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The Periodic Table
History
Electron configurations
●The organization of the modern table
●Some different types of tables
●Periodic trends in the table
●Some example problems
●
The Periodic Table
Dmitri Mendeleev 1834-1907
This photograph is in the public domain.
The Periodic Table
ordered elements by atomic mass
●
saw a repeating pattern of chemical and physical
properties
●
Periodic Law – When the elements are arranged in
order of increasing atomic mass, certain sets of
properties recur periodically
●
put elements with similar properties in the same
column
●
The Periodic Table
used patterns to predict properties of
undiscovered elements
●
where atomic mass order did not fit other
properties, he re-ordered by other properties
ie. Te & I
●
The Periodic Table
This diagram is in the public domain.
The Periodic Table
Predicted Properties of New Elements
Property
Ekaaluminium
atomic mass
68
density (g/cm³)
6.0
melting point (°C) Low
oxide
Ea2O3
chloride
Ea2Cl6
Gallium
69.72
5.904
29.78
Ga2O3
Ga2Cl6
The Periodic Table
Predicted Properties of New Elements
Property
Ekasilicon
atomic mass
72
density (g/cm³)
5.5
melting point (°C) high
oxide
EaO2
chloride
EaCl4
germanium
72.61
5.35
947
GeO2
GeCl4
The Periodic Table
Electron Configurations
The four quantum numbers n, l, ml and ms enable
us to label completely an electron in any
orbital in any atom. It is as though the set of
quantum numbers for an individual electron are
an 'address' for the electron. For example, the
set of quantum numbers for a 2s electron would
be (n, l, ml, ms) 2, 0, 0, +1/2 or -1/2.
The Periodic Table
Electron Configurations
The hydrogen atom is a particularly simple
system since it contains only one electron. The
electron may reside in the 1s orbital (the
ground state) or it may be found in some higher
energy orbital (an excited state). For manyelectron atoms, however, we must know the
electron configuration of the atom ... in other
words, how the electrons are distributed among
the various atomic orbitals.
The Periodic Table
Electron Configurations
The Aufbau Principle: electrons occupy orbitals
in a way that minimizes the energy of the atom.
In other words, electrons fill the atomic
orbitals starting with 1s and proceeding to the
next highest energy orbital
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
The Pauli Principle: two electrons with the same
spin cannot occupy the same orbital. In other
words, no two electrons can have the same set of
quantum numbers (the same 'address').
Hund's Rule: the most stable arrangement of
electrons in subshells is the one with the
greatest number of parallel spins. In other
words, when degenerate orbitals are available
for filling, the electrons will occupy them
singly with the same spin.
The Periodic Table
Electron Configurations
Hydrogen – 1 electron
The Aufbau principle indicates that the
electron will go into the lowest energy
orbital, the 1s orbital:
H[
]
or 1s
1
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
Helium – 2 electrons
He[
]
or 1s
2
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
Lithium – 3 electrons
Li [
] [
]
2
or 1s 2s
1
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
Berylium – 4 electrons
Be[
]
[
]
][
][
2
or 1s 2s
2
Boron – 5 electrons
B[
]
1s
[
] [
2s
2p
]
2
2
or 1s 2s 2p
1
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
Carbon – 6 electrons
C[
]
1s
[
] [
2s
][
][
2p
]
2
2
or 1s 2s 2p
2
The Periodic Table
energy
Electron Configurations
© 2011 K. Brown
The Periodic Table
Electron Configurations
Nitrogen – 7 electrons
N[
] [
1s
] [
2s
2
2
3
2
2
4
][
]
or 1s 2s 2p
][
]
or 1s 2s 2p
][
2p
Oxygen – 8 electrons
O[
]
1s
[
]
2s
[
][
2p
The Periodic Table
Electron Configurations
Fluorine – 9 electrons
F[
] [
1s
] [
][
2s
][
]
or 1s 2s 2p
2
2
5
]
or 1s 2s 2p
2
2
6
2p
Neon – 10 electrons
Ne[
] [
1s
]
2s
[
][
][
2p
The Periodic Table
Electron Configurations
Most often we are interested only in the
outermost shell of electrons ... these are
the ones that are most involved in chemical
reactions. This is reflected in the way the
electron configuration is often written.
Instead of explicitly showing all electrons
we write the symbol for the previous noble
gas followed by the outer shell or valence
shell electron configuration.
