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Chapter 1

This

**preview**shows pages 1-2. to view the full**6 pages of the document.**Chapter 1: Atomic Structure

The Origin of Elements

â€¢Strong Force: a short-range but powerful attractive force between protons and

neutrons bound these particles together into nuclei.

â€¢Electromagnetic Force: a relatively weak but long-range force between electric

charges bound electrons to nuclei to form atom.

â€¢Atomic Number (Z): the number of protons in the nucleus of an atom of the element.

â€¢Isotopes: atoms with the same atomic number but different atomic masses.

â€¢Mass Number (A): the total number of protons and neutrons in the nucleus.

â€¢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.

â€¢Nuclide: a nucleus of specific atomic number Z and mass number A is designated AZE.

â€¢Î²-Particle: denoted as 0-1e.

â€¢Positron: denoted as 01e, it is a positively charged electron with a zero mass number.

â€¢Elements up to Z = 26 were formed inside stars, the products of nuclear fusion

reactions.

â€¢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.

â€¢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 ïƒ 9942Mo + Î³

oFollwed by Î² decay accompanied by neutrino emission: 9842Mo ïƒ 9943Tc + e- +

v

â€¢Daughter Nuclide: the product of a nuclear reaction

The Structures of Hydrogenic Atoms

â€¢Hydrogenic Atoms or Hydrogen-Like Atoms: have only one electron and are free of

the complicating effects of electron-electron repulsions (He+ and C5+).

â€¢Many-Electron Atoms: atoms with more than one electron.

â€¢All wavelengths can be described by the expression: 1/Î» = R(1/n12 â€“ 1/n22), where R is

the Rydberg constant with a value of 1.097 x 107 m-1.

â€¢Wave-Particle Duality: electrons can behave as particles or as waves.

â€¢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 Plankâ€™s

constant Â½Ä§, where Ä§ = h/2Ï€.

â€¢Wavefunction (Î¨ or psi): a mathematical function of the position coordinates x, y and

z which describes the behavior of an electron.

Only pages 1-2 are available for preview. Some parts have been intentionally blurred.

â€¢Schrodinger Equation: kinetic energy contribution + potential energy contribution =

total energy contribution

â€¢Quantization of Energy: an electron can possess only certain discrete energies in an

atom.

â€¢Probability Density: probability of finding an electron at a given location is

proportional to the square of the wavefunction at that point, Î¨2.

â€¢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.

â€¢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).

â€¢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,â€¦, n-1) ïƒ 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 â€“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, +Â½ for anticlockwise spin (spin-up) and â€“Â½ for

clockwise spin (spin-down).

oNodes: the positions where either component of the wavefunction passes

through zero.

ï‚§Radial Nodes: occur where radial component of the wavefunction

passes through zero.

ï‚§Angular Nodes: occur where the angular component of the

wavefunction passes through zero.

ï‚§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 â€“ l â€“ 1 radial nodes.

oRadial Distribution Function (P(r)) = r2R(r)2, where r is the radius.

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