AP Chemistry: Unit 7 : Atomic Structure and Periodicity
How does atomic structure affect behavior?
Electromagnetic Radiation
One of the ways that energy travels through space is by electromagnetic radiation.
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Wave Characteristics
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Wavelength (λ) –
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Frequency (ν) —
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Speed (c) –
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The Nature of Matter
Wave and Particle Duality
Planck found that matter could only absorb or emit energy in whole number multiples of
the quantity hν.
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h
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ΔE =
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Transfer of energy is not continuous but is quantized and can occur only in discrete
amounts called quantum. Thus energy has particle properties as well as wave
properties.
Wave and Particle Duality
Einstein proposed that electromagnetic radiation was also quantized and could be viewed
as a stream of “particles” called photons.
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Ephoton
The Photoelectric Effect
The photoelectric effect refers to the phenomenon in which electrons are emitted from the
surface of a metal when light strikes it.
1.
2.
3.
4.
These observations can be explained by assuming that electromagnetic radiation is
quantized (consists of photons), and that the threshold frequency represents the minimum
energy required to remove the electron from the metal’s surface.
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KEelectron =
2
Planck and Einstein Conclusions
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Electromagnetic radiation, which was previously thought to exhibit only wave
properties, seems to show certain characteristics of particulate matter as well. This
phenomenon is sometimes referred to as the dual nature of light.
The main significance of the equation E = mc2 is that energy has mass.
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Louis de Boglie (1892-1987)
Since light, which previously was thought to be purely wavelike, was found to have
certain characteristics of particulate matter. But is the opposite also true? Does
matter have that is normally assumed to be particulate exhibit wave properties?
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de Broglie’s equation allows us to calculate the wavelength for a particle:
de Broglie’s Proof
Conclusion: Energy is really a form of matter, and all matter shows the same types
of properties. All matter exhibits both particulate and wave properties.
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The Atomic Spectrum of Hydrogen
Spectrum
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A continuous spectrum
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An emission spectrum produces
Hydrogen Line Spectrum
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The significance of the line spectrum is that it indicates that only certain energies
are allowed for the electron in the hydrogen atom. In other words the energy of the
electron in the hydrogen atom is quantized
The Bohr Model
Niels Bohr (1885 – 1962)
Bohr developed a quantum model for the hydrogen atom that allowed for only
specific energy levels around the atom that corresponded with specific radii.
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The most important equation to come from Bohr’s model is the expression for the
energy levels available to the electron in the hydrogen atom.
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Example: What is the change in energy if an electron in level 6 (excited state)
returns to level 1 (ground state) in a hydrogen atom?
What is the corresponding wavelength for the energy produced from the electron
jump?
Bohr Model Conclusions
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The model correctly fits the quantized energy levels of the hydrogen atom and
postulates only certain allowed circular orbits for the electrons.
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As the electron becomes more tightly bound, its energy becomes more negative
relative to the zero-energy reference state. As the electron is brought closer to the
nucleus, energy is released from the system.
The energy levels calculated by Bohr closely agreed with the values obtained from
the hydrogen emission spectrum but does not apply well to other atoms. The Bohr’s
model is fundamentally incorrect but is very important historically because it paved
the way for our current theory of atomic structure.
The Quantum Mechanical Model of the Atom
Quantum Mechanics
Quantum Mechanics or Wave Mechanics were developed by three physicists:
Heisenberg, de Broglie, and Schrodinger.
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Emphasis was given to the wave properties of the electron.
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The electron bound to the nucleus behaves similar to a standing wave.
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Like a standing wave, electrons can travel in patterns that allow for a common node.
In other words, wave patterns around the nucleus must be in whole number wave
patterns. But their exact movement is not known
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Heisenberg Uncertainty Principle
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There is a fundamental limitation to just how precisely we can know both the
position and momentum of a particle at a given time. This limitation is small for
large particles but substantial for electrons.