The Periodic Table
Electron Configurations
Carbon
2
2
1s 2s 2p
2
2
[ He] 2s 2p
long form
compact form
valence electrons
Aluminum
2
2
2
6
2
1s 2s 2p 3s 3p
long form
1
2
1
[ Ne] 3s 3p
compact form
The Periodic Table
Electron Configurations
In the periodic table the left-most columns
include the alkali metals and the alkaline
earth metals. In these elements the valence
s orbitals are being filled
On the right hand side, the right-most block
of six elements are those in which the
valence p orbitals are being filled
The Periodic Table
Electron Configurations
In the middle is a block of ten columns that
contain the transition metals. These are
elements in which d orbitals are being
filled
Below this group are two rows with 14
columns. These are commonly referred to the
f-block metals or the lanthanides and
actinides. In these columns the f orbitals
are being filled
The Periodic Table
Electron Configurations
© 2013 K. Brown
The Periodic Table
Electron Configurations
The electron configurations of some of the
transition metal elements do not follow the
regular rules.
Chromium
2
4
[ Ar ] 4s 3d
expected
configuration
1
5
[ Ar ] 4s 3d
actual
configuration
The Periodic Table
Electron Configurations
The electron configurations of some of the
transition metal elements do not follow the
regular rules.
Copper
2
9
[ Ar ] 4s 3d
expected
configuration
1
10
[ Ar ] 4s 3d
actual
configuration
The Periodic Table
Electron Configurations
The electron configurations of ions are
generally straightforward with some exceptions.
For anions we simply add the extra electrons
using the regular rules for constructing
electron configurations.
For cations we remove electrons in reverse order
for the main group elements. For transition
metal elements however we remove electrons from
the orbital with the highest value of n first.
The Periodic Table
Electron Configurations
2
2
O : 1s 2s 2p
2
2
4
2-
→
2
2
6
-
2
6
2
5
→
6
2
6
1
Cl : 1s 2s 2p 3s 3p
2
O : 1s 2s 2p
K :1s 2s 2p 3s 3p 4s
2
2
6
2
6
1
2
2
6
2
Cl : 1s 2s 2p 3s 3p
→
Cr : 1s 2s 2p 3s 3p 4s 3d
5
+
2
2
6
6
2
K : 1s 2s 2p 3s 3p
→
3+
2
2
6
6
2
6
Cr :1s 2s 2p 3s 3p 3d
3
The Periodic Table
Electron Configurations
Species with the same electron configuration are
isoelectronic.
2-
2
2
O : 1s 2s 2p
-
2
2
F : 1s 2s 2p
2
6
2
Ne : 1s 2s 2p
+
2
6
2
6
Na :1s 2s 2p
2+
2
2
6
Mg : 1s 2s 2p
6
The Periodic Table
Shielding
Why do we see periodic physical and chemical
properties of the elements?
We need to first understand the concept of
shielding before this can be satisfactorily
explained.
The Periodic Table
Shielding
in a multi-electron system, electrons are
simultaneously attracted to the nucleus and
repelled by each other
●
outer shell electrons are shielded from full
strength of nucleus by the inner shell electrons
●
effective nuclear charge is net positive charge
that is attracting a particular electron in an
outer shell
●
The Periodic Table
Shielding
This diagram is licensed under Creative Commons.
The Periodic Table
Shielding
Z is nuclear charge, S is electrons in lower
energy levels
●
Z effective = Z− S
The Periodic Table
Valence Electrons and Periodic Properties
the electrons in all the subshells with the
highest principal energy shell are called the
valence electrons
●
electrons in lower energy shells are called core
electrons
●
The Periodic Table
Valence Electrons and Periodic Properties
the Group number corresponds to the number of
valence electrons
●
the length of each “block” is the maximum number
of electrons the sublevel can hold
●
the Period number corresponds to the principal
energy level, n, of the valence electrons
●
The Periodic Table
Valence Electrons and Periodic Properties
the number of valence electrons largely
determines the behavior of an element
●
since the number of valence electrons follows a
periodic pattern, the properties of the elements
should also be periodic
●
quantum mechanical calculations show that 8
valence electrons should result in a very
unreactive atom, an atom that is very stable –
and the noble gases, that have 8 valence
electrons are all very stable and unreactive
●
The Periodic Table
Valence Electrons and Periodic Properties
conversely, elements that have either one more
or one less electron should be very reactive –
and the halogens are the most reactive nonmetals
and alkali metals the most reactive metals
●
many metals and nonmetals form one ion, and the
charge on that ion is predictable based on its
position on the Periodic Table
●
●
Group 1A = +1, Group 2A = +2, Group 7A = -1,
Group 6A = -2, etc.
