Download General and Inorganic Chemistry I.

General and Inorganic Chemistry I.
Lecture 1
István Szalai
Eötvös University
István Szalai (Eötvös University)
Lecture 1
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Lecture 1
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Outline
1
Periodic Table
István Szalai (Eötvös University)
Periodic Table
Electron Configurations
The electrons occupy the orbitals in the way that gives the lowest
energy for the atom.
Pauli Exclusion Principle: No two electrons in an atom may have
identical sets of four quantum numbers.
Hund’s Rule: Electrons occupy all the orbitals of a given subshell
singly before pairing begins. These unpaired electrons have parallel
spins.
Electrons are assigned to orbitals in order of increasing value of
(n + l).
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Periodic Table
Electron Configurations
Li 1s 2 2s 1
C 1s 2 2s 2 2p 2
Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6
Cu 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10
Ce 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 4f 1
Valence shell: the electrons in the outer shell, those that were not present
in the precding noble gas orbitals.
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Periodic Table
Periodic Table
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Periodic Table
The properties of the elements are periodic functions of
their atomic number
2 Li + 2 H2 O → 2 Li+ + 2 OH− + H2
2 Na + 2 H2 O → 2 Na+ + 2 OH− + H2
2 K + 2 H2 O → 2 K+ + 2 OH− + H2
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Periodic Table
History of the periodic table
About 330 B.C Aristotle proposed that everything is made up of a
mixture of one or more of four ”roots”. The four elements were
earth, water, air and fire.
In 1661 Boyle defined an element as a substance that cannot be
broken down into a simpler substance by a chemical reaction.
Lavoisier published a list of elements in 1789, or substances that
could not be broken down further, which included oxygen, nitrogen,
hydrogen, phosphorus, mercury, zinc, and sulfur. Lavoisier’s
descriptions only classified elements as metals or non-metals.
In 1817, Johann Wolfgang Döbereiner began to formulate one of the
earliest attempts to classify the elements. He found that some
elements formed groups of three with related properties. He termed
these groups ”triads”. (Cl-Br-I, Ca-Sr-Ba, S-Se-Te. . . )
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Periodic Table
History of the periodic table
Alexandre-Emile Béguyer de Chancourtois, a French geologist, was
the first person to notice the periodicity of the elements — similar
elements seem to occur at regular intervals when they are ordered by
their atomic weights. He devised an early form of periodic table,
which he called the telluric helix. With the elements arranged in a
spiral on a cylinder by order of increasing atomic weight, de
Chancourtois saw that elements with similar properties lined up
vertically.
John Newlands was an English chemist who in 1865 classified the 56
elements that had been discovered at the time into 11 groups which
were based on similar physical properties.
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Periodic Table
Dmitri Mendeleev (1837-1907)
Dmitri Mendeleev, a Siberian-born Russian chemist, was the first scientist
to make a periodic table much like the one we use today. His table was
published in 1869. It stated:
The elements, if arranged according to their atomic weights, exhibit
an apparent periodicity of properties.
Elements which are similar as regards to their chemical properties
have atomic weights which are either of nearly the same value (e.g.,
Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs).
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Periodic Table
Dmitri Mendeleev (1837-1907)
The arrangement of the elements, or of groups of elements in the
order of their atomic weights, corresponds to their so-called valencies,
as well as, to some extent, to their distinctive chemical properties; as
is apparent among other series in that of Li, Be, Ba, C, N, O, and Sn.
We must expect the discovery of many yet unknown elements–for
example, elements analogous to aluminium and silicon–whose atomic
weight would be between 65 and 75.
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Periodic Table
Germanium
Properties
Atomic weight
Density
Oxide
Chloride
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Mendeleev prediction
72
5,5 g/cm3
EO2
ECl4
Lecture 1
Observed values
72,6
5,35 g/cm3
GeO2
GeCl4
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Periodic Table
Refinements to the periodic table
In 1914 Henry Moseley found a relationship between an element’s
X-ray wavelength and its atomic number (Z), and therefore
resequenced the table by nuclear charge rather than atomic weight.
Thus Moseley placed argon (Z=18) before potassium (Z=19) based
on their X-ray wavelengths, despite the fact that argon has a greater
atomic weight (39.9) than potassium (39.1). The new order agrees
with the chemical properties of these elements, since argon is a noble
gas and potassium an alkali metal.