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Probability Distribution
A probability distribution is used to indicate the probability of finding an electron in
a specific position.
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For the hydrogen 1s orbital, the maximum radial probability occurs at a distance of
5.29x10-2nm or .529Å from the nucleus. This is the exact radius of the innermost
orbit calculated in the Bohr Model.
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The definition most often used by chemists to describe the size of the hydrogen 1s
orbital is the radius of the sphere that encloses 90% of the total electron probability
Quantum Numbers
Each orbital is characterized by a series of numbers called quantum numbers, which
describe various properties of an orbital:
1.Principal quantum number (n)-
2.Angular momentum quantum number (l) –
3.Magnetic quantum number (ml)
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4.Electron spin quantum number (ms)-
In a given atom no two electrons can have the same set of four quantum numbers (n,
l, ml , ms). This is called the Pauli exclusion principle; an orbital can only hold two
electrons, and they must have opposite spins.
Orbital Shapes and Energies
S orbitals
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The s orbitals have a characteristic spherical shape and contain areas of high
probability separated by areas of zero probability. These areas are called nodal
surfaces, or nodes.
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P Orbitals
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P orbitals each have two lobes separated by a node at the nucleus. The p orbitals
are labeled according to the axis of the xyz coordinate system along which the lobes
lie.
D Orbitals
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The five d orbitals first occur in energy level 3. They have two fundamental shapes.
Four of the orbitals (dxz, dyz, dxy, and dx2-y2) have four lobes centered in the plane
indicated in the orbital label. dx2-y2 lie along the x and y axes and dxy lie between the
axes. The fifth orbital dz2 has a unique shape with two lobes along the z axis and a
belt centered in the xy plane.
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F Orbitals
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The f orbitals first occur in level 4 and have shapes more complex than those of the
d orbitals. These orbitals are not involved in the bonding in any of the compounds
that we will consider.
Orbital Energies
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For the hydrogen atom, the energy of a particular orbital is determined by its value
of n. Thus all orbitals with the same value of n have the same energy – they are said
to be degenerate.
Polyelectronic Atoms
Polyelectronic atoms are atoms with more than one electron. To look at these
atoms, three energy contributions must be considered:
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Kinetic energy of the electrons as they move around the nucleus.
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The potential energy of attraction between the nucleus and the electrons.
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The potential energy of repulsion between the two electrons.
Since electron pathways are unknown, dealing with the repulsions between electrons
cannot be calculated exactly.
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The electron correlation problem occurs with all polyelectronic atoms. To deal with
this, we assume each electron is moving in a field of charge that is the net result of the
nuclear attraction and the average repulsions of all the other electrons.
In other words,…..
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A valence electron is attracted to the highly charged nucleus and still repelled by the
other ‘inner’ electrons. The net effect is that the electron is not bound nearly as
tightly to the nucleus as it would be if it were alone.
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Because of this shielded affect. orbitals within a principal energy level do not have
the same energy (degenerate). Sublevels vary in energy within a principal quantum
level.
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History of the Periodic Table
Early Greeks
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Earth
Air
Fire
Water
Dobereiner (1780-1849)
Johann Dobereiner was the first chemist to recognize patterns and found several
groups of three elements that have similar properties.
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chlorine, bromine and iodine
called triads.
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Newlands
John Newlands suggested that elements should be arranged in octaves, based on the
idea that certain properties seemed to repeat for every eighth element in a way
similar to the musical scale.
Meyer and Mendeleev
The present form of the periodic table was conceived independently by two
chemists: Meyer and Mendeleev. Usually Mendeleev is given most of the credit,
because it was he who emphasized how useful the table could be in predicting the
existence and properties of still unknown elements.
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In 1872 when Mendeleev first published his table, the elements gallium, scandium,
and germanium were unknown. Mendeleev correctly predicted the existence and
properties of these elements from the gaps in his periodic table. Mendeleev also
corrected the atomic masses of several elements.