The Periodic Table
Valence Electrons and Periodic Properties
these atoms form ions that will result in an
electron configuration that is the same as the
nearest noble gas
●
The Periodic Table
Periodic Properties of the Elements
●
Atomic radius increases going down a group
Atomic radius decreases across period (left
to right)
●
The Periodic Table
Periodic Properties of the Elements
licensed under Gnu Free Documentation.
The Periodic Table
Periodic Properties of the Elements
The Periodic Table
Periodic Properties of the Elements
Ion size increases down the group
●
Cations smaller than neutral atom; Anions bigger
than neutral atom
●
Cations smaller than anions
●
Larger positive charge = smaller cation
●
Larger negative charge = larger anion
●
The Periodic Table
Periodic Properties of the Elements
radius of positive ions is smaller than that of
a neutral atom, which in turn is smaller than
that of a negative ion. Thus the sizes of the
isoelectronic atoms and ions have the following
order:
●
●
S2- > Cl- > Ar > K+ > Ca2+
The Periodic Table
Periodic Properties of the Elements
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The Periodic Table
Periodic Properties of the Elements
The Periodic Table
Periodic Properties of the Elements
The Periodic Table
Periodic Properties of the Elements
electron configurations that result in unpaired
electrons mean that the atom or ion will have a
net magnetic field. This is called paramagnetism
●
electron configurations that result in all
paired electrons mean that the atom or ion will
have no magnetic field – this is called
diamagnetism
●
The Periodic Table
Periodic Properties of the Elements
both Zn atoms and Zn2+ ions are diamagnetic,
showing that the two 4s electrons are lost
before the 3d
●
2
10
Zn:[ Ar ] 4s 3d
2+
0
10
Zn :[ Ar ] 4s 3d
The Periodic Table
Periodic Properties of the Elements
Since the number paired and unpaired electrons
is determined by the electron configuration and
the electron configuration is integrally linked
to the modern periodic table, diamagnetism and
paramagnetism are calculable properties of the
elements in the table.
2
7
Co :[ Ar ] 4s 3d
2
Mg :[ Ne] 3s
paramagnetic
diamagnetic
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
The ionization energy, or ionization potential,
is the energy required to completely remove an
electron from a gaseous atom or ion.
© 2013 K. Brown
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
The first ionization energy is the energy
required to remove one electron from the parent
atom. The second ionization energy is the energy
required to remove a second valence electron
from the univalent ion to form the divalent ion,
and so on. Successive ionization energies
increase. The second ionization energy is always
greater than the first ionization energy.
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
Ionization energies increase moving from left to
right across a period (decreasing atomic radius
and increasing effective nuclear charge).
Ionization energy decreases moving down a group
(increasing atomic radius, same effective
nuclear charge).
Group I elements have low ionization energies
because the loss of an electron forms a stable
octet of electrons.
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
The effective nuclear charge can be used to
explain the overall observations concerning the
periodicity of ionization energies. As Zeff
increases across a period the outer shell
electrons are held more tightly and require more
energy to remove from the atom.
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
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The Periodic Table
Periodic Properties of the Elements
Ionization Energy
Going down a group, the average distance that
the outer shell electrons are from the nucleus
increases as Zeff remains constant. Consequently,
the outer shell electrons are held less tightly
and can be removed with less energy than those
above them in the group.
The Periodic Table
Periodic Properties of the Elements
Ionization Energy
© 2013 K. Brown
The Periodic Table
Periodic Properties of the Elements
Successive Ionization Energies
removal of each successive electron costs more
energy
●
regular increase in energy for each successive
valence electron
●
large increase in energy when start removing
core electrons
●
The Periodic Table
Periodic Properties of the Elements
Successive Ionization Energies
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The Periodic Table
Periodic Properties of the Elements
Electron Affinity
Electron affinity reflects the ability of an
atom to accept an electron. It is the energy
change that occurs when an electron is added to
a gaseous atom.
© 2013 K. Brown
The Periodic Table
Periodic Properties of the Elements
Electron Affinity
Atoms with stronger effective nuclear charge
have greater electron affinity. Some
generalizations can be made about the electron
affinities of certain groups in the periodic
table.
The Periodic Table
Periodic Properties of the Elements
Electron Affinity
The Group IIA elements, the alkaline earths,
have low electron affinity values. These
elements are relatively stable because they have
filled s subshells. Group VIIA elements, the
halogens, have high electron affinities because
the addition of an electron to an atom results
in a completely filled shell.
