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Lecture 1
• Contents:
• Introduction
• What is Inorganic Chemistry ?
• Main Group Chemistry
• Reading the Periodic Table: Classification
• Diagonal Relationship
• Electronegativity
• Electron affinity- ionization energy
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Introduction
What is Inorganic Chemistry ?
• “The chemistry of everything that is NOT organic…”
• “The chemistry of all of the elements and their compounds
except for the hydrocarbons and their derivatives.”
• “The branch of chemistry falling between and overlapping
with physical chemistry and organic chemistry.”
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Why Should You Study Inorganic Chemistry ?
• Essentially the entire universe is Inorganic.
Elemental Composition of the Sun and the Universe
Sun
Universe
Hydrogen
Helium
All Others
92.5 %
7.3 %
0.2 %
90.87 %
9.08 %
0.05 %
• The Earth is predominantly Inorganic.
Elemental Composition of the Earth’s Crust
Oxygen
Silicon
Aluminum
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45.5 %
27.2 %
8.30 %
Iron
Calcium
All Others
6.20 %
4.66 %
8.14 %
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 Main Group Elements: s and p block elements
 Transition Elements: d elements
 Lanthanides and Actinides: f block elements
p
s
d
In which d orbit is being filled
In which p orbit is being filled
In which s orbit is being filled
f
In which f orbit is being filled
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Main Group Chemistry
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Lanthanides
Actinides
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Reading the Periodic Table: Classification
• Nonmetals, Metals, Metalloids, Noble gases
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In the periodic table of elements, the
number of the orbital electrons, that is the
atomic number, and their arrangement
determines the chemical and physical
properties of an element. Each group of
elements has a characteristic electronic
arrangement and therefore elements within
one group exhibit similar physical and
chemical properties
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s-block elements
1-The alkali metals: in which all elements have one s
electron in their outer shell, appear in a vertical column named
Group 1. ns1
2-the alkaline earth metals: which have two s electrons in
their outer shell, form Group 2. ns2
p-block elements
Groups 13
ns2np1
Groups 14
ns2np2
Groups 15
ns2np3
Groups 16
ns2np4
Groups 17
ns2np5
Groups 18 or 0
ns2np6
(except He:1s2)
The inert gases which appear in a group labeled Group 0
have the most stable arrangement of electrons because their
outer shell of electrons is full. This explains their lack of
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chemical reactivity.
 Core Designation - A designation of electronic
configuration wherein the outer shell electrons
are shown along with the “core” configuration of
the closest previous noble gas.
Li
Be
[He] 2s2
Na [Ne] 3s1
Mg
[Ne] 3s2
[Ar] 4s1
Ca
[Ar] 4s2
Rb [Kr] 5s1
Sr
[Kr] 5s2
K
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[He] 2s1
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Electronic Configuration and the Periodic Table
 s-Block Elements
 p-Block Elements
 d-Block Elements
 f-Block Elements
Electronic Configuration for positive ions (cations) Cations are formed by removing electrons in order
of decreasing n value. Electrons with the same n
value are removed in order of decreasing l value.
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 Atomic Size Atomic radii are considered to be 1/2 of the
average distance between centers of identical
atoms that are touching each other. This will
vary with the chemical environment the atom
is in.
1.42 Å
1.54 Å
Fluorine
Diamond
C – 0.77 Å
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F – 0.71 Å
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Trends in Atomic Radii:
1. Atomic radii increase from top to bottom in a
family or group.
The number of electrons and the nuclear
charge are increasing! - Tends to shrink atom.
But extra electron are added to new shells that
are further from the nucleus and more
effectively shielded from the nucleus - Tends to
make the atom larger.
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2. Atomic radii decrease from left to right across a row or
period.
The number of electrons and the nuclear charge are
increasing! - Tends to shrink atom.
The electrons are being added to the same shell and are
not well shielded and thus, the atoms get smaller.
3. Summary of trends
Down a Group - Larger
Across a Period - Smaller
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What Affects Atomic/Ionic Sizes?
 The Charge on the Nucleus
 Shielding - This reduces the actual nuclear charge
resulting in an “effective” nuclear charge.
In general, Zeff (effective nuclear charge) increases
across a period but remains about the same or
slightly decrease down a group.
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 Ionic Size Based on the internuclear distance of cations and anions
in ionic crystals.
Cations - Monatomic cations are smaller than
their parent atoms.
The whole outer shell is typically removed.
The effective nuclear charge is increased.
Na atom
1.86 Å
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Na+ ion
1.02 Å
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Anions - Monatomic anions are larger than
their parent atoms.
The extra electrons are typically added to the same
shell where they are repelled by the other electrons
already present, making the ion bigger than its parent
atom.
F Atom
0.71 Å
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Fluoride Ion
1.36 Å
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 Ionization Energy - The energy required to remove an
electron from a gaseous ground-state atom or ion.
A. First Ionization Energy - The energy required to
remove the most loosely bound electron from the valence
shell.
B. Second Ionization Energy - The energy required to
remove the second electron after the first one is gone.
