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The Periodic Table
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Dmitri Mendeleev (1871)
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Organized the first periodic table
He arranged the periodic table in order of
increasing atomic mass
He placed the elements in columns based
upon their similar properties.
He left blank spaces where he thought
undiscovered elements would go, based on
their properties
Henry Moseley (1911)
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Arranged periodic table in order of increasing
atomic number
Periodic law = properties of elements are
periodic functions of their atomic number
The Periodic Table
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Rows (horizontal) are called periods
-indicated by using a number 1-7
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Columns (vertical) are called groups
or families
Groups
- numbered with an A or B
– The A group elements 1A 8A are the
main group elements (the s and p blocks)
= Representative Elements
– The B group elements are the d-block and
f-block
– Groups can also be numbered 1-18
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Metals
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left of staircase line
- 1A (or 1) are called the alkali metals;
highly reactive except H
- 2 A (or 2) are called the alkaline earth
metals
- d-block are called transition metals
- f-block are called inner-transition metals
Nonmetals
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right of staircase line
- 7A (or 17) are called the halogens; very
reactive
- 8A (or 18) are called the noble gases;
very unreactive
Metalloids
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Elements bordering the staircase line
-have properties of both metals and non
metals.
Using the Periodic Table
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All the elements in the same group have the
same number of valence electrons
Group number (1A-8A) = number of valence
e- **this only works for the Representative
Elements (s & p block)**
Elements with the same number of valence
e- have similar properties
level of valence e- = period
number of element
 Energy
Example: Aluminum is in period 3;
valence e- are in 3rd energy level
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Important Concepts
 Nuclear
Charge = positive
charge of the nucleus
 Shielding Effect = blocking of
valence e- from nucleus by
inner e12
Periodic Trends
Periodic Trends – trends found on the periodic
table that compare properties of the
elements
 Atomic Radius (Atomic Size)
 Ionic Radius (Ionic Size)
 Ionization Energy
 Electronegativity
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Atomic Radius
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One half the distance between the nuclei of 2 like
atoms in a diatomic molecule
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You can’t measure the radius exactly because the
electron cloud doesn’t have a definite end.
Atomic Radius Trend
 Increases
down a group↓ due to
increase in occupied energy levels
 Decreases across a period→ due
to increase in nuclear charge
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Atomic Radius Example
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Atomic Radius
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Ions
 Ion=
atom that has gained or
lost e– Metals lose e- and form positive
ions (cations)
– Nonmetals gain e- and form
negative ions (anions)
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Ionic Radius Trend
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No general trend because some
atoms gain e-, while others lose e1.
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3.
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Write e- configuration after gaining/losing
eCompare energy levels
If energy levels are equal, compare
nuclear charge
Ionic Radius Example
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Ionization Energy
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Energy required to remove an electron from
an atom
A different amount of energy is required each
time an e- is removed
1st ionization energy – remove 1st electron
2nd ionization energy – remove 2nd electron
3rd ionization energy – remove 3rd electron
Ionization Energy (cont’d)
 Every
element wants an econfiguration like a noble gas (full
outer s & p)
 The easier it is to lose an e-, the
lower the ionization energy is
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Ionization Energy Trend
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Decreases down a group ↓ – greater
shielding effect
Increases across a period → – greater
nuclear charge
Ionization Energy Example
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Ionization Energy
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Electronegativity
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The tendency for the atoms to attract
electrons
Trend :
– Decreases down a group ↓
– Increases across a period →
Flourine is the Most Electronegative Element
Why aren’t noble gases included in
electronegativity?
Electronegativity
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