• 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.

• 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 13-18 on your periodic table)

26
Fe
55.845
28
Ni
58.6934
27
Co
58.9932
51
Sb
121.760
53
I
126.9045
Mendeleev
52
Te
127.60

• 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.

26
Fe
55.845
27
Co
58.9332
28
Ni
58.6934
51
Sb
121.760
52
Te
127.60
Moseley
53
I
126.9045

• Elements can be classified as metals, nonmetals, or metalloids (semimetals).
Metals Nonmetals
 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
Metalloids
(Semimetals)
 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


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


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

• http://www.youtube.com/watch?v=sNdijknRx fU
• http://www.youtube.com/watch?v=DFQPnHk
QlZM


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


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


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

• http://www.youtube.com/watch?v=u2ogMUD
Baf4
• http://www.youtube.com/watch?v=QLrofyj6a
2s
• http://www.youtube.com/watch?v=3TYuym2j qjM

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

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

• http://www.youtube.com/watch?v=IFmAhhia m9g

• Group # tells how many valence electrons are in the outer shell of an atom
• This only applies to groups 1-2 and groups 13-
18. 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

*Group # indicates the # of valence electrons in outer shell of the atom

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

• 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

• 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

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

• 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.

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.

• 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.

• Ionization energy measures the quantity of energy needed to remove an electron from an atom or ion

• 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.

• 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.

• 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

• 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

• 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

• Electronegativity is the ability of an atom to attract an electron from another element during bonding.
• It is closely related to electron affinity.

• 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.

• 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.