The Periodic Table
Periodic Properties of the Elements
Electron Affinity
© 2013 K. Brown
The Periodic Table
Periodic Properties of the Elements
Metals and Non-metals
Metals
● malleable & ductile
● shiny, lusterous, reflect light
● conduct heat and electricity
● most oxides basic and ionic
● form cations in solution
● lose electrons in reactions - oxidized
●
The Periodic Table
Periodic Properties of the Elements
Metals and Non-metals
Nonmetals
● brittle in solid state
● dull
● electrical and thermal insulators
● most oxides are acidic and molecular
● form anions and polyatomic anions
● gain electrons in reactions - reduced
●
The Periodic Table
Periodic Properties of the Elements
Metals and Non-metals
© 2013 K. Brown
The Periodic Table
Periodic Properties of the Elements
Alkali Metals
atomic radius increases down the column
●
ionization energy decreases down the column
●
very low ionization energies
●
Reactivity increases down the column
●
electron affinity decreases down the column
●
melting point decreases down the column
●
The Periodic Table
Periodic Properties of the Elements
The Halogens
atomic radius increases down the column
●
ionization energy decreases down the column
●
very high electron affinities
●
reactivity decreases down the column
●
react with hydrogen to form HX (X – halogen) and
chlorides with metals
●
The Periodic Table
Periodic Properties of the Elements
The Noble Gases
atomic radius increases down the column
●
ionization energy decreases down the column
●
very unreactive
●
The Periodic Table
The Stories of the Elements
The prehistorical elements:
Copper
Iron
Lead
Platinum
Sulfur
Gold
Carbon
Mercury
Silver
Tin
The Periodic Table
The Stories of the Elements
Hydrogen – discovered by Cavendish, named
by Lavoisier
Oxygen – discovered by Scheele and
Priestly, named by Lavoisier
Nitrogen – discovered by Rutherford, named
by Lavoisier
Phosphorus – discovered by Brand in 1669.
The Periodic Table
The Stories of the Elements
The Periodic Table
The Stories of the Elements
Yttrium - discovered 1794 by Johann Gadolin
Terbium - discovered 1843 by Mosander
Erbium - discovered 1843 by Mosander
Ytterbium - discovered 1878 by Marignac
Gadolinium - discovered 1880 by Marignac
Holmium - discovered 1878 by Cleve & Soret
Thulium - discovered 1879 by Cleve
The Periodic Table
The Stories of the Elements
The Periodic Table
The Stories of the Elements
The Periodic Table
The Stories of the Elements
The Periodic Table
© 2013 K. Brown
The Periodic Table
The Periodic Table
The Periodic Table
The Periodic Table
The Periodic Table
The Periodic Table
The Periodic Table
The Periodic Table
Example Problems
1. Arrange the following in order of increasing
first ionization energy: Na, Cl, Al, S, Cs.
Cs < Na < Al < S < Cl
Example Problems
2. Arrange the following species in
isoelectronic pairs: O+, Ar, S2-, Ne, Zn, Cs+,
N3-, As3+, N, Xe.
O+, N
Ar, S2Ne, N3Zn, As3+
Cs+, Xe
5s24d105p6
1s22s22p3
[Ne] 3s23p6
1s22s23p6
[Ar] 4s23d10
[Kr]
Example Problems
3. Based on valences and typical anion and
cation charges, predict the likely formulas for
the reaction between francium and oxygen,
strontium and bromine, radium and selenium.
Fr2O
ie. Fr+ and O2-
SrBr2
Sr2+
and Br-
RaSe
Ra2+
and Se2-
Example Problems
4. Arrange the following atoms in order of
decreasing atomic radius: Na, Al, P, Cl, Mg.
Na > Mg > Al > P > Cl
Example Problems
5. Arrange the following isoelectronic species
in order of increasing a) radius and b)
ionization energy: O2-, Mg2+, F-, Na+
a) Mg2+ < Na+ < F- < O2b)
O2- < F- < Na+ < Mg2+
Example Problems
6. Element 120 (El) has just been discovered.
What will the most likely formula of its oxide
be? Will it be larger or smaller than the
previous element in the table?
Element 120 will have electron configuration
[Rn] 7s26d107p6 8s2 ... it has two valence
shell electrons in the 8s orbital which puts
it in group two. The 'standard' ion charge
for a group two element is 2+ so the formula
for the oxide of element 120 should be ElO.
Example Problems
6. Element 120 (El) has just been discovered.
What will the most likely formula of its oxide
be? Will it be larger or smaller than the
previous element in the table?
The previous element would be element 119
which would have to be in group 1. The trend
in atomic radii is for them to decrease from
left to right across the table so that element
120 should be smaller than element 119.