Tentative material to be covered for Exam 2

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Tentative material to be covered for Exam 2
(Wednesday, October 27) Chapter
17 Many-Electron Atoms and Chemical
Bonding 17.1 Many-Electron Atoms and the
Periodic Table 17.2 Experimental Measures of
Orbital Energies 17.3 Sizes of Atoms and
Ions 17.4 Properties of the Chemical
Bond 17.5 Ionic and Covalent Bonds 17.6 Oxidatio
n States and Chemical Bonding Chapter
18 Molecular Orbitals, Spectroscopy, and Chemical
Bonding 18.1 Diatomic Molecules 18.2 Polyatomic
Molecules 18.3 The Conjugation of Bonds and
Resonance Structures 18.4 The Interaction of
Light with Molecules 18.5 Atmospheric Chemistry
and Air Pollution
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Quantum mechanics provides an intellectual
structure for describing all of the properties of
atoms and molecules. For atoms quantum mechanics
the concept of orbitals (wavefunctions) provides
a description of the energies, the sizes of atoms
and the basis for bonding of atoms and the
construction of the periodic table. The orbitals
for the H atom, which are known precisely, are
used as starting approximation for building up
the electron configuration of multielectron
atoms. Every electron in an atom is assigned
four quantum numbers (n, l, ml and ms) that
uniquely define its spatial distrbution and spin
state. Thus, we can envision every electrons in
terms of a characteristic energy, size, shape,
orientation and spin.
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Properties of electrons in atoms Quantum numbers
of electrons Electron configurations Core
electrons Valence electrons Energy required to
remove an electron Energy required to add an
electron Size of atoms
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Building up the ground state configuration of
atoms Every atom possesses the SAME set of
available orbitals Every electron of an atom
MUST be in one of these orbitals 1s, 2s, 2p,
3s, 3p, 3d, etc. The energy and size of these
orbitals depend on the atom (Z), but the shape
and orientation is space of any orbitals of the
same l are the same for all atoms. The energy
ranking of the orbitals for the representative
elements is generally 1s lt 2s lt 2p lt 3s lt
3p. From this point on the next lowest energy
orbital may be 4s or 3d, depending on the number
of electrons in the neutral atom.
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The properties of the atoms of the elements vary
periodically with the atomic weights of the
elements. All chemical and physical properties
of the elements depend on their atomic weights
and therefore vary periodically with atomic
weight. The ground state electron
configuration of the atoms of elements vary
periodically with the atomic number Z. All
chemical and physical properties of the elements
that depend on electron configurations vary
periodically with atomic number.
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Ground state electron configuration Z electrons
(Z atomic number of the atom) are placed
seriatim into the orbitals according to the
following guidelines. Aufbau principle
electrons go into lowest energy orbitals
first. Pauli principle No more than two
electrons in any one orbital. Filled orbitals
have spins paired. Hunds rule When there are
orbitals of equal energy in a subshell to fill,
the electrons first go into different orbitals
with parallel spins one at a time.
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Valence electrons, Lewis structures and
electronic configurations The valence electrons
are electrons in the s and p orbitals valence
electrons snpm Atom Configuration Comment 3L
i He2s Paramagnetic 4Be He 2s2 Closed
shell (diamagnetic) 5B He 2s22p1 Paramagnetic
6C He 2s22p2 Paramagnetic 7N He 2s22p3
Paramagnetic 8O He 2s22p4 Paramagnetic 9F
He 2s22p5 Paramagnetic 10Ne He
2s22p6 Closed shell (diamagnetic)
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Correlation of valence electron and Lewis
structures
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Building up the third row of the periodic
table From Na to Ar
Atom Configuration Comment 11Na Ne2s Paramag
netic 12Mg Ne 2s2 Closed shell
(diamagnetic) 13Al Ne 2s22p1 Paramagnetic 14Si
Ne 2s22p2 Paramagnetic 15P Ne 2s22p3
Paramagnetic 16S Ne 2s22p4
Paramagnetic 17Cl Ne 2s22p5
Paramagnetic 18Ar Ne 2s22p6 Closed shell
(diamagnetic)
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d orbitals From photoelectron spectroscopy, the
3d subshell for elements 21 through 29 (Sc
through Cu) lies well above the 3d subshell.
However, the energy of the 3d subshell is very
close in energy to the 4s subshell 3p ltlt 3d
4s 1s ltlt 2s lt 2p ltlt 3s lt 3p lt 4s 3d Thus is
some cases the specifics of orbital
configurations place 3d below 4s and in other
cases th 4s is below the 3d.
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The fourth row of the periodic table Atom Confi
guration 19K 18Ar2s 20Ca 18Ar2s2 ____
_____________________________ d orbitals fill
up 31Ga 18Ar 2s22p1 32Ge 18Ar
2s22p2 33As 18Ar 2s22p3 34Se 18Ar
2s22p4 35Cl 18Ar 2s22p5 36Kr 18Ar
2s22p6 What about 21M through 30M?
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The electron configurations of the transition
elements 21Sc 18Ar4s23d 22Ti 18Ar4s23d2 23V
18Ar4s23d3 24Cr 18Ar4s23d4 instead
18Ar4s13d5 25Mn 18Ar4s23d5 26Fe
18Ar4s23d6 27Co 18Ar4s23d7 28Ni
18Ar4s23d8 29Cu 18Ar4s23d9 instead
18Ar4s13d10 30Zn 18Ar4s23d10 The
surprises for 24Cr and 29Cu are due to ignored
electron-electron repulsions. For 24Cr the
stability of half shells trumps one filled
subshell and a partially filled subshell. For
29Cu the stability of a full d shell and half
filled 4s subshell trumps a partially filled 3d
subshell.
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Example IE drops dramatically from He to Li.
Why? He 1s2 versus Li 1s22s. 2s on average
further away from nucleus for same average charge
(after screening by 1s2).
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Compare Be He2s2 versus B He2s22p Compare
N He2s22px2py2pz versus O
He2s22px22py2pz
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Bond length the distance between the centers
(nucleus) of bonded atoms.
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Atomic radius the atomic radius of a neutral
atom generally decreases from left to right
across a period (larger Z) and increases down a
group (increase in n).
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The electron affinity (EA) of an atom is the
energy change which occurs when an atom gains an
electron. X(g) e- Xe- (g) Electron
affinities of the representative elements What
are the correlations across and down?
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Electronegativity a measure of the power of an
atom to attract electrons to itself in a bond.
Most electronegative atoms F gt O gt Cl gtN Br gt I