1. Unit 4: The Periodic Table How is the periodic table a useful tool?

2. Introduction to the Periodic Table 1

3. Johann Döbereiner- 1817
Arranged elements into groups of three (triads) 1

5. John Newlands- 1864 The “Law of Octaves”
When arranged in order of increasing atomic weight, an element showed properties similar to those of an element 8 places ahead and 8 places behind
This arrangement did not work once the noble gases were discovered (they were discovered much later because they are stable/unreactive and are not found in compounds) 1

6. Dmitri Mendeleev- 1869 1

7. Dmitri Mendeleev- 1869 Published a table of 63 elements
Used properties to sort elements into groups
Arranged the elements in order of increasing atomic mass
Left blanks in table for undiscovered elements 1

8. Henry Moseley- 1913
Arranged elements according to atomic number (modern periodic table) 1

9. Groups vs. Periods Groups: Vertical columns1
Groups: Vertical columns
All elements in the same group have the same number of valence electrons
Groups 1-2 and are considered main group elements (also noted IA-IIA and IIIA- VIIIA)
Groups 3-12 are also noted IIIB- IIB
Periods: Horizontal rows
All elements in the same period have the same number of occupied principle energy levels (PELs)

10. Types of Elements
Metals, Nonmetals, and Metalloids 1

11. Physical Properties of Metals
Good conductors of heat and electricity
Have luster (shiny)
All are solid at STP, except for Hg (liquid at STP)
STP stands for Standard Temperature and Pressure (Reference Table A)
Ductile: can be drawn into wires
Malleable: can be hammered into thin sheets without breaking
High melting points
High density 1

12. Examples of Alloys 1

13. What is an alloy?
Homogenous mixture composed of two or more elements, at least one of which is a metal
Properties are often superior to those of their component elements
Ex. Steel, Bronze, Brass 1

14. Physical Properties of Nonmetals1
Most are gases at STP
A few are solids at STP (ex. sulfur and phosphorus)
Bromine is the only liquid non-metal at STP
Look dull
NM solids are poor conductors
NM solids brittle and hard
Have low densities
Have lower melting points

15. Metalloids (Semi-metals)
Have some properties of metals and some properties of nonmetals
B, Si, As, Te, Ge, Sb 1

16. What is an ion?
A charged particle that results from the LOSS or GAIN of electrons
Positively charged  cation
Negatively charged  anion
Atoms gain/lose electrons in order to obtain a stable valence e- configuration like the noble gases
The Octet Rule: All other elements want a full valence shell of 8 e- like the noble gases (except for He which is stable with 2 valence e-) 1

17. Chemical Properties of Metals and NMs1
Metals tend to LOSE electrons and become positively charged ions
Non-metals tend to GAIN electrons and become negatively charged ions

18. Have one valence electron Highly reactive with water
Group 1: Alkali Metals
Have one valence electron
Highly reactive with water
Lose one electron to become +1 charged ions 1

19. Group 2: Alkaline Earth Metals Have two valence electrons
Reactive with water
**Elements in Groups 1 and 2 are so reactive that they are never found alone in nature (only found in stable compounds)
Lose two electrons to become +2 charged ions 1

20. Transition Metals Groups 3-12: Transition Metals
Can have multiple oxidation states
Colored ions must be transition metals 1

21. Inner Transition Metals (Rare Earth Metals)
Lanthanide series: an extension of period 6
Actinide series: an extension of period 7 1

22. The Halogens Have 7 valence electrons
Group 17- The Halogens
Have 7 valence electrons
Gain one electron to become -1 charged ions 1
Start here period 5 on Friday

23. The Noble (Inert) Gases
Group 18- Noble (Inert) Gases
Stable/ Non-reactive
Have a complete valence shell of 8 e- (except He which has 2 valence e-) 1

24. What is an allotrope?
Allotrope: Different forms of the same element that have different physical and chemical properties
Ex) Carbon: Coal, Graphite, Diamond, Fullerene
Oxygen: O2 and O3
Phosphorus: Red and White 1
Start here Period 1 Friday

25. What trends exist within the PT?
There are trends as you go DOWN a group and ACROSS a period 1

26. Atomic Radius
Atomic radius: distance from the nucleus to the outermost occupied energy level 1
n=2
n=1
n=3

27. Ionization Energy
Ionization Energy: the amount of energy it takes to remove an electron from an atom
Units: kJ/mol (values are on Table S) 1

28. Electronegativity
Electronegativity: the ability of an atom to attract electrons toward itself (also on Table S)
Based on the Pauling Scale ( ) 1

29. Group Trends (based on pg.15)
As you go DOWN a group:
1) Atomic radius increases
Why? Increase in the number of occupied PELs
2) Ionization energy decreases
Why? The shielding effect
Shielding (inner) electrons reduce the pull the nucleus has for the outermost electrons, making it easier to remove an electron
3) Electronegativity decreases
Shielding makes it harder for the nucleus to attract electrons toward itself 1

30. Group Trends 4) Reactivity increases
As you go down a group, it becomes easier to lose electrons due to the increased number of shielding electrons
5) Metallic character increases
As you go down a group, it becomes easier to lose electrons, which is a property of metals. 1

31. Alkali Metals Reactivity 1

32. Period Trends (based on pg. 16)
Atomic radius decreases
Why? Increasing nuclear charge, due to the increasing number of protons in the nucleus, pulls e- closer
Ionization energy increases
Why? Increasing nuclear charge makes it harder to remove electrons
3) Electronegativity increases
Why? Increasing nuclear charge makes it easier to attract e-
4) Metallic character decreases
As you go across, it becomes harder to lose electrons, which is a property of metals. 1

33. Other Trends: Ionic Radii
Ionic radius: the radius of an atom’s ion
Cations are always smaller than the atoms from which they form
Why? Protons > electrons…so the nucleus can pull electrons in closer 1

34. Other Trends: Ionic Radii
Anions are always larger than the atoms from which they form
Why? Electrons> protons….The nucleus can’t pull the extra electrons as close together 1

35. Warm-Up Draw the Lewis Dot diagram for the ATOM of phosphorus
Draw the Lewis Dot diagram for the ION of phosphorus
Draw the Lewis Dot diagram for the ATOM of calcium
Draw the Lewis Dot diagram for the ION of calcium
Explain WHY it is easier to remove an electron from barium in comparison to calcium. 1