PERIODICITY
Table of Contents
Electron Configurations pg. 3
History of the Periodic Table pg. 11
Periodic Properties pg. 16
Atomic Radius pg. 17
Ionic Radius pg. 24
Ionic Charge pg. 27
Ionization Energy pg. 32
Electron Affinity pg. 38
Octet Rule pg. 40
Electronegativity pg. 43
Other Periodic Trends pg. 45
Questions pg. 49
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
1s22s22p63s23p3
P: [Ne]
Must be a
Noble gas
(One just before
Element)
Electron Configurations
Let’s do a couple more:
Ba: [Xe]
2
6s
Hg: [Xe] 6s2 4f14 5d10
V: [Ar] 4s2 3d3
Exceptions to the order of filling
Electron Configurations
• The chemistry of an atom occurs at
the set of electrons called valence
electrons
• The valence electrons are electrons
in an atom’s highest energy level.
– For the Group – A elements, it is the
outermost s & p e- of the atom.
– Specifically the 2 s electrons + 6 p
electrons (octet electrons)
• The arrangement of the valence elead to the element’s properties.
History of the Periodic Table
• 70 elements had been discovered by
the mid-1800’s, but until Dmitri
Mendeleev, no one had a come with
a way to organize the elements.
– Mendeleev came up with the first
working system of filing the elements.
• He listed the elements in columns in
order of increasing atomic mass, and
then put columns together that were
similar
History of the Periodic Table
• Mendeleev left gaps in the table since
there were no current elements that
seemed to fit those spots
– Those elements were eventually
discovered and they fit perfectly into
an open spot.
• The 1st scientist that set the table in
order of atomic number was Henry
Moseley
History of the Periodic Table
• The modern PT is arranged by
increasing atomic number
– Increases from left to right, and top to
bottom
• This establishes the periodic law
– When the elements are arranged in
order of increasing atomic #, there is
a periodic repetition of their phys &
chem properties
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
Electron affinity
Electronegativity
Ionization energy
Ionic Size
Periodic Properties
• 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.
Periodic Properties
• As we examine atomic radius from left
to right across the PT we see a gradual decrease in atomic size.
– As e- are added to the s and p
sublevels in the same energy level,
they are gradually pulled closer to the
highly positive nucleus
• The more e-’s in the atom the less
dramatic this trend looks
Periodic Properties
• The change in atomic radii across the
PT is due to e- shielding or to the
effective nuclear charge
– As we move across
the PT we are adding
e- into the same general vol. in which case
they will shield or
interact with each
other (repulsion)
Periodic Properties
– We are also adding protons into the
nucleus which increases the p+-einteraction (attraction)
• So the nucleus gains strength while
the e- aren’t gaining much distance,
so the atom is drawn in closer and
closer to the nucleus.
– Decreasing the overall radius of the
atom
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 a layer
therefore, its radius
decreases.
Periodic Properties
• 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 layer,
therefore, its radius
increases.
Periodic 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
Periodic Properties
O: [He]
2s2
Wants a complete set
2p4
Gains
2 electrons
[He]
Periodic Trend of Ionic Charges
The Transition Elements are almost
unpredictable, and sometimes have
more than one possible charge -- due to
d orbitals --
Tend to lose
electrons to
become
positive
Tend to gain
electrons to
become
negative
Periodic Properties
• Another periodic trend on the table is
ionization energy (a.k.a. potential)
– Which is the energy needed to
remove one of an atoms e-s.
– Or a measure of how strongly an
atom holds onto its outermost e-s.
• If the e-s are held strongly the atom
will have a high ionization energy
Periodic Properties
• The ionization energy is generally
measured for one electron at a time
• You can also measure the amount of
energy needed to reach in and pluck
out additional electrons from atoms.
– 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-
Periodic Properties
• There is simply not enough energy
available or released to produce an
Na2+ ion to make the compnd NaCl2
– Similarly Mg3+ and Al4+ require too
much energy to occur naturally.
• Chemical formulas should always
describe compounds that can exist
naturally the most efficient way
possible
Periodic Properties
• Another periodic trend dealing with
an e- is electron affinity
– Which is a measure of the ability of
an atom to attract or gain an electron.
• Atoms that tend to accept an e- are
those that tend to give a neg. charge.
– The closer to a full outer shell an
atom has, the higher the affinity
(more neg. the measurement)
Periodic Properties
• An atoms ability to lose an e- or gain
an e- can be used to understand the
Octet Rule
• Octet Rule: atoms tend to gain, lose,
or share electrons in order to acquire
a full set of valence electrons.
– 2 e- in the outermost s sublevel + 6 e–
in the outermost p sublevel= a full
valence shell
Periodic Properties
• Electronegativity is a key trend.
– It reflects the ability of an atom to
attract electrons in a chemical bond.
– F is the most electronegative
element and it decreases moving
away from F.
• Electronegativity correlates to an
atom’s ionization energy and electron
affinity
BOILING POINT & MELTING POINT VS. ATOMIC NUMBER
INCREASES
INCREASES
Questions
•
1.) Write out the noble gas electron configuration of Tungsten, Tin, Selenium,
Uranium, and Silver.
•
2.) What are the valence electrons of the elements from the previous question?
•
3.) Describe how the modern periodic table is arranged.
•
4.) Arrange Lithium, Potassium, and Cesium in order from smallest to largest
atomic radius and briefly explain your reasoning for the ordering.
•
5.) Arrange Sulfur, Sodium, and Aluminum in order from smallest to largest
atomic radius and briefly explain your reasoning for the ordering.
•
6.) Which has a larger radius, neutral bromine or bromide ion and why?
•
7.) Which has a larger radius, neutral calcium or calcium ion and why?
Questions
•
8.) Name three elements that typically have a charge of -1, +2, and can take
more than one charge.
•
9.) Arrange Chlorine, Sodium, Argon, and Bromine from smallest to largest
Ionization energy.
•
10.) Why does KBr2 not exist in nature?
•
11.) Explain why Ionization Energy, Electron Affinity, and Electronegativity all
follow the same periodic trend.
•
12.) Arrange Selenium, Platinum, Chlorine, and Nickel from least dense to
most dense.
•
13.) Hypothesize why the periodic trends for Atomic Radius and Electron
Affinity are opposite to each other.