AP Chemistry: Chapter 6 – Electronic Structure & The Periodic Table
6.1 Light, Photon Energies, & Atomic Spectra______________________
-Light travels through space as a wave. Waves are made up of crests (which rise above
the midline) and troughs which sink below it. Waves have three primary characteristics:
wavelength, frequency, and amplitude. Two of these characteristics are of special focus
in regards to atomic structure:


Wavelength (greek letter lamda) – the distance from two
consecutive crests or troughs, most often measured in meters or
nanometers.
Frequency (greek letter nu or f ) – the number of wave cycles that
pass a given point in unit time, measured in Hertz (1 Hz =
1cycle/s).
-Wavelength and frequency are related by the following formula:
c = the speed of light (in a vacuum) = 2.998 x 108 m/s
-Visible light is a tiny portion of the electromagnetic spectrum, covering a wavelength
region of 400 nm to 700nm.
-Hundreds of years ago it was believed that all properties of light could be explained in
terms of it’s wave nature. Max Planck & Albert Einstein conducted a series of
experiments that discredited this idea. Today it is considered that light is generated as a
stream of particles called photons, whose energy E is represented by the following
equation:
E = hf =
E = energy in Joules (J)
h = Planck’s constant = 6.626 x 10-34 J*s
f = frequency = s-1 (reciprocal seconds)
Pg 132 – Example 6.1
Pg 133 – Example 6.2
Energy, Wavelength, & Frequency:
Classwork:
2
Homework: 4,6,8
6.2 The Hydrogen Atom_________________________________________
-Niels Bohr assumed that the hydrogen atom consists of a central proton around which an
electron moves in a circular path. He related that the electrostatic force of the proton for
the electron (effective nuclear charge) to the centrifugal force due to the circular motion
of the electron. Bohr was able to express the energy of that tome in terms of the radius of
the electron’s orbit.
-There are three points to be made with the Bohr Model:
1. Bohr designated zero energy as the point at which the proton and electron are
completely separated.
2. Ordinarily the hydrogen electron is in its lowest energy state, referred to as the
ground state (n=1). When an electron absorbs enough energy, it moves to a
higher, excited state. For hydrogen, the first excited state is n=2, then n=3.
3. When an excited electron drops back to its lower energy state it gives off energy
as a photon of light.
-Using these three points it is possible to relate the frequency of light emitted to the
quantum numbers, nhi and nlo, of the two states:
f = RH
h
RH = 2.180 x 10-18 J
h = 6.626 x 10-34 J*s
-All of the lines in the Balmer series come from transitions to n = 2 from higher levels (n
= 3,4,5,….)
-The Bohr modeled had a flaw in that it could only be applied to hydrogen. The idea that
the electron moved around the nucleus in a well defined orbit at a fixed distance from the
nucleus was not sufficient.
-Louis de Broglie reasoned that if light could behave has particles (photons) as well as
waves, then why can’t electrons behave like a wave. This idea led to the development of
wave mechanics, now called quantum mechanics. This deferred from the Bohr model in a
number of ways:


The kinetic energy of an electron is inversely related to the volume of the region
to which it is confined.
It is impossible to specify the precise position of an electron in an atom at a given
instant.
Pg 135 – Example 6.3
Bohr Model of Hydrogen:
Classwork:
10
Homework: 14
6.3 Quantum Numbers__________________________________________
-Erwin Schrodinger proposed a rather complex differential equation to express wave
properties of an electron in an atom. The Shrodinger equation can be solved,
approximately, for atoms with two or more electrons. There are many solutions for the
wave function, each associated with a set of quantum numbers.
-The principal quantum number:
 Symbolized by n, basically the energy level the electron is in.
 The energy of the electron depends mainly on this number. As n increases, so
does the value of the electron.
 n = 1,2,3,4,5,…….
-The orbital quantum number:
 Symbolized by l, basically represents the sublevel the electron is in s,p,d, or f.
 s: l = 0, p: l = 1, d: l = 2, f: l = 3
 l = 0,1,2,…… (n-1)
 Sublevels increase in energy as well. Sublevel 2p is higher in energy than sublevel
2s.
