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CHEMISTRY 1001
Part 9
Electrons and Photons
In Chemistry-Electronic
Chemical Properties (the
Periodic Table)-PT
DR Kresimir Rupnik
LSU
2005, Intelliome LLC , Prentice Hall
Periodic Trends in the
Properties of the Elements
so,
the periodic table is actually the
table of electronic structure of
atoms of different elements !!!
Chemistry depends on the
electronic structure!!!!
Electron Configurations and the Periodic
Table
The periodic table can be used as a guide for electron
configurations.
The period number is the value of n.
Groups 1A and 2A have the s-orbital filled.
Groups 3A - 8A have the p-orbital filled.
Groups 3B - 2B have the d-orbital filled.
The lanthanides and actinides have the f-orbital filled.
Note that the 3d orbital fills after the 4s orbital.
Similarly, the 4f orbital fills after the 5d orbital.
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Electron Configurations and the Periodic
Table
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Periodic Trends in atomic Properties
One of the most obvious atomic propertyelectron structure dependence: size of the
atom
Start from the Electron Shells in Atoms!!
As the principal quantum number increases, the size
of the orbital increases.We can plot the 2
(probability of finding an electron) versus distance,
r, is called a radial plot of electron density (see
electronic structure Ch 3.).
Atomic size varies consistently through the periodic
table.
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Electron Shells and the Sizes of Atoms
Electron Shells in Atoms
The ns orbitals all have the same shape, but have
different sizes and different numbers of nodes.
Consider: He: 1s2, Ne: 1s2 2s22p6, and Ar: 1s2 2s22p6
3s23p6.
The radial electron density is the probability of finding
an electron at a given distance.
For He there is only one maximum (for the two 1s electrons).
For Ne there are two maxima: one largely for the 1s electrons
(close to the nucleus) and one largely for the n = 2 electrons
(further from the nucleus).
For Ar there are three maxima: one each largely for n = 1, 2,
and 3.
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Electron Shells and the Sizes of Atoms
Electron Shells in Atoms
Consider a simple diatomic
molecule.
The distance between the two
nuclei is called the bond
distance.
If the two atoms which make
up the molecule are the same,
then half the bond distance is
called the covalent radius of the
atom.
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Electron Shells and the Sizes of Atoms
Atomic Sizes
As the principle quantum number increases (i.e., we
move down a group), the distance of the outermost
electron from the nucleus becomes larger. Hence, the
atomic radius increases.
As we move across the periodic table, the number of
core electrons remains constant. However, the nuclear
charge increases. Therefore, there is an increased
attraction between the nucleus and the outermost
electrons. This attraction causes the atomic radius to
decrease.
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Electron Shells and the Sizes of Atoms
Atomic Sizes –Atomic radii (Table 4.5 textbook)
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Trends in Atomic Size
• either volume or radius
treat atom as a hard marble
• Increases down a group
valence shell farther from nucleus
effective nuclear charge fairly close
• Decreases across a period (left to right)
adding electrons to same valence shell
effective nuclear charge increases
valence shell held closer
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Trends in Atomic Size
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Group IIA
2e2e-
Be (4p+ & 4e-)
4 p+
2e-
Mg (12p+ & 12e-)
8e2e12 p+
2e8e-
Ca (20p+ & 20e-)
8e2e-
+
16
p
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Period 2
1e2e3 p+
Li (3p+ & 3e-)
4e2e-
2e2e4 p+
Be (4p+ & 4e-)
6e2e-
6 p+
8 p+
C (6p+ & 6e-)
O (8p+ & 8e-)
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3e2e5 p+
B (5p+ & 5e-)
8e2e10 p+
Ne (10p+ & 10e-)
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Covalent Radius, elements 1 - 58
250
150
100
50
Atomic Number
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57
53
49
45
41
37
33
29
25
21
17
13
9
5
0
1
Radius, pm
200
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Example 9.6 – Choose the
Larger Atom in Each Pair
•
•
•
•
C or O
Li or K
C or Al
Se or I
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Example 9.6 – Choose the
Larger Atom in Each Pair
•
•
•
•
C or O
Li or K
C or Al
Se or I?
