1. Intro To The Periodic Table

2. How Was The Periodic Table Created?
Throughout history elements were discovered at different times but they were not put in any order to classify them.
By 1828, 53 elements were discovered.
To bring ORDER out of CHAOS, to classify is a basic need of man. Therefore, scientists began looking for similarities among the elements as more were continuing to be discovered.

4. Who Is Dmitri Mendeleev?
In 1871, Russian scientist Dmitri Mendeleev organized the known elements of the time according to increasing atomic mass.
He left spaces on the periodic table where elements were yet to be discovered.
Mendeleev also recognized that certain similarities in the chemical properties of elements occurred at regular intervals. He implemented John Newland’s law of octaves.
 Law of Octaves – every 8 elements tend to have the same chemical properties (groups 1-2 and on your periodic table)

5. 26 Fe
55.845 28 Ni 27 Co
Mendeleev 51 Sb 53 I 52 Te
127.60

6. Who Is Henry Moseley?
In 1911 English scientist Henry Moseley discovered that the elements on the periodic table fit into patterns better when they are arranged in order of increasing atomic number, or the number of protons in the nucleus.
Moseley’s periodic table represents the periodic law which is what we follow when we study chemistry today.
Periodic Law - the physical and chemical properties of the elements are periodic functions of their atomic numbers
*When elements are arranged in order of increasing atomic number, elements with similar properties occur at regular intervals.

7. 26 Fe
55.845 27 Co 28 Ni
Moseley 51 Sb 52 Te
127.60 53 I

8. Organization Of The Periodic Table

9. How are the Elements Classified?
Elements can be classified as metals, nonmetals, or metalloids (semimetals).
Metals
Nonmetals
Metalloids (Semimetals)
Located to the left of the staircase on the periodic table
Tend to lose electrons and form positive (+) ions to look like a noble gas
Have a shiny appearance (luster)
Good conductors of electricity
Located to the right of the staircase on the periodic table
Tend to gain electrons and form negative (-) ions to look like a noble gas
Dull in appearance
Poor conductors of electricity
Located along the staircase on the periodic table
May form + or – ions gas OR share electrons to look like a noble gas
Appearance varies
Intermediate conductor of electricity

10. Alkali Metals Located in group 1 on the periodic table
 Highly reactive metallic elements that rapidly react with water to form hydrogen and an alkaline solution
 Must be stored in mineral oil to prevent it from reacting with water vapor in the air and EXPLODING!
 Tend to lose the 1 electron on its valance (outermost) shell to look like a noble gas

11. Alkaline Earth Metals Located in group 2 on the periodic table
Reactive metallic elements but not as reactive as the alkali metals
Tend to lose the 2 electrons on its valance (outermost) shell to look like a noble gas

12. Alkali & Alkaline Earth Metal Videos

13. Transition Metals Located in the center section of the periodic table
This group contains precious metals (Cu, Au, Ag, Pt) 
This group contains magnetic elements (Fe, Ni, Co)
Tends to lose electrons (amount varies) on its valance (outermost) shell to look like a noble gas 
Electrons usually fill the d subshells

14. Halogens  Located in group 17 on the periodic table
Nonmetallic group of elements that has the nickname “Salt Formers” 
Tend to gain 1 electron on its valance (outermost) shell to look like a noble gas
Most reactive nonmetals on the periodic table

15. Noble Gases Located in group 18 on the periodic table
 Nonmetallic group of elements that are characterized by low reactivity because of their full valance shell (all noble gases have 8 valance electrons with the exception of He) 
Every element on the periodic table wants to look like a noble gas because of its stability achieved by having a full valence shell

16. Halogen & Noble Gases Videos

17. Lanthanides Elements with the atomic numbers 58-71
Rare earth elements that occur in minerals
Electrons fill the 4f subshell

18. Actinides Elements with atomic numbers 90-103
Rare earth elements that are radioactive (because they are very heavy) 
Electrons fill the 5f subshell

19. Lanthanides & Actinides Videos

20. Periodic Table Trends
Group # tells how many valence electrons are in the outer shell of an atom
This only applies to groups 1-2 and groups Remember for groups 13-18, subtract 10 from the group # to tell how many valence electrons are present
Period # tells how many energy levels an atom has and where the valence shell is located

