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Chapter 1

CHMB31 Chapter 1


Department
Chemistry
Course Code
CHMB31H3
Professor
Alen Hadzovic
Chapter
1

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Chapter 1: Atomic Structure
The Origin of Elements
Strong Force: a short-range but powerful attractive force between protons and
neutrons bound these particles together into nuclei.
Electromagnetic Force: a relatively weak but long-range force between electric
charges bound electrons to nuclei to form atom.
Atomic Number (Z): the number of protons in the nucleus of an atom of the element.
Isotopes: atoms with the same atomic number but different atomic masses.
Mass Number (A): the total number of protons and neutrons in the nucleus.
Nuclear reactions are very much more energetic than normal chemical reactions
because the strong force is much stronger than the electromagnetic force that binds
electrons to nuclei.
Nuclide: a nucleus of specific atomic number Z and mass number A is designated AZE.
β-Particle: denoted as 0-1e.
Positron: denoted as 01e, it is a positively charged electron with a zero mass number.
Elements up to Z = 26 were formed inside stars, the products of nuclear fusion
reactions.
Binding Energy: represents the difference in energy between the nucleus itself and
the same numbers of individual protons and neutrons (Ebind = delta-m x c2).
oA positive binding energy corresponds to a nucleus that has a lower, more
favorable energy than its constituent nucleons.
oEven-Z nuclides are slightly more stable than their odd-Z neighbors.
oA large binding energy signifies a stable nucleus.
Nuclei close to iron (Fe) are the most stable and heavier elements are produced by a
variety of processes that require energy.
oNeutron Capture: 9842Mo + 10n 9942Mo + γ
oFollwed by β decay accompanied by neutrino emission: 9842Mo 9943Tc + e- +
v
Daughter Nuclide: the product of a nuclear reaction
The Structures of Hydrogenic Atoms
Hydrogenic Atoms or Hydrogen-Like Atoms: have only one electron and are free of
the complicating effects of electron-electron repulsions (He+ and C5+).
Many-Electron Atoms: atoms with more than one electron.
All wavelengths can be described by the expression: 1/λ = R(1/n12 – 1/n22), where R is
the Rydberg constant with a value of 1.097 x 107 m-1.
Wave-Particle Duality: electrons can behave as particles or as waves.
Uncertainty Principle: it is impossible to know the linear momentum and the location
of an electron simultaneously so the product of the uncertainty in momentum and
the uncertainty in position cannot be less than a quantity of the order of Plank’s
constant ½ħ, where ħ = h/2π.
Wavefunction or psi): a mathematical function of the position coordinates x, y and
z which describes the behavior of an electron.

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Schrodinger Equation: kinetic energy contribution + potential energy contribution =
total energy contribution
Quantization of Energy: an electron can possess only certain discrete energies in an
atom.
Probability Density: probability of finding an electron at a given location is
proportional to the square of the wavefunction at that point, Ψ2.
Constructive Interference: a positive region of one wavefunction may add to a
positive region of the other wavefunction to give a region of enhanced amplitude.
Destructive Interference: a positive region of one wavefunction may be cancelled by
a negative region of the second wavefunction (reduces the probability that an
electron will be found in that region).
Atomic Orbital: the wavefunction of an electron in an atom.
oPrincipal Quantum Number (n): energy of the bound electron; indicates the
size of the orbital.
oOrbital Angular Momentum Quantum Number (l): magnitude of the orbital
angular momentum and indicates the angular shape of the orbital (for a given
value of n, l can have values of 0, 1,…, n-1) the value of l is the subshell
designation (0 = s, 1 = p, 2 = d and 3 = f).
oMagnetic Quantum Number (ml): orientation of that angular momentum (for a
given value of l, ml can have integer values from +l to –l).
oDegenerate: all orbitals with the same value of n have the same energy.
oS Orbital: only one orbital with l = 0 and ml = 0.
oP Orbitals: three orbitals with l = 1 and ml = +1, 0 and -1.
oD Orbitals: five orbitals with l = 2 and ml = +2, +1, 0, -1 and -2.
oSpin: described by two quantum numbers, s and ms.
oMs can take only two values, for anticlockwise spin (spin-up) and –½ for
clockwise spin (spin-down).
oNodes: the positions where either component of the wavefunction passes
through zero.
Radial Nodes: occur where radial component of the wavefunction
passes through zero.
Angular Nodes: occur where the angular component of the
wavefunction passes through zero.
The numbers of both types of node increase with increasing energy
and are related to the quantum numbers n and l.
oA 1s orbital, the wavefunction with n = 1, l = 0 and ml = 0, decays
exponentially with distance from the nucleus and never passes through zero.
oA 2s orbital has n = 2, l = 0 and ml = 0, having one radial node.
oA 2p orbital has n = 2, l = 1 has no radial nodes because its radial
wavefunction does not pass through zero anywhere but it is zero at the
nucleus (like all orbitals other than s orbitals).
oAs the principal quantum number (n) of an electron increases, it is found
further away from the nucleus and its energy increases.
oAn orbital with quantum numbers n and l in general has n – l – 1 radial nodes.
oRadial Distribution Function (P(r)) = r2R(r)2, where r is the radius.
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