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Property: Ionization Energy of Atoms
The first ionization energy, I1, is the amount of energy
required to remove an electron from a free atom:
Na(g)  Na+(g) + e-.
The second ionization energy, I2, is the energy required
to remove an electron from a gaseous ion:
Na+(g)  Na2+(g) + e-.
The larger ionization energy, the more difficult it is to
remove the electron.
There is a sharp increase in ionization energy when a
core electron is removed.
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Ionization Energy
• minimum energy needed to remove an electron
from an atom
gas state
endothermic process
valence electron easiest to remove
M(g) + 1st IE  M1+(g) + 1 eM+1(g) + 2nd IE  M2+(g) + 1 efirst ionization energy = energy to remove electron from
neutral atom; 2nd IE = energy to remove from +1 ion; etc.
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Ionization Energy
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Ionization Energy of Elements 1-56
2500
1500
1000
500
Elements by Atomic Number
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Xe
Sb
d
C
h
R
o
M
Y
Kr
As
Zn
n
M
Ti
K
S
Al
e
N
N
Be
0
H
Ionization Energy, kJ/mol
2000
20
Ionization Energy of Group IA
1400
H
Ionization Energy, kJ/mol
1200
1000
800
600
Li
Na
K
400
Rb
Cs
200
0
H
Li
Na
K
Elements by Period Number
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Rb
Cs
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Covalent Radii of Group IA
250
6 Cesium
5 Rubidium
Radii, pm
200
4 Potassium
3 Sodium
150
2 Lithium
100
50
1 Hydrogen
0
Hydrogen
Lithium
Sodium
Potassium
Rubidium
Cesium
1
2
3
4
5
6
Group Number
Ionization Energy, Group IA
1400
H
1200
Energy, kJ/mol
1000
800
600
Li
Na
K
400
Rb
Cs
200
0
H
Li
Na
K
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Group Number
Rb
Cs
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Ionization Energy of Periods 2 & 3
Ionization Energy, kJ/mol
2500
Ne
2000
F
Ar
1500
N
O
Cl
C
P
1000
S
Be
B
Al
Li
500
Si
Mg
Na
0
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Elements by Group Number
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Ionization Energy
Periodic Trends in Ionization Energy
Ionization energy decreases down a group.
This means that the outermost electron is more readily
removed as we go down a group.
As the atom gets bigger, it becomes easier to remove an
electron from the most spatially extended orbital.
Ionization energy generally increases across a period.
As we move across a period, Zeff increases. Therefore, it
becomes more difficult to remove an electron.
Two exceptions: removing the first p electron and
removing the fourth p electron.
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Ionization Energy
Periodic Trends in Ionization Energy
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Trends in Ionization Energy
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Trends in Ionization Energy
• as atomic radius increases, the IE generally
decreases
because the electron is closer to the nucleus
• 1st IE < 2nd IE < 3rd IE …
• 1st IE decreases down the group
valence electron farther from nucleus
• 1st IE generally increases across the period
effective nuclear charge increases
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Example 9.7 – Choose the Atom with the
Highest Ionization Energy in Each Pair
•
•
•
•
Mg or P
As or Sb
N or Si
O or Cl
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Example 9.7 – Choose the Atom with the
Highest Ionization Energy in Each Pair
•
•
•
•
Mg or P
As or Sb
N or Si
O or Cl?
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Metallic Character
• how well an element’s properties match the
general properties of a metal
• 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
• Nonmetals
 brittle in solid state
 dull
 electrical and thermal insulators
 most oxides are acidic and molecular
 form anions and polyatomic anions
 gain electrons Prentice
in reactions
- reduced
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Trends in Metallic Character
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Example 9.8 – Choose the
More Metallic Element in Each Pair
•
•
•
•
Sn or Te
Si or Sn
Br or Te
Se or I
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Example 9.8 – Choose the
More Metallic Element in Each Pair
•
•
•
•
Sn or Te
Si or Sn
Br or Te
Se or I?
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Metals, Nonmetals, and Metalloids
Metals
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Property: Electron Affinities
Electron affinity is the energy change when a gaseous
atom gains an electron to form a gaseous ion:
Cl(g) + e-  Cl-(g)
Look at electron configurations to determine whether
electron affinity is positive or negative.