21. Visuals to Group & Period # Trends
*Group # indicates the # of
valence electrons in outer
shell of the atom

22. The carbon atom
has 2 energy levels
because it Is located
on period 2
The cesium atom
has 6 energy
levels because it
is located on
period 6

23. What does Atomic Radius Measure?
Atomic radius is calculated by measuring the distance between the nuclei of 2 identical bonded atoms and then cutting this distance in half
3.72 angstrums/2 = 1.86 angstrums
3.72 angstrums

24. Atomic Radii Group Trend
Atoms become larger as one moves down a group (column)
As you move down a group, more e-’s lie between the nucleus and the highest energy level/outermost valence shell.
These inner shell e-’s SHIELD the outermost valence e-’s from the positively charged nucleus

25. Atomic Radii Group Trend
Versus
Cesium
Cesium has a larger atomic radius because there is more energy levels which
SHIELD the effect of the positive nucleus from attracting to the negatively charged
outermost electrons
Lithium has a smaller atomic radius because there are less energy levels, resulting in
less SHIELDING between the positively charged nucleus and the outermost valence
electrons. The nucleus attracts and pulls the valence electrons closer to it, resulting
in a smaller atom

26. Atomic Radii Period Trend
Atoms become smaller as one moves across a period from left to right.
As you move across a period, the inner shells of the atom do not get any “thicker.”
More e-’s are added to the outer shell of the atom and more protons are added to the. nucleus
As a result, the nuclear charge is greatly increased and the greater attraction between the protons and electrons shrinks the atom.

27. Atomic Radii Period Trend
Lithium Z=3 has a larger atomic radius than fluorine Z=9 because of the
increased nuclear charge as we move across period 2. A greater attraction of
electrons and protons causes the atoms to shrink in size when moving across
a period.

28. Summary of Atomic Radii Trend

29. Cation & Anion Trend
Negative ions (anions) are larger than their parent atoms.
A larger radius will result from greater repulsion on the valence shell due to the addition of e-’s.
Positive ions (cations) are smaller than their parent atoms.
A smaller radius will result from less repulsion on the valence shell due to the removal of e-’s.

31. What does Ionization Energy Measure?
Ionization energy measures the quantity of energy needed to remove an electron from an atom or ion

32. Ionization Energy Group Trend
As we go down a group, IE decreases due to the SHIELDING effect.
It is easier and doesn’t take as much energy to remove an electron from larger elements because of the minimal pull from the nucleus and larger distance to the outermost electrons.
The inner electrons shield or protect the valence electrons from the attraction from the nucleus.

33. Ionization Energy Period Trend
As we go across a period, IE increases due to an increase in nuclear charge.
It is harder to remove or pull away an electron from a smaller atom because of the close proximity of the valence electrons to the nucleus and the strong attraction that results.

34. What does Electron Affinity Measure?
Electron affinity measures the amount of energy that is released when electrons are added to the outer shell of an atom or ion.
The greater the negative EA, the easier it is for that atom or ion to accept an atom

35. Electron Affinity Group Trend
As we go down a group, EA decreases due to electron SHIELDING.
There is less attraction between the nucleus and the valence electrons due to the distance between them.
It is harder to add and hold another electron to the outermost shell because of the weak pull from the nucleus.
Only a small amount of energy (if any) would be released because most of it would be used to hold on to the added electrons

36. Electron Affinity Period Trend
As we go across a period, EA increases because of greater nuclear charge.
The addition of protons to the nucleus and electrons to the valence shell causes great attraction to one another.
It becomes easier to add an electron and a larger amount of energy will be released in the process

37. What does Electronegativity Measure?
Electronegativity is the ability of an atom to attract an electron from another element during bonding.
It is closely related to electron affinity.

38. Electronegativity Group Trend
As we go down a group, the ability of atoms or ions to attract another electron and form a bond will decrease.
There is less attraction between the nucleus and the valence electrons due to the distance between them.
It is harder to add and hold another electron to the outermost shell because of the weak pull from the nucleus.

39. Electronegativity Period Trend
As we go across a period, the ability of atoms or ions to attract another electron and form a bond will increase.
The addition of protons to the nucleus and electrons to the valence shell causes great attraction to one another.
It becomes easier to add an electron and form a bond due to the increase in nuclear charge.