Transcript History of Periodic Table

History of Periodic
Table
Chapter 5
History

1860s – 60 elements discovered
– Cannizzaro - agreed on method to
measure atomic mass
– Search for relationships between
properties of elements
Dimitri Mendeleev
Organized elements by increasing
atomic mass
 Noticed chemical and physical
properties followed trend, or pattern 
Periodic

Mendeleev’s Table
Henry Moseley
Worked with Rutherford looking at linespectras
 Noticed better pattern when elements
were organized by increasing atomic #
 Periodic Law: the physical and
chemical properties of elements are
periodic functions of their atomic #s

Regions of Periodic
Table
Group Project
Main Group Elements
s and p block
elements
Group 1A are the alkali metals
 Group 2A are the alkaline earth metals

Group 7A is called the Halogens
 Group 8A are the noble gases

The group B are called the
transition metals
Top: Lanthanide Series
Bottom: Actinide Series
Periodic Properties
Atomic Radii (Atomic Size)

Def: half the distance between the
nuclei of identical atoms that are
bonded together
}
Radius
Atomic Radii - Group trends
As we go down a
group
 Another energy
level…
 So the atoms get
bigger.

H
Li
Na
K
Rb
Atomic Radii - Periodic Trends
Go across a period the radius gets
smaller.
 Same energy level.
 More nuclear charge.
 Outermost electrons pulled in closer

Na
Mg
Al
Si
P
S Cl Ar
Ionization Energy (IE)

An e- can be removed from any atom if
there is enough energy
A + energy  A+ + e-
Ion: atom or group of bonded atoms
that has a (+) or(-) charge
 Process that results in ion formed is
ionization

Valence Electrons
Def: The e- available to be lost, gained
or shared to form chemical compounds
 e- found in the outermost s and p
sublevels

Ionization Energy (IE)
Def: the energy needed to remove one
e- from an atom (IE1 – first IE)
 Atoms with HIGH ionization energy 
hold on tight to their electrons

IE – Group Trends
As you go down a group IE decreases
 Electron further away from nucleus
 Less attraction to nucleus, easier to
take e
IE – Periodic Trends
IE generally increases from left to right
 Increasing nuclear charge
 More nuclear charge holds on tight to e Exact opposite of atomic radius

IE2 and IE3
Energy required to remove additional e Energies keep getting higher and higher
 e- that are left are being held closer to
nucleus  harder to remove
 Pg. 155
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Ionic Radii (Ionic Size)

Cation: positive ion
– Always smaller than atom
– Lost e-, now nucleus pulling in more on
remaining e-s

Anion: negative ion
– Always bigger than atom
– Gaining e-, now e- are crowded and spread
out (repulsion of like charges)
Ionic Radii – Group Trends
Same as Atomic Radii
 More energy levels as go down  size
increases

Ionic Radii – Periodic Trends
2 sections
 Metals on LEFT make CATIONs
 Nonmetals on RIGHT make ANIONS
 Cations (1A – 4A) Anions (5A – 8A)

– Decrease as go across (L-R) due to
increase nuclear charge
Electronegativity
Valance e- are involved in forming
bonds
 Some atoms in a chemical bond attract
the valance e- more than the other (tug
of war)
 Linus Pauling – electronegativity –
measure of the ability of an atom in a
chemical compound to attract e- from
another atom in the compound

Electronegativity – Group Trends
Tend to decrease down a group or
remain about the same
 Noble gases are NOT assigned
electronegativities

Electronegativity – Periodic
Trends
Tend to increase as you go across the
table
 F – most electronegative
 Fr – least electronegative
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