2. Do not copy any notes in green lettering for this unit!

3. Dmitri Mendeleev (1869)
First organize the elements in groups according to their physical and chemical properties
Predicted undiscovered elements and their properties

4. Henry Moseley
Reorganized Mendeleev’s table in order of increasing atomic number
Just like today’s modern Periodic Table

5. Periodic Law
When elements are arranged according to atomic numbers, elements with similar properties appear at regular intervals

6. Groups – vertical column of elements; share chemical properties
Aka: Families
Periods – horizontal row of elements
VC – Periodic Table Overview

7. Valence Electrons Electrons in outer most energy level
Determine the chemical properties of an element
Show students how to count valence electrons using periodic table.
Explain elements in the same group have the same chemical properties b/c they have the same
number of valence electrons (This also explains the periodic law!)

8. Some elements more reactive than others!
The closer an atom is to having 8 valence electrons, the more reactive it is
Atoms with 8 valence electrons are unreactive (aka stable)
Use Bohr models of sodium vs magnesium & sulfur vs chlorine to explain why group 1 is
more reactive than group 2 & why group 17 is more reactive than group 16

9. Metals (in blue) Good conductors of heat & electricity
Shiny surface appearance
Malleable – able to be hammered into sheets
Ductile – able to be drawn into a wire
VC – Properties of Metals: Malleability & Ductility

10. Nonmetals (in green) Poor conductors of heat and electricity
Dull surface appearance
Very brittle

11. Metalloids Characteristics of metals & nonmetals
Semiconductors (conduct electricity only at high temperatures)
Surface can appear shiny or dull
More brittle than metals, but more malleable than nonmetals
VC – Comparing Metals, Nonmetals, & Metalloids

12. Main Group Elements Groups 1,2, 13-18
Very wide range of physical and chemical properties
S and P blocks
AKA: Representative Elements

13. Alkali Metals Group 1 Extremely reactive
(1 valence electron)
React with water to make alkaline (basic) solutions
Soft, low density

14. Alkaline-Earth Metals
Group 2
Highly reactive (2 valence electrons)
Harder than alkali metals

15. Halogens Group 17 Nonmetals Highly reactive
(7 valence electrons)
React with most metals to produce salts
Greek – “salt maker”

16. Noble Gases Group 18 Unreactive (8 valence electrons)
Helium used in balloons
Others used in lamps

17. Transition Metals
Groups 3 – 12
D and F blocks
Relatively unreactive

18. Lanthanide & Actinide series
F block
Lanthanides – aka rare-earth series
Actinides are radioactive –unstable nucleus spontaneously breaks apart
Uranium used in nuclear power plants & bombs

19. Alloy – homogeneous mixture of two or more metals
Gold & silver in jewelry
Copper & zinc make brass
Iron mixed with a variety of elements to produce different types of steel

20. Hydrogen Most common element in the universe one proton + one electron
Behaves unlike all other elements
Found in all living things

21. Diatomic
The seven elements that exist in nature as two atoms of the same element bonded together
BrINClHOF or Br2 I2 N2 Cl2 H2 O2 F2
7-up

22. Periodic Trends

23. Atomic Radius – half the distance between the nuclei of two identical atoms that are not bonded together
Use Bohr Models to explain trend

24. Ionization Energy – amount of energy required to remove an electron from an atom
Electron Shielding
Use Bohr Models and magnets to explain why

25. Electron Shielding
Electrons on inner energy levels reduce the attraction between the nucleus and valence electrons (valence electrons held loosely to the atom)
Causes atoms to get bigger down a group
Causes ionization energy to decrease down a group
No affect across periods
Use Bohr Models to explain why

26. Electronegativity – ability of an atom in a chemical compound to attract electrons
Use Bohr Models to explain why

27. Periodic Trends Decreases Decreases Increases Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases

28. Electron Configuration
The arrangement of electrons in atoms

29. Electron Cloud made of Orbitals
Orbital – three-dimensional region around the nucleus that indicates the probable location of an electron
2 electrons per orbital
4 types of orbitals: s, p, d, f

31. Orbital blocks on periodic Table (pg 119)

32. energy level 1s2 that orbital
The Break Down
# of electron in
energy level 1s2 that orbital
type of orbital

33. The Rules Every box on the P.T. represents 1 electron
Each row represents an energy level
Each block represents an orbital
Start with H, read left to right across the rows until the superscripts add up to the atomic # of the element you want
Be careful when you get to the D and F blocks

34. Try These! Write the electron Configuration for:
Boron, Phosphorus, Calcium

35. Follow the arrows 4f14 5f14 3d10 4d10 5d10 6d10 2p6 3p6 4p6 5p6 6p6

36. DO NOT DO LIKE THE TEXT BOOK DOES for elements after calcium:
s-d-p is correct
d-s-p is wrong

37. Try These! Write the electron Configuration for:
Nickel, Strontium, Iodine

38. Noble Gas Notation
Same rules as electron configuration, except don’t start with H, start with the last noble gas before the element you want

39. Try These! Write the Noble Gas Configuration for:
Boron, Phosphorus, Calcium, Nickel,
Strontium, Iodine