PP - Unit 10 Notes_Periodic Table (1).ppt

History of the Periodic Table
Antoine Lavoisier (1743 – 1794)
– Published Elements of Chemistry in 1789
• Included a list of “simple substances”
(which we now know to be elements)
• Formed the basis for the modern list
of elements
– Only classified substances as
metals or nonmetals

Johann Döbereiner (1780 – 1849)
– Classified elements into “triads”
• Groups of three elements with related
properties and weights
• Began in 1817 when he realized Sr was
halfway between the weights of Ca and
Ba and they all possessed similar traits
– Döbereiner’s triads:
• Cl, Br, I  S, Se, Te
• Ca, Sr, Ba  Li, Na, K

John Newlands (1837 – 1898)
– Law of Octaves (1863)
• Stated that elements repeated their
chemical properties every eighth element
• Similar to the idea of octaves in music

Dmitri Mendeleev (1834 – 1907)
– Russian chemist (“The father of the periodic table”)
– Arranged elements based on accepted
atomic masses and properties that he
observed
– Listed elements with similar
characteristics in the same family/group
• Left blank spots for predicted
elements (Ted-Ed Video)

Property Mendeleev’s
Prediction for
“eka-silicon” in
1871
Observed
Properties of
Germanium
(discovered in
1886)
Atomic
Weight
72 72.59
Density
(g/cm3
)
5.5 5.35
Melting Point
(°C)
High 947
Color Dark gray Grayish white
Formula of XO GeO

Henry Moseley (1887 – 1915)
– English physicist
– Arranged elements based on increasing
atomic number
• Remember: atomic number = # of p+
in nucleus
– Periodic table looked similar to
Mendeleev’s design since as
atomic number increases, so
does the atomic mass

Periodic Law
Periodic Law
• Periodic – occurring at regular intervals
– Relates to trends on the periodic table of
elements
• Modern Periodic Law
– When elements are arranged in order of
increasing atomic number, there is a
periodic repetition of their properties
• Just like Mendeleev suspected!!

Reading the Periodic Table
•Periods - “Horizontal Rows”
•Groups (or Families) - “Vertical Columns”

• Valence electrons are periodic!
• Notice the similarities
– Ex.) Write the noble gas configurations for:
• F [He]2s2
2p5
7 valence electrons
• Cl [Ne]3s2
3p5
7 valence electrons
• Br [Ar]4s2
3d10
4p5
7 valence electrons
• I [Kr]5s2
4d10
5p5
7 valence electrons
– GROUPS have similar valence electron
configurations!

Groups of Elements
Groups of Elements
• Group 1 = Alkali Metals
– Located in Group 1 (except Hydrogen)
– Extremely reactive
• Want to lose 1 e- to become “noble gas-like”
• Group 2 = Alkaline Earth Metals
– Also very reactive
– Both Group 1 & 2 occur naturally as
compounds not elements

Groups of Elements
Groups of Elements
• Group 17 = Halogens
– Very active nonmetals
• Want to gain 1 e-
to become like a noble
gas

Groups of Elements
Groups of Elements
• Group 18 = Noble Gases
– Sometimes called “inert gases” since they
generally don’t react
• Mainly true, but not always (Kr, Xe will
react sometimes)
• Have a full valence shell (8 e-
)
Mythbusters Noble Gas Demo

Groups of Elements
Groups of Elements
• Transition Metals
– Located in the center of the Periodic Table
– 10 elements wide (“d” orbitals)
– Semi-reactive, valuable, crucial to many life
processes
• Lanthanides and Actinides
– Located at the bottom of the Periodic Table
– 14 elements wide (“f” orbitals)
– Some are radioactive, though not all
– Lanthanides = Period 6 (4f)
– Actinides = Period 7 (5f)

Alkali Metals =
Alkaline Earth Metals =
Transition metals =
Metalloids = Lanthanides =
Halogens = Actinides =
Noble Gases =

• Electronegativity
– Ability of an atom to pull e-
towards itself
– Increases going up and to the right
• Across a period  more protons in nucleus =
more positive charge to pull electrons closer
• Down a group  more electrons to hold onto =
element can’t pull e-
as closely
Periodic Properties & Trends

• Atomic Radius
– Distance between the nucleus and the
furthest electron in the valence shell
– Increases going down and to the left
• Down a group  more e-
= larger radius
• Across a period  elements on the right can
pull e-
closer to the nucleus (more
electronegative) = smaller radius
• *Remember*
– LLLL  Lower, Left, Large, Loose

• Ionization Energy
– Energy required to remove an e- from the
ground state
– 1st
I.E. = removing 1 e-
, easiest
– 2nd
, more difficult
– 3rd
, even more difficult
• Ex.) B --> B+
+ e- I.E. = 801 kJ/mol
• Ex.) B+
--> B+2
+ e- I.E.2 = 2427 kJ/mol
• Ex.) B+2
--> B+3
+ e- I.E.3 = 3660 kJ/mol

Ionization Energy
• Increases going up and to the right
– Down a group  more e-
for the nucleus
to keep track of = easier to rip an e-
off
– Across a period  elements on the right
can hold electrons closer (more
electronegative) = harder to rip an e-
off

• Metallic Character
– How “metal-like” an element is
• Metals lose e-
– Most Metallic: Cs, Fr
– Least: F, O
– Increases going down and to the left
Think about where the metals & nonmetals are
located on the periodic table to help you remember!

• Ionic Radius
– Radius of an atom when e-
are lost or
gained different from atomic radius
– Ionic Radius of Cations
• Decreases when e- are removed
– Ionic Radius of Anions
• Increases when e- are added

Sizes of Ions
• CATIONS are SMALLER than the
atoms from which they are formed.
• Size decreases due to increasing he
electron/proton attraction.
Li,152 pm
3e and 3p
Li +, 78 pm
2e and 3 p
+

Sizes of Ions
• ANIONS are LARGER than the
atoms from which they are formed.
• Size increases due to more
electrons in shell.
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
-

Overall Periodic Trends
Property Group Trend Period Trend
Atomic Radius
Increases going
down
Increases to the
left
Ionization Energy
Increases going up Increases to the
right
Electronegativity
Increases going up Increases to the
right
Metallic Character
Increases going
down
Increases to the
left

Practice:
Rank the elements from lowest to highest…
Electronegativity - C, F, Mg
Atomic Radius - Ir, Re, Bi
Metallic Character - Rb, Mn, P
Ionization Energy - B, Ga, In

PP - Unit 10 Notes_Periodic Table (1).ppt

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