Download Chapter 7 - UDChemistry

7.1: Electromagnetic Radiation
• Types of Electromagnetic Radiation: Gamma, X-ray, UV, visible,
IR, micro, radio
Properties of Electromagnetic Waves:
Wavelength ( λ ) :
• Distance Between two consecutive peaks or troughs in a wave
• Measured in meters
Frequency (ν)
• Number of waves that pass a given point per second
• Measured in Hertz
Speed (c)
• Speed of light
• Measured in meters/ second
Relationship Between Properties
Shortest wavelength = highest frequency
Longest wavelength = lowest frequency
INVERSE RELATIONSHIP
7.2 The Nature of Matter
Max Plank & Quantum Theory: Energy is gained/lost in whole numbers
multiples of the quantity hv ( frequency=v, Planck’s constant =h)
Planck’s Constant: h = 6.62606957 × 10-34 m2 kg / s ( J/s)
Planck discovered that energy is transferred to matter in packets of
energy called quantum, rather than energy of matter being continuous.
Einstein’s Photoelectric Effect
Phenomenon in which electrons are emitted from the surface of a metal
when light strikes it
His observations are explained by assuming electromagnetic radiation is
quantized (photons) and the threshold frequency is the minimum
energy required to remove the electron.
7.2: The Dual Nature of Matter
• Dual Nature of Light:
• Light travels through space as a wave
• Light transmits energy as a particle
• Particles have wavelength, exhibited by diffraction patterns
• De Broglie’s Equation: Allows calculation of wavelength for a particle
• λ = h/mv
• Diffraction: results when light is scattered from a regular array of
points or lines
• Diffraction Patterns: The interference pattern that results when a
wave or a series of waves undergoes diffraction, as when passed
through a diffraction grating or the lattices of a crystal. The pattern
provides information about the frequency of the wave and the
structure of the material causing the diffraction.
7.3: The Atomic Structure of Hydrogen
Continuous Spectrum: results when white light is passed
through a prism. Contains all wavelengths of visible light
Line Spectrum: only see a few lines, each of which
corresponds to discrete wavelength when passed thorough a
prism. (Hydrogen emission spectrum)
7.4: The Bohr Model
• Quantum Model : electron in a hydrogen atom moves
around the nucleus only in certain allowed circular orbits
• Bright line spectra confirms that only certain energies exist
in the atom, and atom emits photons with definite
wavelengths when the electron returns to a lower energy
state.
• Energy levels available to the electron in the hydrogen atom:
• n= an integer
• z= nuclear charge
• J= energy in Joules
7.4: The Bohr Model
• Calculating the energy of the emitted photon
• Calculate electron energy in outer level
• Calculate electron energy in inner level
• Calculate the change in the energy
• ΔΕ= energy of final state- energy of initial state
• hc/ ΔΕ : to calculate the wavelength of emitted photon
• Energy Change in Hydrogen atoms
• Calculate the energy change between any two energy levels:
• Limitations of the Bohr Model
• Bohr’s model does not work for atoms other than hydrogen
• Electrons do not move in circular orbits
7.5: Quantum Mechanical Model
• Electron bound to nucleus similar to standing waves
• The exact path of the electron is not known
• Heisenberg Uncertainty Principle- a limitation to the position and
momentum of a particle at a given time
Physical Meaning of ψ
- Square of the function is the probability of finding an electron near a
particular point
- Represented as a probability distribution
- aka electron density map, electron density, electron probability
Radial Probability Distribution
Since the orbital size cannot be calculated, the size of the orbital is the
radius of the sphere that an electron is in for 90% of the time
7.6: Quantum Numbers
Principal quantum number (n)
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•
•
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Main energy level
1, 2, 3, …
Size and energy of orbital
When n increases: orbital becomes larger, electron is further
from the nucleus, higher energy b/c electron is less tightly bound
to the nucleus so the energy is less negative
Angular momentum quantum number/Azimuthal QN (l)
• Sublevels, subshell
• 0...n-1 for each value of n
• Shape of atomic orbitals
Magnetic quantum number (𝒎𝒍 )
• Integral values from l to –l
• Orientation of the orbital in space
7.7: Orbital Shapes and Energies
s Orbitals
• Spherical shape
• Nodes for s orbitals of n=2 or greater
p Orbitals
• Two lobes each
• Occur in levels n=2 and greater
• Each orbital lies along an axis
Size of orbital:
• Defined as the surface that contains
9-% of the total electron probability.
• As n increases orbitals of the same
shape grow larger.
d Orbitals
• Occur in levels n=3 or greater
• Four orbitals with four lobes each centered in the plane indicated in the orbital label
• Fifth orbital has two lobes along z axis and a belt centered in the xy plane
f Orbitals
• Occur in levels n=4 and greater
• Complex shapes
• Usually not involved in bonding in compounds
Orbital Energies
• All orbitals with the same value of n have the same energy for hydrogen atoms
(Degenerate)
• The lowest energy state = ground state
• When the atom absorbs energy the electrons can move to higher energy orbitals
• “excited state”
7.8: Electron Spin & the Pauli Principle
• Electron Spin Quantum Number
• An orbital can only hold two electrons, must have
opposite spins.
• Spin can have +1/2 or -1/2
• Pauli Exclusion Principal
• In a given atom no two electrons can have the same set
of four quantum numbers
7.9: Polyatomic Atoms
Polyatomic Atoms: Atoms with more than one electron
3 energy contributions must be considered in description of the atom:
1) Kinetic energy of electrons as they move around the nucleus
2) Potential energy of attraction between nucleus and electrons
3) The potential energy of repulsion between the two electrons
Electron correlation problem: Electron pathways are not known, so
electron repulsive forces cannot be calculated exactly
• Average repulsions are approximated by...
• Treat each electron as it were moving in a field of charge that is
the net result of the nuclear attraction and average repulsions of
all other electrons
Screening or Shielding
• Electrons are attracted to the nucleus
• Electrons are repulsed by other electrons
• Electrons would be bound more tightly if other electrons weren’t
present
Closer proximity to the nucleus = lower energy
7.10 History of the periodic table
• Originally constructed to
represent patterns observed in
chemical properties of elements
• Mendeleev and Meyer both
independently conceived present
periodic table
• Mendeleev also corrected
several atomic masses
7.11: Aufbau Principle & the Periodic Table
Aufbau Principle: “As protons are added one by one to the nucleus to
build up elements, electrons are similarly added to these hydrogen like
orbitals”
Hunds Rule: “The lowest energy configuration for an atom is the one
having the maximum number of unpaired electrons allowed by Pauli
principle in a particular set of degenerate orbitals
Periodic Table Vocab:
Valence electrons: electrons in outermost principal quantum level of
an atom
Transition metals: “d” Block
Lanthanide and Actinide Series : “f” block
Representative Elements: Group 1A through 8A
Metalloids: Border between metals and nonmetals, exhibit
properties of both
7.12 Periodic Trends in Atomic Properties
Ionization energy: energy required to remove an electron from an
atomic (increase across period, decreases with increasing atomic
number within a group)
Electron affinity: energy change associated with the addition of an
electron (decrease down period, increase across period)
Atomic Radius: Determination of radius (increases down group,
decreases across period)
7.13: The Properties of a Group: Alkali Metals
• Easily lose valence electrons
• Reducing agents
• React with water
• Large hydration energy
• Positive ionic charge