*The letters s,p,d, and f come from the adjectives used to describe spectral lines:
sharp, principal, diffuse, fundamental.
-The magnetic quantum number:
 Symbolized by ml, this determines the direction in space of the electron cloud
surrounding the nucleus.
 Sublevel s has only 1 possible orbital
 Sublevel p has 3 possible orbitals
 Sublevel d has 5 possible orbitals
 Sublevel f has 7 possible orbitals
 All of the orbitals in a sublevel have the same energy
s __
p __ __ __
d __ __ __ __ __
f __ __ __ __ __ __ __
-The spin quantum number:
 Symbolized by ms, this represents the electron spin.
 An electron has magnetic properties that correspond to those of a charged particle
spinning on its axis, either clockwise or counterclockwise.
 ms can equal +1/2 or -1/2
*The Pauli Exclusion Principle: No two electrons in an atom can have the same set of
four quantum numbers – it requires that only 2 electrons can fit into an orbital.
Pg 140 – Example 6.4
Pg 140 – Example 6.5
Energy Levels & Quantum #’s:Classwork: 18, 22, 27
Homework: 20,28-34 all, 38
6.5 Electron Configurations in Atoms_________________ ____________
-Electron configuration describes the arrangement of electrons.
For example: Ne 1s22s22p63s23p6
-Abbreviated electron configurations may be used by preceding with the noble gas for the
element. For example: [Ne] 3s23p4
-Electron Configuration & The Periodic Table
 The atoms of elements in a group of the periodic table have the same
distribution of electrons in the outermost principal energy level.
 The elements in groups 13-18 fill p sublevels (starts at 2p).
 The transition metals fill d sublevels (starts at 3d).
 The two sets of elements listed separately at the bottom of the table are
filling f sublevels (starts at 4f).
Pg 146 – Example 6.7
6.6 Orbital Diagrams of Atoms___________________ _______________
-Show how electrons are distributed using arrows (to represent electrons) and parentheses
or lines (to represent the orbitals).
-Hund’s Rule:
 When several orbitals of equal energy are available, as in a given sublevel,
electrons enter singly with parallel spins.
 In all filled orbitals, the two electrons have opposed spins.
 In accordance with this rule, within a given sublevel, there are as many half filled
orbitals as possible.
-Paramagnetic: If there are unpaired electrons present the solid will be attracted into the
field
-Diamagnetic: If the atoms in the solid contain only paired electrons it is slightly
repelled by the field.
Pg 148 – Example 6.8
Orbital Diagrams; Hund’s Rule:
Classwork: 40
Homework: 42, 46
6.7 Electron Arrangements in Ions ________________ _______________
-In general when a monatomic ion is formed electrons are added to or removed from
sublevels in the highest principal energy level. Many monatomic ions (especially the
main group elements) have noble gas configurations.
-Transition metal cations to the right of the scandium group do not form ions with nobelgas configurations, they would have to lose four or more electrons to do so. In transition
metals the outer s electrons are usually lost first to form positive ions.
For example: Mn:
Mn+2:
-In ions like Fe, electrons will be lost from 4s first then the 3d. This is usually referred to
as the “first in, first out” rule.
Pg 150 – Example 6.9
Electron Arrangment in Ions:
Homework: 50
6.8 Periodc Trends In The Properties of Atoms _____________________
-The Periodic Law: The chemical and physical properties of elements are a periodic
function of atomic number.
-Atomic Radius: one half the distance of closest approach between atoms in an elemental
substance.
 Decreases across a period from left to right on the periodic table
 Increases down a group on the periodic table
-Ionic Radius:
 Positive ions are smaller than the metal atoms from which they are formed
 Negative ions are larger than the nonmetal atoms from which they are formed
-Ionization Energy: the energy required to remove an electron.
 Increases across the periodic table from left to right
 Decreases down the periodic table
-Electronegativity: the ability to attract an electron.
 Increases across the periodic table from left to right
 Decreases down the periodic table
Pg 154 – Example 6.11
Trends in the Periodic Table:
Classwork: 54,56
Homework: 58,60