In 1945, Glenn T. Seaborg proposed a significant change to
Mendeleev’s table: the actinide series. Seaborg’s actinide concept of
heavy element electronic structure, predicting that the actinides form
a transition series analogous to the rare earth series of lanthanide
elements.
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Periodic Table
Classification of Elements
Name
Representative Elements
Alkali Metals
Alkaline Earth Metals
Boron group
Carbon group
Nitrogen group
Oxygen group
Halogens
Noble Gases
Transition Metals
Lanthanides and Acthinides
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Electron structure
ns 1
ns 2
ns 2 np 1
ns 2 np 2
ns 2 np 3
ns 2 np 4
ns 2 np 1
ns 2 np 6
ns 2 (n − 1)d 1−10
ns 2 (n − 1)d 1 (n − 2)f 1−14
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Periodic Table
Periodic Properties
atomic radii, ionic radii
ionization energy
electron affinity
electronegativity
periodic chemical properties
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Periodic Table
Atomic radii
Covalent radius: the nominal radius of the atoms of an element when
covalently bound to other atoms, as deduced the separation between
the atomic nuclei in molecules. In principle, the distance between two
atoms that are bound to each other in a molecule (the length of that
covalent bond) should equal the sum of their covalent radii.
Ionic radius: the nominal radius of the ions of an element in a specific
ionization state, deduced from the spacing of atomic nuclei in
crystalline salts that include that ion. In principle, the spacing
between two adjacent oppositely charged ions (the length of the ionic
bond between them) should equal the sum of their ionic radii.
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Periodic Table
Atomic radii
Metallic radius: the nominal radius of atoms of an element when
joined to other atoms by metallic bonds.
van der Waals radius: in principle, half the minimum distance
between the nuclei of two atoms of the element that are not bound to
the same molecule.
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Periodic Table
Atomic radii
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Periodic Table
Atomic radii
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Periodic Table
Atomic radii
Effective nuclear charge:
Zeff = Z − S where S is the shielding effect of the inner electrons.
l
0
1
2
3
...
1,0
1,0
1,0
1,0
ni−1
0,85
0,85
1,0
1,0
ni
0,3
0,35
0,35
0,35
Na 1s 2 2s 2 2p 6 3s 1 Zeff = 11 − (2 × 1 + 8 × 0, 85) = 2, 2
Al 1s 2 2s 2 2p 6 3s 2 3p 1 Zeff = 13 − (2 × 1 + 8 × 0, 85 + 2 × 0, 3) = 3, 6
Na 1.86 × 10−10 m (186 pm)
Al 1.43 × 10−10 m (143 pm)
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Periodic Table
Ionic radii
Isolectronic species (10e − )
Na+
Mg2+
radius (nm) 0,097 0,066
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Al3+
0,051
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Periodic Table
Ionization Energy
First ionization energy: the minimum amount of energy required to
remove the most loosely bound electron from an isolated gaseous atom.
Be(g) → Be+ (g) + e −
899 kJ/mol
+
2+
−
Be (g) → Be (g) + e
1757 kJ/mol
IE1 < IE2
2+
3+
−
Be (g) → Be (g) + e
14849 kJ/mol IE1 < IE2 IE3
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Periodic Table
Ionization Energy
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Periodic Table
Ionization Energy
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Periodic Table
Electron Affinity
The amount of energy absorbed when an electron is added to an isolated
gaseous atom.
Cl(g) + e − → Cl− (g)
−349 kJ/mol
−
−
O(g) + e → O (g)
−141 kJ/mol
−
−
2−
O (g) + e → O (g) +744 kJ/mol
Generally, nonmetals have more positive EA than metals. Atoms whose
anions are more stable than neutral atoms have a greater EA.
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Periodic Table
Electron Affinity
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Periodic Table
Electron Affinity
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Periodic Table
Electronegativity
The electronegativity of an element is a measure of the relative tendency
of an atom to attract electrons to itself when it is chemically combined
with another atom.
Pauling proposed the concept of electronegativity in 1932 as an
explanation of the fact that the covalent bond between two different
atoms (A–B) is stronger than would be expected by taking the average of
the strengths of the A–A and B–B bonds.
