PERIODICITY and ATOMIC STRUCTURE

1
PERIODICITY and ATOMIC STRUCTURE

• Electromagnetic Radiation (1-3)
• Quantum Mechanics (4-8)
• Electronic Configurations (9-15)

2
STRUCTURE OF ATOM

• Electrons and Protons observed and characterized
• Models of the Atom - pre-20th Century
• Raisin Pudding J.J. Thomson
• Small positive nucleus surrounded by a lot of
  empty space through which the electrons are
  dispersed Rutherford (1911)
• Atom nucleus also includes neutrons Chadwick
  (1932)

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LIGHT or EM RADIATION WAVE

• Classical science (CS) considered electromagnetic
  radiation (EM) or light as a wave based on
  observations of diffraction, reflection,
  interference, refraction.
• Form of energy
• Wavelength, ? c/? nm
• Frequency, ? Hz 1/s
• Speed, c ? ? 3.00E08 m/s

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ELECTROMAGNETIC SPECTRUM

• Light or electromagnetic radiation spans many
  orders of magnitude in E, ?, and ?.
• Figure 5.3
• Visible ROY G. BIV 400-800 nm
• At lower E and ? , ? increases Infrared,
  microwave, radiowave
• At higher E and ? , ? decreases Ultraviolet,
  X-rays, gamma-rays

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ELECTRONS, PROTONS, NEUTRONS PARTICLES

• Classical science considered these subatomic
  particles to be particles with mass (m), velocity
  (v) and momentum (mv)
• It was assumed that an object was either a wave
  (light) or a particle (electron).

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FROM CLASSICAL TO QUANTUM THEORY

• From the late 1800s to the 1920s, many
  experimental observations that could not be
  explained by Classical Science/Theories were
  recorded. These led to the development of
  Quantum Mechanics
• atomic line spectra (Balmer,1885)
• properties of radiation from heated solid or
  blackbody (Planck, 1900)

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TRANSITION (2)

• photoelectric effect (Einstein, 1905)
• heat capacity of solids
• electron diffraction
• As scientists worked to understand these exptal
  results, several conclusions emerged
• Electrons have WAVE and particle properties.
• Light has PARTICLE and wave properties.
• deBroglie Eqn expresses this ? h/mv

8
ATOMIC LINE SPECTRA

• CS Rutherford model of the atom.
• Expt When atoms are excited, they return to
  their stable states by emitting light. This
  light can be recorded to produce an atomic
  spectrum. Early experiments showed that the
  spectra consists of lines and that atoms from
  different elements gave different line spectra.
  Fig 5.6
• What do these spectra tell us about the structure
  of the atom?

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ATOMIC LINE SPECTRA (2)

• Balmer measured the emission spectrum of H and
  fit the observed wavelengths of the emitted light
  to an equation
• ? Rc (1/22 1/n2) where R Rydberg constant
  1.097E-2 1/nm
• The emission lines of the H atom in other regions
  of the EM spectrum fit the Balmer-Rydberg Eqn ?
  Rc (1/m2 1/n2) for n gt m n and m are
  integers. This is an empirical eqn.

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BOHR ATOM (Fig 5.14)

• The Rutherford model could not explain these
  results, but Bohrs planetary model of the atom
  could (1914) .
• This model led to quantized electronic energy
  levels and to an eqn consistent with the
  Balmer-Rydberg Eqn.
• The energy of an electron in the nth energy level
  is quantized and equals
• En - hcRZ2/n2 where n 1, 2, 3...

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BOHR ATOM (2)

• Then when an electron goes from one quantized
  level (n) to another (m), light is emitted or
  absorbed with a wavelength defined by 1/?
  R(1/m2 - 1/n2 )
• The Bohr atom is the basis for the modern theory
  of the atom but it has limitations. For example,
  it is only accurate for 1-electron atoms and ions.

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QUANTUM MECHANICS (Schrodinger, 1926)

• The QM model of the atom replaced the Bohr model.
  This model is based on electrons wave
  properties.
• The electron in an atom was viewed as a standing
  wave around the nucleus.
• These standing waves (?) are called wave
  functions and are interpreted as the allowed
  atomic orbitals for electrons in an atom.

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QUANTUM MECHANICS (2)

• The goal of QM is to solve the Schrodinger Eqn
  and find ? plus its associated (quantized)
  energy.
• ? 2 is related to the probability of finding an
  electron at a particular (x,y,z) location.
• Heisenberg Uncertainty Principle (1927) states
  that we cannot know the position and momentum of
  an electron (considered a wave) exactly
• ?x ?(mv) h/4 p

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QUANTUM MECHANICS (3)

• When the ?s are found, each one is defined by
  three quantum Numbers.
• A set of ?s lead to atomic electronic
  configurations.
• QM is the basis for understanding chemical
  bonding and molecular shapes (Chap.7), chemical
  reactions, physical and chemical properties
  (Chap. 6).

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ATOMIC ORBITALS AO) and QUANTUM NUMBERS (QN)

• AOs are wavefunctions, ?, and are characterized
  by QNs which are related to each other.
• These relationships determine the identity and
  number of AOs in an atom.
• Principal QN, n 1, 2, 3(K, L, M...shell)
  determines energy (quantized) and size of atomic
  orbital.

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AOs and QNs (2)

• Angular momentum QN, l 0, 1, 2n-1 (s, p, d)
  subshell determines shape of atomic orbital.
  For each n value, there are n lvalues.
• Magnetic, ml - l, -2, -1, 0, 1, 2, l
  determines spatial orientation of orbital. For
  each l value, there are 2l 1 ml values.
• Spin, ms 1/2, -1/2 determines orientation
  of electron spin axis.

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AOs and QNs (3)

• There are relationships (limitations) between
  four quantum numbers (Table 5.2)
• For the H atom and other one-electron atoms, all
  AOs with the same n value have the same energy.
  This is called energy degeneracy. (Fig 5.9)

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ATOMS WITH 2 ELECTRONS

• We will apply the 1-electron results to the
  many-electron atom.
• For the 1-electron atom, AO energy depends only
  on n and as n increases, energy increases
  (becomes less positive). So AOs can be ordered
  from low to high energy 1s lt 2s, 2p lt 3s,
  3p, 3d...

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MULTI-ELECTRON ATOMS

• For the many-electron atom, energy depends on n
  and l 1s lt 2s lt 2p lt 3s lt 3p, etc. See Fig 5.9
• This is due to electron-electron repulsions.
• Also, because of the difference in AO shape for
  different l values, electrons with the same n but
  in different subshells experience different
  attractive forces to the nucleus.
• Zeff Zactual electron shielding