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CHEM 511 chapter 1 page 1 of 12
Chapter 1
Basic Concepts: Atoms
What is inorganic chemistry?
The periodic table is made of elements, which are made of ...?
Particle Symbol Mass in amu Charge
1.0073 +1e
1.0087 0
5.486 10-4
-1e
Define: atomic number (Z):
Define: mass number (A):
From general chemistry we recall that atoms of the same element all contain the same number of
protons, but may or may not contain the same number of neutrons.
Isotopes:
Isotopes of hydrogen:
Isobars:
Isotones:
How can you roughly tell, by looking at the periodic table, if an element has just one isotope (i.e.,
the element is a monotope)?
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Distinction of isotope with allotrope.
Examples of allotropes with oxygen? carbon?
sulfur (below)
Inside the atom Early quantum mechanics (~1900-~1925) treated the electron as a particle. Why?
http://web.physik.rwth-aachen.de/~harm/aixphysik/atom/discharge/index1.html
Note that this equation works well for hydrogen and hydrogen-like atoms—ONLY!! Limitations?
Later quantum mechanics (~1925 and on)
Think of the electron NOT as a particle, but as a wave. What else has this “wave-particle duality”?
What is the wavelength of a 68 kg person walking at 4.0 mph (1.8 m/s)?
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Heisenberg’s Uncertainty Principle
HUP: It is impossible to know precisely the momentum and location of an electron simultaneously.
Schrodinger’s Wave Equation (Ψ)
Ψ represents the “wave function” (as opposed to a particle function) of the electron—it is based on
x, y, z coordinates and time (for convenience, it may be better to use polar coordinates (r, θ, Φ))
Ψ describes the behavior of an electron in a specific region of space (i.e., an orbital); it is a
mathematical function and can have positive or negative values—a negative value doesn’t mean
the electron isn’t present!
Atomic Orbitals
To define any orbital we need three quantum numbers: n, l ,ml 1. Principal Quantum Number (n)
Numerical values for n?
2. Orbital Quantum Number (l) aka orbital angular momentum Q. N. aka azimuthal Q. N.
Numerical values for l?
3. Magnetic Quantum Number (ml)
Numerical values for ml?
To plot Ψ versus distance from the nucleus results in graphs like this:
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Another useful feature is to look at the squared function of the wavefunction (Ψ2) (note, it isn’t
actually squared, Ψ is multiplied by a factor such that when integrated over all space, you are
guaranteed to find the electron somewhere in that space). Your book actually applies an additional
factor 4πr2 (surface area of a sphere), as shown below.
What can we infer about the shape (and size) of the s orbitals from these graphs?
Radial distribution plot for 2p & 3p Probability plots for other orbitals
4π
r2R
(r)2
4π
r2R
(r)2
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A note on radial nodes: to determine the number of radial nodes use the formula n-l-1.
(This does not include nodal planes.)
# of nodes for 1s:
# of nodes for 2s:
# of nodes for 3s:
# of nodes for 2p:
# of nodes for 3p:
p-orbitals
d-orbitals
The Orbitron shows informative pictures:
http://winter.group.shef.ac.uk/orbitron/
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One more quantum number: the magnetic spin quantum number (ms) aka spin quantum number)
Moving charges generate magnetic fields—thus two electrons close to each other must have
opposite spins
Pauli Exclusion Principle:
Result of the PEP?
Many Electron Atoms In a hydrogen atom (1 electron), all orbitals of the same n have the same energy
Introduction of just one additional electron changes this and "splits" the energy of the orbitals.
Why?
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Effective nuclear charge (Zeff): That portion of the total nuclear charge that is experienced by a
given electron.
Zeff = Zactual – S(Shielding factor)
Slater’s Rules for determining effective nuclear charge (empirically derived):
Write the electron configuration in the following groupings: (1s), (2s, 2p), (3s, 3p), (3d),
(4s, 4p), (4d), (4f), (5s, 5p), etc.
Electrons higher than the electron of interest do not contribute to the shielding factor
Electrons in s- and p-orbitals:
o Each electron in (ns, np) contribute S = 0.35
o Each electron in the n-1 shell contribute S = 0.85
o Each electron in the n-2 or lower shells contribute S = 1.00
Electrons in d- or f-orbitals
o Each of the other electrons in the other d- or f-orbitals contributes S = 0.35
o Each of the electrons in a lower group contributes S = 1.00
Why useful? It helps to determine electron configuration!
EX. Determine the Zeff for the outermost electron in the following electron configurations:
(a) 1s22s
22p
6 (b) 1s
22s
22p
53s
1
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The Classification of the Elements1
What are the three main categories of elements according to the periodic table?
Important sections of the periodic table:
main group elements
transition metal elements
rare earth elements aka inner transition elements
Important families (aka groups)
Group 1 (or 1A)
Group 2 (or 2A)
Group 15 (or 5A)
Group 16 (or 6A)
Group 17 (or 7A)
Group 18 (or 8A)
Group 11 (or 1B)
1 Other representations of the elements can be found here: http://www.meta-synthesis.com/webbook/35_pt/pt_database.php
Dangerous potential time-sink!
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Brief review of electron configuration
The ground state (lowest energy) electron configuration is determined by the Aufbau principle
1.
2.
3.
Filling orbitals is straight forward until the d-block
elements
Determine the electron configuration of:
B
P
Ti
Cr
Ni
Cu
Se
Bi
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Core vs. valence electrons?
How are cations created?
How are anions created?
Determine the electron configuration of:
Na+
P3-
Mn4+
Xe-
Fe3+
Fe2+
Ag+
Periodic Trends
Atomic radius: a measurement of the unionized form of an atom.
metallic radius: Usual method for metals is to measure the distance between nuclei in the
solid and divide by 2
covalent radius: For nonmetals, measure the distance between nuclei of a binary molecule
and divide by 2
ionic radius: a measurement of an ion's size, usually derived from the distance between an
oxygen nucleus and a metal ion. (an approximation only)!!
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Size trends within the periodic table?
Down a group?
Across a period?
Note: transition metals in the 5th and 6th periods are nearly the same size!! The period 6 atoms
have 32 MORE electrons than the period 5 atoms. How can they be the same size?
Atomic Radii (pm)
Period 5 Zr Nb Mo Tc Ru Rh Pd Ag
160 147 140 135 134 134 137 144
Period 6 Hf Ta W Re Os Ir Pt Au
159 147 141 137 135 136 139 144
What do we know about the relative size of cations to parent atoms? anions to parent atoms?
Cations are smaller than the parent (neutral) atoms
Anions are larger than the parent (neutral) atoms
Ionization energy: Energy needed to remove an electron from an isolated atom in the gas phase
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Based on this figure, what are the periodic trends for Ei1?
Down a group?
Across a period?
What about the dips at Be-B and N-O?
Electron Affinity: The energy gained or released when an electron is added to the valence shell
of an isolated atom in the gas phase.
