Honors Chemistry: Test Objectives for Chapter 4, Atomic Structure II
Name: ______________________________________________________
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Date: _________
Describe the wave properties of electromagnetic radiation. Know the speed of all
electromagnetic radiation (3.0 X 108 m/s).
Explain the mathematical relationship among the speed, wavelength, and frequency of
electromagnetic radiation, c = .
Describe how the wavelength and frequency differ in various regions of the electromagnetic
spectrum, such as IR, visible, and UV.
o As frequency increases, energy increases and wavelength decreases.
Describe developments that led to the modern theory of the atom: photoelectric effect, lineemission spectrum of hydrogen, & dual nature of light.
Use the equation E = h for electromagnetic radiation.
Describe a continuous spectrum. Explain what a line spectrum is, referring to the light emitted
or absorbed by an atom.
Describe the basic features of the Bohr model of the atom. Using reference tables, draw Bohr
diagrams for given elements.
o Electron is treated as a particle (classical treatment).
o Quantization of angular momentum is imposed.
o The orbit closest to the nucleus has the lowest energy, and the energy of the orbits
increases as the distance from the nucleus increases.
o Know the maximum capacity of each orbit (2n2).
o The orbits get closer together as the distance from the nucleus increases.
o Electrons can only exist on orbits, but can move among available orbits by absorbing or
emitting energy.
Define the terms: valence electrons and kernel.
Draw Lewis dot diagrams for atoms.
Understand that electrons absorb heat, light, or electric energy when they move to an excited
state (endothermic process) and release light energy (photons) when they return to the ground
state (exothermic process).
Explain the cause of spectral lines in terms of electrons and energy levels. (Spectral lines are
produced when electrons jump from a high energy level to a low energy level.)
Recognize that each line in the bright-line spectrum of an element can be used to calculate the
energy of a particular electron transition, using E = h or E = hc/.
Understand that the bright-line spectrum of an element can be used to determine if a given
element is present in a sample, by matching the spectral lines of the element to lines in the
emission spectrum of the sample.
Identify an element from a Bohr configuration (excited or ground state). Recognize the ground
state Bohr configuration of an element, using the Reference tables. Identify excited state Bohr
configurations for a given element.
List the successes and failures of the Bohr model.
Discuss deBroglie & the dual nature of matter. Appreciate that the electron exhibits properties
of both particles & waves.
Describe the basic features of the quantum (modern) model of the atom.
o Appreciate that wave mechanics uses the mathematical language of waves to describe
the electron.
o Quantization of electron energies arises naturally from the mathematical language of
waves.
Explain how the Heisenberg Uncertainty principle and the Schrodinger wave equation led to
the idea of atomic orbitals.
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o Describe the uncertainty principle and explain the limitations it places on defining the
simultaneous location and momentum of an electron.
o Explain the concepts of orbital, electron density, & probability in the quantummechanical model of the atom.
List the four quantum numbers, describe their significance and allowed values.
o 3 quantum numbers define an orbital
 n = principle quantum number. Describes energy of orbital and distance from
nucleus. Allowed values are n = 1, 2, 3, 4, …
 l = angular momentum quantum number. Describes shape of orbital. Allowed
values are l = 0, 1, 2, 3, … n-1
 For l = 0, called an s orbital – spherical
 For l = 1, called a p orbital – dumbbell shapes
 For l = 2, called a d orbital – mostly clover shapes
 For l = 3, called an f orbital – more complex shapes
 ml = magnetic quantum number. Describes orientation of orbital. ml = -l, … 0
… +l.
 if l = 0, ml = 0
 if l = 1, ml = -1, 0, 1
 if l = 2, ml = -2, -1, 0, 1, 2
 if l = 3, ml = -3, -2, -1, 0, 1, 2, 3
o 1 quantum number, ms, describes the spin state of an electron in an orbital. It can be
either + ½ or - ½ .
Relate the number of sublevels in each of an atom’s main energy levels (nth level has n
sublevels) , the number of orbitals per sublevel (s is 1, p is 3, d is 5, f is 7), and the number of
orbitals per main energy level (nth level has n2 orbitals and holds 2n2 electrons).
List the total number of electrons needed to fully occupy each principle energy level (2, 8, 18,
32, … 2n2).
For one-electron systems, all the sublevels of a given principle energy level have the same
energy.
For systems with more the one electron, the sublevels of a given principle energy level vary in
energy: : Ens  Enp  End  Enf for polyelectronic atoms.
Recall the rules for filling orbitals & determining electron configurations, including the Aufbau
principle, the Pauli exclusion principle, and Hund’s rule.
o Aufbau principle: as protons are added to the nucleus, electrons are added to the
orbitals. An electron will occupy the lowest-energy orbital that is available.
o Pauli exclusion principle: no two electrons in an atom can have the same set of four
quantum numbers. Since the spin quantum number has only two possible values, this
limits the occupancy of an orbital to two electrons.
o Hund’s rule or rule of maximum multiplicity: Electrons don’t pair up in an orbital until
they have to. All electrons in singly occupied orbitals must have the same spin state.
Write electron configurations for the atoms of any element using orbital notation, sublevel
notation, and, when appropriate, noble gas or kernel notation. Use either the periodic table or
the diagonal rule to write out configurations.
Write electron configurations for simple ions, including transition metal ions. Valence
electrons are lost first (not necessarily the same as the order of fill).
Compare and contrast the Bohr model and the quantum model of the atom.