CH 7 ATOMIC
STRUCTURE AND
PERIODICITY
AP Chemistry
2014-2015
TERMS 7.1-7.4
Electromagnetic radiation: radiant
energy that exhibits wavelike behavior
and travels through space at the speed
of light in a vacuum
Wavelength: the distance between two
consecutive peaks or troughs in a wave.
Frequency: the number of waves in a cycle
that pass a given point in space each second.
TERMS
Planck’s constant: the constant relating the
change in energy for a system to the
frequency of the electromagnetic radiation
absorbed or emitted; equal to 6.626 x 10 -32 Js.
Quanta
Quantization is the concept that energy can
only occur in discrete packets, called
quantum. “Quanta” is the plural of quantum.
Photon: a quantum of electromagnetic
radiation.
Photoelectric effect: the phenomenon in which
electrons are emitted from the surface of a
metal when light strikes it.
 This light must meet or exceed a threshold frequency
in order for the effect to occur.
 If the effect occurs, the number of electrons and
kinetic energy of the electrons emitted both increase
as the intensity of the light increases (for KE, the
relationship is linear)
Dual nature of light: light exhibits both
wave-like and particulate characteristics.
 A photon only has mass in a relative sense; it
cannot be weighed but it exhibits mass
experimentally.
Diffraction: the scattering of light from a
regular array of points or lines.
 Diffraction patterns are patterns of bright spots and
dark areas that form when scattered light interferes
constructively and destructively (respectively). This
phenomenon demonstrates that particles such as
electrons have wavelengths.
All matter exhibits both particulate and wave
properties.
EMR has been shown to behave like a particle;
electrons have been shown to have wavelengths.
Science is solved, everyone go home.
Atomic spectra
 The continuous spectrum contains all the wavelengths of
visible light and is obtained by passing white light
through a prism.
 Line spectra
 An emission spectrum shows lines with wavelengths
corresponding to discrete energy levels in a substance.
 The opposite of an emission spectrum is an absorption
spectrum, which shows light at all wavelengths except the
wavelengths corresponding to discrete energy levels in a
substance.
 The wavelengths at which light can be absorbed or emitted
correspond to the energy levels that are allowed for a
substance, which in turn are related to its electron structure.
Neat! So electron structure determines the energy found in
each possible quantum which determines color!
Quantum model (the Bohr Model)
Niels Bohr developed a quantum model for the
hydrogen atom (see Equations section). The
meaning of this model is that the electron in a
hydrogen atom moves around the nucleus only
in certain allowed circular orbits.
This model only works for the electron in
hydrogen. It does not work for any other
element.
Ground state: the lowest possible energy
level for an electron in an atom.
EQUATIONS
 λν = c
where c is the speed of light, 2.9979 x 10 8 m/s
 ΔE = nhν
where n is an integer (n = 1 corresponds to a
single quantum)
hc
 ΔE photon = hν =
λ
“packets”, or quanta
h
Used for calculating energy
λ=
de Broglie’s equation; allows us to
mν
calculate the wavelength for a particle
 (derivation shown on pages 302 -303 of your book)
 E = -2.178 x 10 -18 J (Z 2 /n 2 )
Used to determine the
energy levels available to the electron in the
 hydrogen atom; Z = nuclear charge, n = an integer; 0 for a
free electron (an electron an infinite distance from the
nucleus)
PRACTICE PROBLEMS
 1 , 2, 3, 4
The brilliant red colors seen in fireworks are due to the
emission of light with wavelengths around 650 nm
when strontium salts such as Sr(NO3)2 and SrCO3 are
heated. Calculate the frequency of red light of
wavelength 6.50 × 102 nm.
TERMS 7.5-7.8
Quantum mechanics (or wave mechanics) is
the result of Heisenberg, de Broglie, and
Schrodinger’s efforts to find a model superior
to the Bohr model. This model emphasizes the
wave properties of the electron.
 The electron bound the nucleus is similar to a
standing wave (a stationary wave; its ends are fixed).
There are limitations on the allowed wavelengths of
standing waves since they have definite length.
Nodes (N) have zero displacement while antinodes (A)
exhibit the amplitude of the wave. There must be a
whole number of half wavelengths in any of the
allowed motions of a standing wave.
 Only certain circular orbits
in an atom have a
circumference into which a
whole number of
wavelengths of a standing
electron will “fit”.
 All other orbits produce
destructive interference and
are not “allowed”.
Schrodinger used this model
to derive the equation to the
right. The Ψ (psi) is called
the wave function. A specific
wave function (solution) for Ψ
is called an orbital.
 The wave function does not
actually tell us how an
electron moves.
 The Heisenberg uncertainty principle states that there is a
fundamental limitation to just how precisely we can know both
the position and momentum of a particle at a given time. The
more accurately we know about a particle’s position, the less
accurately we can know its momentum, and vice -versa.
 This limitation is negligibly small for large objects like baseballs
but is much more significant for small particles like electrons -which is another inadequacy of the Bohr model (it assumes we
know more than we can).
 Δx = uncertainty in a particle’s position, Δp = uncertainty in a
particle’s momentum (also calculated as Δmv), h = Planck’s
constant 6.626 x 10 -34 Js
 The square of a wave function indicates the
probability of finding an electron near a particular
point in space (probability distribution).
 A graph of the total probability of finding an electron
in each spherical shell vs. distance from the nucleus
is called its radial probability distribution. The
maximum in this curve occurs because of two
opposing effects: the probability of finding an
electron at a particular position is greatest near the
nucleus, but the volume of the spherical shell
increases with distance from the nucleus. Therefore,
the total probability increases to a certain radius and
then decreases past it.
