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PERIODICITY and ATOMIC STRUCTURE

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PERIODICITY and ATOMIC STRUCTURE Electromagnetic Radiation (1-3) Quantum Mechanics (4-8) Electronic Configurations (9-15) – PowerPoint PPT presentation

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Title: 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)

3
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

4
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

5
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).

6
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)

7
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?

9
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.

10
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...

11
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.

12
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.

13
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

14
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).

15
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.

16
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.

17
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)

18
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...

19
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
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