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Periodic Table
Electronegativity
To calculate Pauling electronegativity for an element, it is necessary to
have data on the dissociation energies of at least two types of covalent
bond formed by that element.
p
∆ = EAB − (EAA × √
EBB )
ENA − ENB = 0.102 ∆
ENF = 4.0
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Periodic Table
Electronegativity
Mulliken proposed that the arithmetic mean of the first ionization energy
and the electron affinity should be a measure of the tendency of an atom
to attract electrons.
EN =
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IE + EA
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Periodic Table
Electronegativity
Allred and Rochow considered that electronegativity should be related to
the charge experienced by an electron on the ”surface” of an atom: the
higher the charge per unit area of atomic surface, the greater the tendency
of that atom to attract electrons.
EN = 0, 359
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Zeff
+ 0, 744
r2
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Periodic Table
Electronegativity
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Periodic Table
Chemical properties: hydrides
Ionic hydride → covalent hydrides
LiH
NaH
KH
BeH2
MgH2
CaH2
B2 H6
(AlH3 )x
Ga2 H6
CH4
SiH4
GeH4
NH3
PH3
AsH3
H2 O
H2 S
H2 Se
HF
HCl
HBr
LiH + H2 O → Li+ + OH− + H2
HCl + H2 O → H3 O+ + Cl−
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Periodic Table
Chemical properties: oxides
Metal oxides → Nonmetal oxides
Li2 O
Na2 O2
KO2
BeO
MgO
CaO
B2 O2
Al2 O3
Ga2 O3
CO2
SiO2
GeO2
N2 O5
P4 O10
As2 O5
SO3
SeO3
OF2
Cl2 O7
Br2 O7
CaO + H2 O → Ca2+ + 2 OH−
SO3 + 3 H2 O → 2 H3 O+ + SO2−
4
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Periodic Table
Ionic Bonding
2 Na(s) + Cl2 (g) → 2 NaCl(s)
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Periodic Table
Ionic Bonding
2 Na(s) + Cl2 (g) → 2 NaCl(s)
Na
[Ne]
Cl
[Ne]
↑
3s
↑↓
3s
↑↓ ↑↓ ↑
3p
→
Na+
[Ne]
→
Cl−
[Ne]
↑↓
3s
↑↓ ↑↓ ↑↓
3p
Q2 α
EB =
4π0 r
Electrostatic interaction
(Coulomb force)
α(NaCl) = 1.7475
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Periodic Table
Ionic Bonding
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Periodic Table
Types of Ions
Noble gas configuration
s 2 : H− , Li+ , Be2+
s 2 p 6 : pl. Na+ , Ca2+ , Sc3+ , Cl− , O2− . . .
d 10 s 2 configuration
Sn ([Kr]5s 2 4d 10 5p 2 ) → Sn2+ ([Kr]5s 2 4d 10 ) + 2e−
Tl+ ,Pb2+ , Bi3+ , . . .
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Periodic Table
Types of Ions
Ions of transition metals
The first transition series is the result of the 3d orbitals being filled after
the 4s orbital. However, once the electrons are established in their orbitals,
the energy order changes - and in all the chemistry of the transition
elements, the 4s orbital behaves as the outermost, highest energy orbital.
The reversed order of the 3d and 4s orbitals only applies to building the
atom up in the first place. In all other respects, the 4s electrons are always
the electrons you need to think about first.
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Periodic Table
Types of Ions
Ions of transition metals
d 10 configuration
Zn ([Ar]4s 2 3d 10 ) → Zn2+ ([Ar]3d 10 ) + 2e−
Cu+ , Ag+ , Cd2+ , Tl3+ , . . .
[Ar]3d 1 :
[Ar]3d 2 :
[Ar]3d 3 :
[Ar]3d 4 :
[Ar]3d 5 :
[Ar]3d 6 :
[Ar]3d 7 :
[Ar]3d 8 :
[Ar]3d 9 :
Ti3+
V3+
Cr3+
Mn3+
Mn2+ , Fe3+
Fe2+ , Co3+
Co2+
Ni2+
Cu2+
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Periodic Table
Types of Ions
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