C. Third Ionization Energy - Etc., Etc., Etc.
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Li (g)
Li+
+
e-
IE1 = +520 kJ/mol
Li+
Li2+ +
e-
IE2 = +7298 kJ/mol
Na (g)
Na+
+
e-
IE1 = +496 kJ/mol
Na+
Na2+ +
e-
IE2 = +4564 kJ/mol
Mg (g)
Mg+
+
e-
IE1 = +737 kJ/mol
Mg+
Mg2+ +
e-
IE2 = +1447 kJ/mol
Mg2+
Mg3+ +
e-
IE3 = +7738 kJ/mol
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Trends in Electron Affinities Increases up a group.
Increases from left to right in a period.
H
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Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
Ga
Ge
As
Se
Br
Rb
Sr
In
Sn
Sb
Te
I
Cs
Ba
Tl
Pb
Bi
Po
At
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Summary of Trend
• Periodic Table and Periodic Trends
• 1. Electron Configuration
3. Ionization Energy: Largest toward NE of PT
4. Electron Affinity: Most favorable NE of PT
2. Atomic Radius: Largest toward SW corner of PT
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Metals Blue (underlined)
Non-metals green (italics)
Metalloids red (Bold)
1
2
3
4
5
6
7
Li
Be
B
C
N
O
F
Na
Mg
Al
Si
P
S
Cl
K
Ca
Ga
Ge
As
Se
Br
Rb
Sr
In
Sn
Sb
Te
I
Cs
Ba
Tl
Pb
Bi
Po
At
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p-block and the frontier zone
• So this gives rise to frontier between metals and non-
•
metals and it is in this zone are the metalloids that we
have encountered earlier. The metalloid have some of
the characteristics of metals and non-metals and this
frontier is denoted by this zigzag line
3
4
5
6
7
B
C
N
O
F non
Al
Si
P
S
Cl metals
Ga
Ge
As
Se
Br
In
Sn
Sb
Te
I
Tl
Pb
Bi
Po
As
Metals
Metalloids
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Electrongetaivities for selected
elements (Pauling Scale)
H
2.1
Li
1.0
Na
0.9
K
0.8
Be
1.5
Mg
1.2
Ca
1.0
B
2.0
Al
1.5
C
2.5
Si
1.8
N
3.0
P
2.1
Ga
Ge
As
O
3.5
S
2.5
Se
2.5
F
4.0
Cl
3.0
Br
2.8
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Diagonal
Relationships
The elements in each
encircled pair have
several similar
properties.
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Diagonal Relationship
• Li
Na
Be
Mg
B
Al
C
Si
• Elements that are linked by the arrows in the
diagram above are said to be diagonally linked.
These pairs of elements often show similar
chemical properties e.g. Li and Mg both form
nitrides.
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Group IA: The Alkali Metals
The metals in Group IA (Li, Na, K, Rb, Cs, and Fr) are called the
alkali metals because they all form hydroxides (such as NaOH)
that were once known as alkalies.
The electron configurations of the alkali metals are characterized
by a single valence electron. As a result, the chemistry of these
elements is dominated by their tendency to lose an electron to
form positively charged ions (Li+, Na+, K+).
Li:
[He] 2s1
Rb:
[Kr] 5s1
Na:
[Ne] 3s1
Cs:
[Xe] 6s1
K:
[Ar] 4s1
Fr:
[Rn] 7s1
The alkali metals lose electrons so easily that sodium dissolves
in liquid ammonia at temperatures below the boiling point of
ammonia (-33oC) to give Na+ ions and electrons.
Na(s)
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NH3(l)
Na+ + e28
Group IIA: The Alkaline-Earth Metals
The elements in Group IIA (Be, Mg, Ca, Sr, Ba and Ra) are all metals,
and all but Be and Mg are active metals.
The term alkaline reflects the fact that many compounds of these
metals are basic or alkaline. The term earth was historically used to
describe the fact that many of these compounds are insoluble in
water.
Three points should be kept in mind, however.
1-The alkaline-earth metals tend to lose two electrons to form M 2+
ions (Be2+, Mg2+, Ca2+, and so on).
2-These metals are less reactive than the neighboring alkali metal.
Magnesium is less active than sodium; calcium is less active than
potassium; and so on.
3-These metals become more active as we go down the column.
Magnesium is more active than beryllium; calcium is more active than
magnesium; and so on.
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Inert pair effect
Al and Tl are both metals in group 3 of the
periodic table, but Al ions are only ever found in
the +3 state. (Al3+ cations), but Tl is known to
form compounds in which there can be Tl+ or
Tl3+ cations. This tendency for elements at the
bottom of groups 3, 4 and 5 to form compounds
in which their outermost s electrons are not
involved in bonding is called the inert pair
effect. The basic reason is that the s electrons
see much more of the nucleus than the p
electrons so they are more stable (the s
electrons are more penetrating)
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Properties and Trends
in Group 3A
• Boron tends to form covalent compounds
rather than ionic compounds.
• The rest of group 3A elements can form 3+
ions.
• Gallium, indium, and thallium also often form
1+ ions by retaining their ns2 electrons; this is
called the inert pair effect.
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