The extra electron in Ar needs to be placed in the 4s orbital
which is significantly higher in energy than the 3p orbital.
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Electron Affinities
The added electron in Cl is placed in the 3p orbital to form
the stable 3p6 electron configuration.
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Properties of Main-Group Elements
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Metals, Nonmetals, and Metalloids
Metals
Metal character refers to the properties of metals (shiny
luster, malleable and ductile, oxides form basic ionic
solids, and tend to form cations in aqueous solution).
Metal character increases down a group.
Metal character decreases across a period.
Metals have low ionization energies.
Most neutral metals are oxidized rather than reduced.
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Active Metals
Group 1A: The Alkali Metals
Alkali metals are all soft.
Chemistry dominated by the loss of their single s
electron:
M  M+ + eReactivity increases as we move down the group.
Alkali metals react with water to form MOH and
hydrogen gas:
2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
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Active Metals
Group 1A: The Alkali Metals
Alkali metal produce different oxides when reacting
with O2:
4Li(s) + O2(g)  2Li2O(s)
(oxide)
2Na(s) + O2(g)  Na2O2(s)
(peroxide)
K(s) + O2(g)  KO2(s)
(superoxide)
Alkali metals emit characteristic colors when placed in
a high temperature flame.
The s electron is excited by the flame and emits energy
when it returns to the ground state.
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Active Metals
Group 1A: The Alkali Metals
Li line: 2p  2s
transition
Na line (589 nm):
3p  3s transition
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K line: 4p  4s
transition
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Active Metals
Group 2A: The Alkaline Earth Metals
Alkaline earth metals are harder and more
dense than the alkali metals.
The chemistry is dominated by the loss of two s
electrons:
M  M2+ + 2e-.
Mg(s) + Cl2(g)  MgCl2(s)
2Mg(s) + O2(g)  2MgO(s)
Be does not react with water. Mg will only react
with steam. Ca onwards:
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
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Important Nonmetals
Hydrogen
Hydrogen is a unique element.
Most often occurs as a colorless diatomic gas, H2.
It can either gain another electron to form the hydride
ion, H-, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(g)
H+ is a proton.
The aqueous chemistry of hydrogen is dominated by
H+(aq).
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Important Nonmetals
Group 6A: The Oxygen Group
As we move down the group the metallic
character increases (O2 is a gas, Te is a metalloid,
Po is a metal).
There are two important forms of oxygen: O2 and
ozone, O3. Ozone can be prepared from oxygen:
3O2(g)  2O3(g) H = +284.6 kJ.
Ozone is pungent and toxic.
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Important Nonmetals
Group 6A: The Oxygen Group
Oxygen (or dioxygen, O2) is a potent oxidizing
agent since the O2- ion has a noble gas
configuration.
There are two oxidation states for oxygen: 2- (e.g.
H2O) and 1- (e.g. H2O2).
Sulfur is another important member of this
group.
Most common form of sulfur is yellow S8.
Sulfur tends to form S2- in compounds (sulfides).
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Important Nonmetals
Group 7A: The Halogens
The chemistry of the halogens is
dominated by gaining an electron to form
an anion:
X2 + 2e-  2X-.
Fluorine is one of the most reactive
substances known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)
H = -758.7 kJ.
All halogens consists of diatomic
molecules, X2.
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Important Nonmetals
Group 7A: The Halogens
Chlorine is the most industrially useful halogen. It is
produced by the electrolysis of brine (NaCl):
2NaCl(aq) + 2H2O(l)  2NaOH(aq) + H2(g) + Cl2(g).
The reaction between chorine and water produces
hypochlorous acid (HOCl) which disinfects pool water:
Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq).
Hydrogen compounds of the halogens are all strong
acids with the exception of HF.
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Group 8A: The Noble Gases
These are all nonmetals and monatomic.
They are notoriously unreactive because they have
completely filled s and p sub-shells.
In 1962 the first compound of the noble gases was
prepared: XeF2, XeF4, and XeF6.
To date the only other noble gas compound known is
KrF2.
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