PERIODIC TRENDS
Unit 5, Periodic Trends and Electron
Configuration
s, p, d, f Blocks
Remember in the lab we did with electron
configuration, there was a pattern on the
periodic table.
This pattern is the basis for modern chemistry.
While protons define an atom, electrons
define the atom’s reactivity.
Why does this matter?
Valence electrons are those electrons found in
the highest energy level in an atom.
All elements in a group have the same number
of valence electrons in their highest energy
level.
Problems
Given each of the following valence electron
configurations, determine which block of the
periodic table the element is in:
a. s2p4
c. s1
b. s2d1
d. s2p1
Describe how each are related:
Group number and
number of valence
elections:
Principal energy level of
valence electrons and
period number:
Group number in Roman
numerals (with A) is the
same as the number of
valence electrons.
Period number is the
energy level in which
you find valence
electrons for elements
in that period.
Electron Configurations
• Don’t have to write out the entire electron
configuration.
• There is a short-cut:
– Keeps focus on valence electrons
– An atom’s inner electrons are represented by the
symbol for the nearest noble gas with a lower atomic
number.
K: [Ar]4s1
Electron Configurations
For the element Phosphorus
-- 15 electrons
2
3
2
2
6
2
3
1s 2s 2p 3s3s3p3p
P: [Ne]
Must be a
Noble gas
(One just before
Element)
Electron Configurations
Let’s do a couple more:
Ba: [Xe] 6s2
Hg: [Xe] 6s2 4f14 5d10
2
3
V: [Ar] 4s 3d
Periodic Properties
• An element’s properties can go hand in hand
with electron arrangement
• We can use an element’s location on the PT to
predict many properties.
– Atomic radius
– Ionic Size
– Ionization energy
– Electronegativity
Atomic Radius
• The radius of an atom is defined by the edge of
its last energy level.
– However, this boundary is fuzzy
• An atom’s radius is the measured distance
between the nuclei of 2 identical atoms
chemically bonded together - divided by 2.
Atomic Radius
• As we examine atomic radius, we see a gradual
decrease in atomic size from left to right across
the PT.
– As e- are added in the same energy level, they are
pulled closer to the highly positive nucleus.
Atomic Radius
• Atomic radius increases going down a family.
• The change in atomic radii going down is due to
e- shielding
– As we move down a group on the PT, we add an
energy level of e-‘s pushing the valence e’s farther
away from the nucleus.
Periodic Properties
• How does the size of an atom change when
electrons are added or removed?
As an atom loses 1 or
more electrons
(becomes positive), it
loses an energy layer
therefore, its radius
decreases.
Ion Radius
• How does the size of an atom change
when electrons are added or
removed?
As an atom gains 1 or
more electrons
(negative), it fills its
valence energy level,
therefore, its radius
increases.
Positive Ion Radius Properties
• Elements in a group tend to form ions of the
same charge.
– Modeled by electron configurations.
K: [Ar]
4s
Loses 1
electron
[Ar]
Wants a full set of e-
4s
Negative Ion Radius
O: [He]
2s2
Wants a complete set
2p4
Gains
2 electrons
[He]
Periodic Trend of Ionic Charges
Ion Radius Trends
Ion radius increases going down a group due to
added energy levels.
Ion radius across a period depends on the
charge of the ion; positive ions are smaller
than their neutral atoms (due to missing
energy level) and negative ions are larger than
their neutral atoms (due to increased
repulsion of the added electrons.
Tend to lose
electrons to
become
positive
Tend to gain
electrons to
become
negative
Ionization Energy Properties
The energy needed to remove an atom’s
electron.
• If the e-s are held strongly the atom will have a
high ionization energy
• The first electron is always the easiest to
remove.
• The more valence electrons, the more difficult
to remove one.
• Noble gas configurations are exceptionally high.
Ionization Energy
•IE decreases down a group due to the
increasing number of energy levels
shielding the valence electrons.
•IE increases left to right across a period
because as an atom gets closer to an
octet, the attraction for valence
electrons increases
– There is generally a large jump in energy
necessary to remove additional electrons
from the atom.
The amount of energy required to remove a
2p e– (an e- in a full sublevel) from a Na ion is
almost 10 times greater than that required
to remove the sole 3s e-
Electronegativity
Electronegativity reflects the ability of an
atom to attract electrons in a chemical
bond.
--Decreases going down a group due to
shielding from added energy levels.
–Increases left to right across a period
due to increased strength of attraction
between nucleus and valence electrons.
Electronegativity correlates to an atom’s
ionization energy.
Electronegativity
Electronegativity
Electron Dots
Draw the symbols of atoms with dots to
represent the valence-shell electrons for
period 3.
1A 2A

Na

Mg
3A
4A
5A

Al

 Si 


P

6A


S:

7A

:
Cl

8A

:Ar
:

Bonding
• Atoms are generally found in nature in
combination held together by chemical
bonds.
– A chemical bond is a mutual electrical
attraction between the nuclei and outer
electrons of different atoms that binds the
atoms together so that they behave as one
unit.
• There are two types of chemical bonds:
ionic and covalent.
Introduction to Ionic Bonding
• Two atoms in an ionic bond transfer
electrons from one element to another.
– Generally, electrons move from a metal to a
nonmetal.
– The electron is transferred from the atom
with the low electronegativity to the atom
with the high electronegativity.
Introduction to Covalent Bonding
• Two atoms involved in the covalent bond
share electrons in order to achieve the
arrangement of a noble gas.
–Generally between two nonmetals.
–Because electronegativities are similar
so neither wants to give up electrons.
metal w/nonmetal = usually ionic
nonmetal w/nonmetal = usually covalent
Metallic Bond
• The electron sea model proposes that all the
atoms in a metallic solid contribute their
valence electrons to form a “sea” of electrons.
• A metallic bond is the attraction of a metallic
ion for the “sea” of delocalized electrons.
Metals Form Alloys
Metals do not combine chemically with metals.
They form alloys which are solutions of a metal in
a metal.
•Delocalized electrons are free to move heat or
electricity easily through metals.
•Number of delocalized electrons and strength of
metallic bond determine melting point.
Types of Alloys
There are two kinds of alloys.
Substitutional alloys are made of elements of
similar size (sterling silver, brass, pewter).
Interstitial alloys are made of elements with
different size atoms so that little ones fit
between larger ones (carbon steel).
(stop)
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