Atomic radius
Octet Rule
Ionic radius
Ionization energy
Electronegativity

13. How are the elements organized?
Periodic Table: a complete chart of all elements in the universe
Arranged according to physical and chemical properties
Each box on the table contains the atomic number, atomic mass, and chemical symbol

14. How are the elements organized?
Groups: also known as families; are the columns
Have similar properties
Some have specific names:
Family 1: Alkali Metals
Family 2: Alkaline Earth Metals
Families 3 – 12: Transition Metals
Family 13: Boron Family

15. How are the elements organized?
Family 14: Carbon Family
Family 15: Nitrogen Family
Family 16: Oxygen Family
Family 17: Halogen Family
Family 18: Noble Gases
Periods: the rows on the periodic table
Do not have similar properties

16. The Modern Periodic Table

17. Metals To the left of the stair-step line 88 elements
Tend to lose electrons
Most reactive in the s block
Includes alkali metals and alkaline earth metals

18. Properties of Metals Shiny luster Good conductors of heat
Good conductors of electricity
Usually solids at room temp
Malleable
Ductile

19. gold
lead
copper
nickel

20. Nonmetals On right side of stair-step line Tend to gain electrons
Most reactive group is halogens
Least reactive is Noble gases

21. Properties of Nonmetals
Dull luster
Poor conductors of heat
Poor conductors of electricity
Brittle
Many are gases at room temp

22. Carbon (graphite)
bromine
sulfur

23. Metalloids On either side of stair-step line
Have properties of metals and nonmetals
Includes all elements that touch the line except Al and Po
Many are used in transistors

24. antimony
germanium

25. Alkali Metals Group 1 (except H) All have only 1 valence electron
Most reactive metals; never found in pure state in nature
Soft, shiny, have relatively low melting points

26. Alkaline Earth Metals Group 2 All have 2 valence electrons
Are the second most reactive metals; never found naturally in pure state
Harder, denser, stronger than alkali metals
Have higher melting points than alkali metals

27. Transition Metals Groups 3 – 12
All have 1 or 2 valence electrons (in s sublevels)
Do not fit into any other group or family
Have many irregularities in their electron configurations

28. Boron Family Group 13 Have 3 valence electrons
Boron is a metalloid, while all of the others are metals

29. Carbon Family Group 14 All have four valence electrons
Carbon is a nonmetal; Si and Ge are metalloids; Sn and Pb are metals

30. Nitrogen Family
Group 15
All have 5 valence electrons (in s and p sublevels)
N and P are nonmetals; As and Sb are metalloids; Bi is a metal

31. Oxygen Family
Group 16
All have 6 valence electrons (in s and p sublevels)
All are nonmetals except Te, which is a metalloid, and Po, which is a metal.

32. Halogens Means “salt former” Group 17
All have 7 valence electrons (in s and p sublevels)
Most reactive nonmetals
All are nonmetals except At, which is a metalloid

33. Noble Gases
Group 18
Complete, stable electron configuration (no valence electrons)
Most unreactive elements

34. Rare Earth Elements Found in 2 rows at bottom of Periodic Table
Also known as the inner transition metals
Lanthanide series: starts with La
Actinide series: starts with Ac
Little variation in properties
Actinides are radioactive; only first three and Pu are found in nature

35. Periodic Trends
Trends among elements in the periodic table include their size and their ability to lose or attract electrons

36. Atomic Radius
For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms.

37. Atomic Radius

38. Atomic Radius
Atomic radius generally increases as you move down a group.
The outermost orbital size increases down a group, making the atom larger.

39. Ionic Radius
An ion is an atom or bonded group of atoms with a positive or negative charge.
When atoms lose electrons and form positively charged ions, they always become smaller for two reasons:
The loss of a valence electron can leave an empty outer orbital resulting in a small radius.
Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.

40. Ionic Radius
When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.

41. Ionic Radius

42. Ionic Radius
The ionic radii of positive ions generally decrease from left to right.
The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.

43. Ionization Energy The energy needed to remove one of its electrons
Decreases as you move down a group
Increases as you move across a period
Successive ionization energies increase for every electron removed

44. Ionization Energy

45. Octet Rule (Rule of 8)
Atoms tend to gain, share, or lose in order to acquire a full set of valence electrons (in most cases, this is 8)

46. Electronegativity
Reflects an atom’s ability to attract electrons in a chemical bond
Related to its ionization energy and electron affinity
Increases as you move across a period
Increases as you move up a group

47. Electronegativity