 The size of an orbital (as we usually use the term) is
the radius of the sphere that encloses 90% of the
total electron probability .
EACH ORBITAL IS CHARACTERIZED BY A
SET OF QUANTUM NUMBERS.
 Principal quantum number (n)
 Has integral values 1, 2, 3…
 Related to the size and energy of the orbital; as n increases, energy increases
 Angular momentum quantum number (l)
 Has integral values from 0 to n - 1 for each value of n.
 Related to shape of atomic orbital




s orbital
p orbital
d orbital
f orbital
l=0
l=1
l=2
l=3
 Each set of orbitals with a given value of l is called a subshell. These
subshells are described using the principal quantum number and the angular
momentum quantum number.
 Magnetic quantum number (m l )
 Has integral values between l and -l, including zero.; related to the orientation
of the orbital in space relative to the other orbitals in the atom.
 Spin quantum number (m s )
 ½ or -½ (positive is “filled” before negative); two electrons with opposite spin
can occupy any one orbital.
Orbital Shapes and Energies
 Nodal surface: an area of low
probability in
a probability distribution (also
called a node).
 s orbital = spherical, found in all n
 p orbital = peanut, found in n >1
 d orbital = daisy, found in n > 2
 f orbital = fancy, found in n >3
All orbitals with the same n have the
same energy and are said to be
degenerate.
All 3 p orbitals in a given level have the
same energy as each other, etc.
Electron spin and the Pauli Exclusion
Principle
An electron has a magnetic moment with two
possible orientations when the atom it is
found in is placed in an external magnetic
field. Therefore, electrons have two (opposite)
spin states.
Pauli Exclusion Principle: in a given atom no
two electrons can have the same set of four
quantum numbers.
PRACTICE PROBLEM
5
TERMS 7.9-7.11
 Polyelectronic atoms: atoms with more than one
electron
 Three energy contributions are considered in the description of
polyelectronic atoms
 KE of the electrons as they move around the nucleus
 PE of attraction between the nucleus and the electrons
 PE of repulsion between two electrons
 Since the results of the Schrodinger equation cannot be solved
exactly, electron repulsion cannot be calculated exactly. This is
called the electron correlation problem, and requires that we
make approximations in describing the movements of
electrons.
 Shielding (or screening) occurs when an electron in a
polyelectronic atom is repelled away from the
nucleus due to the other electrons in the atom. This
leads to the electron binding to not be as tight as it
would have been if the atom was not polyelectronic.
 When electrons are placed in a particular quantum
level, they “prefer” the orbitals in the order s, p, d,
and then f. This is due to the energy of each
sublevel (s = lowest, f = highest).
 Aufbau principle: as protons are added one by one to
the nucleus to build up the elements, electrons are
similarly added to orbitals (starting with the lowest energy orbital and working upwards). By the way,
“aufbau” is German for “building up” and not
someone’s last name.
Hund’s rule: the lowest energy
configuration for an atom is the one having
the maximum number of unpaired
electrons allowed by the Pauli Exclusion
Principle in a particular set of degenerate
orbitals. (these images are called orbital
diagrams)
Electrons can be
grouped together in
many ways, including
Valence electrons, which
are found in the outermost
principal quantum level of
an atom; important to
chemists because they are
involved in bonding; the
elements in the same
group have the same
number of valence
electrons
Core electrons, the inner
electrons
 Electron configuration: description of electronic structure
of an atom that uses the principal quantum number, the
angular momentum quantum number, and the number of
electrons in a subshell. Can be written longhand or
using noble gas configuration. When writing the electron
configuration for an ion, take the number of electrons
lost or gained into account.
Orbital diagram:
description of
electronic structure of
an atom that displays
the number electrons
in each orbital.
Cannot be written
shorthand.
Parts of the periodic table to know
Group 1A = alkali metals
Group 2A = alkaline earth metals
Group 7A = halogens
Group 8A = noble gases
d-block = transition metals
f-block = lanthanides (4f) and actinides (5f)
s-block + p-block = representative elements
PRACTICE PROBLEMS
 6 and 7
TERMS 7.12
 Ionization energy is the energy required to remove an electron
from a gaseous atom or ion (assumed to be in its ground state).
 The first ionization energy, I 1 , is the energy required to remove the
highest-energy electron of an atom.
 The second ionization energy, I 1 , is the energy required to remove the
next highest-energy electron of an atom after the highest -energy electron
has been removed.
 Ionization energy increases from left to right, and decreases going down
a group.
 Electrons added in the same principal quantum level do not completely shield
the increasing nuclear charge caused by added protons. Thus electrons in the
same principal quantum level are generally more strongly bound as we move
to the right on the periodic table.
 The main reason for the decrease in ionization energy in going down a group is
that the electrons being removed are, on average, farther from the nucleus.
Electron affinity is the energy change
associated with the addition of an electron
to a gaseous atom. If the addition is
exothermic, the corresponding value for
electron affinity will carry a negative sign.
Becomes more exothermic from left to right
across a period until group 8A.
Sort of becomes less exothermic down a group,
but there are a lot of exceptions and the change
going down a group is relatively small.
 The atomic radius is measured indirectly (by measuring
the distances between atoms in chemical compounds —
often called covalent atomic radii because of the way
they are determined; the radii for metal atoms, called
metallic radii, are obtained from half the distance
between metal atoms in solid metal crystals).
 Smaller than might be expected from the 90% electron density
volumes of isolated atoms, because when atoms form bonds, their
electron “clouds” interpenetrate.
 Decreases from left to right across a period, due to increase in
effective nuclear charge.
 Increases down a group because of the increases in the orbital
sizes in successive principal quantum levels.
EXERCISES
 8, 9, 10, 43