The Aufbau Principle and the Periodic Table
Three rules for Orbital Configuration
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Aufbau principle – As protons are added, so are electrons, and fill in orbitals in order
of energy levels.
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Pauli Exclusion – Two electrons with opposite spins can occupy an orbital.
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Hund’s rule – The lowest energy configuration for an atom is the one with one
unpaired electrons in each degenerate orbital. (Electrons don’t like roommates)
Valence Electrons
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Valence electrons are the electrons in the outermost principal quantum level of an
atom. These are the most important electrons because they are involved in bonding.
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The elements in the same group have the same valence electron configuration.
Elements with the same valence electron configuration show similar chemical
behavior.
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Transition Metals
Transition metals have electron configurations that fill in the order of 4s before 3d.
Copper and Chromium have a configuration that is observed different than what is
expected.
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Additional Orbital Rules
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The (n+1)s orbital always fills before the nd orbitals. The s orbitals fill prior to the d
orbitals due to the vicinity of the nucleus.
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After lanthanum, which has the configuration of [Xe] 6s25d1, a group of 14 elements
called the lanthanide series, or the lanthanides occurs. This seris of elements
corresponds to the filling of the seven 4f orbitals.
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After actinium, a group of 14 elements called the actinide series or actinides occurs.
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The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence
electrons for the atoms in these groups.
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After actinium, a group of 14 elements called the actinide series or actinides occurs.
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The groups 1A, 2A, 3A…, the group numbers indicate the total number of valence
electrons for the atoms in these groups.
Periodic Trends in Atomic Properties
Ionization Energy
Ionization energy is the energy required to remove an electron from a gaseous atom or ion
when the atom or ion is assumed to be in its ground state:
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Ionization Energy
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It is always the highest-energy electron (the one bound least tightly) that is removed
first. The first ionization energy (I1) is the energy required to remove that first
electron. The second ionization energy (I2) is considerably larger.
The first electron is removed from a neutral atom, the second from a +1 cation. The
increase in positive charge binds the electrons more firmly and the ionization
energy increases. The trend continues for consecutive electrons removed.
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Electron Affinity
Electron Affinity is the change in energy change associated with the addition of an electron
to a gaseous atom:
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Electron Affinity
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If the addition of the electron is exothermic the corresponding value for electron
affinity will carry a negative sign.
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o As with Ionization energy. Some exceptions occur due to repulsions and
electron configuration.
Atomic Radius
Atomic radii are measured by the distances between atoms in chemical compounds.
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Covalent atomic radii are assumed to be half the distance between atoms in covalent
bonds.
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For metallic atoms, the metallic radii are obtained from half the distance between
metal atoms in a solid metal crystal
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The Properties of a Group: The Alkali Metals
Information and the Periodic Table
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It is the number and type of valence electrons that primarily determine an atom’s
chemistry
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The organization of the period table allows the prediction of electron configuration
without memorization.
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Groups on the periodic table have specialized names: Alkali metals, Alkaline earth
metals, Halogens, …etc.
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The most basic division of elements in the periodic table is into metals and nonmetals. This division affects chemical properties.
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Information and the Periodic Table
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Metalloids are elements along the division line and exhibit both metallic and
nonmetallic properties under certain circumstances. These elements are sometimes
called semimetals.
The Alkali Metals
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Lithium, sodium, potassium, rubidium, cesium, and francium are the most
chemically reactive of the metals. Hydrogen is found in group 1 but behaves as a
nonmetal because its very small and the electron is bound tightly to the nucleus.
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There is a smooth decrease in melting point and boiling points in Group 1 that is not
typical for other groups.
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o The hydration energy is greatest with Li+ because it has the most charge
density (charge per unit volume). This means that polar water molecules are
more strongly attracted to the small Li+ ions
o The order of reducing abilities in an aqueous reaction is Li > K > Na
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