AP CHAPTER SEVEN OUTLINE
ELECTRON CONFIGURATION AND THE PERIODIC TABLE

Electrons in the outermost “shell” of an atom are the valence electrons; their
number and location are the chief factors that determine chemical reactivity.
I.
Electromagnetic Radiation and Matter
 All electromagnetic radiation consists of oscillating perpendicular
electrical and magnetic fields that travel through space at the same rate
(the “speed of light”): 186,000 miles per second, or 2.998 x 108 m/s in a
vacuum.
 The wavelength (λ) is the distance between adjacent crests or troughs in a
wave.
 The frequency (υ) is the number of complete waves that pass a given point
in a second. The units of frequency is waver per second. 1 hertz (Hz) = 1
wave per second.
 Wavelength and frequency are inversely proportional – as the wavelength
increases the frequency decreases, and as the wavelength decreases the
frequency increases.
 The amplitude of a wave is its height, and the higher a waves amplitude
the more intense is the radiation.
 The speed of light = wavelength times frequency
 c = υλ
II.




III.




Planck’s Quantum Theory
Planck stated that when an atom in a hot object emits radiation, it does so only
in packets having a minimum amount of energy. He called a packet of energy
a quantum.
The energy of a quantum is equal to its frequency times a constant, h, which is
6.626 x 10-34 J•s.
Equantum = hυ
The energy per quantum of radiation increases as the wavelength gets shorter.
The Photoelectric Effect
Certain metals exhibit a photoelectric effect: they emit electrons when
illuminated by light of certain wavelengths.
For each photosensitive metal there is a threshold wavelength below which no
photoelectric effect is observed.
Albert Einstein said Planck’s quanta were massless particles of light and
called them photons.
To remove one electron from a photosensitive metal surface requires a certain
minimum quantity of energy, Emin. Since each photon has energy E =h υ, only
photons whose E is greater than Emin will be able to knock an electron loose.
IV.

V.
The Dual nature of Light
All forms of electromagnetic radiation appear to have either wave or particle
characteristics depending upon the experimental circumstances.
The Bohr Model of the Hydrogen Atom
 In 1913 Neils Bohr developed a mathematical model to explain the
behavior of excited atoms emitting quantized wavelengths of light.
 In the Bohr model of the hydrogen atom, the electron could circle the
nucleus in orbits of only certain radii, which correspond to specific
energies.
 Bohr referred to these energy levels as orbits and represented the energy
difference between any two adjacent orbits as a specified quantity of
energy.
 Each orbit was assigned an integer, n, known as the principal quantum
number.
 The energy of the electron and the size of its orbit increase as the value of
n increases.
 Any atom with its electrons in their lowest energy levels is said to be in its
ground state
 When the electron absorbs a quantized amount of energy and moves to an
orbit with n greater than 1, it is in an excited state
 Any excited state is unstable. When the electron returns to its ground state
it emits the energy it had gained in the form of light.
 ∆ E = Ef – Ei = - 2.179 x 10-18 J ( 1/nf2 – 1/ni2)

If nfinal is greater than ninitial energy is absorbed
 ν = ∆E/h = (2.179 x 10-18 J)/h x (1/ni2 – 1/nf2)
VI.
Beyond Bohr: The Quantum Mechanical Model of the Atom
 In 1924 De Broglie proposed the idea that electrons could have wave-like
properties
 λ = h/mv, where h is Plancks constant
 Heisenberg’s Uncertainty Principle states that it is impossible to
simultaneously determine the exact position and the exact momentum of an
electron
VII. Quantum Numbers, Energy Levels, and Orbitals
 The region in which an electron can be found within an atom is known as an
orbital.
 In the quantum mechanical model, the principal quantum number, n, is a
measure of the most probable distance of the electron from the nucleus.
 A collection of orbitals with the same principal quantum number, n, is called
an electron shell.
 Schrödinger determined that four quantum numbers are needed to describe the
three-dimensional coordinates of an electrons motion in the orbital of any
atom
1. First Quantum Number: n: Prinicpal Energy Level
This indicates the energy level of the orbital(s)
2. Second Quantum Number: l : subshells (s,p,d,f)
This indicates the shape of the orbital: s orbitals are spheres, p orbitals are shapted
like dumbbells; d and f orbitals are complex structures. The number of subshell
types is equal to n for each energy level.
3. Third Quantum Number: ml : relates to its orientation (along the x, y, or zaxis)
4. Fourth Quantum Number, ms : relates to the “spin” of the electron, clockwise
or counterclockwise, +½, -½.
 Pauli Exclusion Principle: No more than two electrons can be
assigned to the same orbital in an atom, and these electrons must have
opposite spins. This means that no two electrons in an atom can have
the same four quantum numbers.
VIII.





IX.



X.



XI.
Atom Electron Configuration
Electron configuration: the complete description of the orbitals occupied by
all the electrons in an atom or an ion
Aufbau “building up principle” :For an atom in its ground state, electrons are
found in the energy shells, subshells, and orbitals that produce the lowest
possible energy for the atom.
Hund’s rule states that electrons pair in an orbital only after each orbital in a
subshell is occupied by a single electron
Valence electrons are electrons in an atoms outermost shell; electrons in the
inner shells are called core electrons.
Groups 1A and 2A are s-block elements and groups 3A-8A are p-block;
transition metals are d-block elements
Periodic Trends: Atomic Radii
For the main group elements, atomic radii increase going down a group and
decrease going across a period
Transition metal have similar radii, so they are alike in their properties.
Effective Nuclear Charge: the nuclear positive charge experienced by the
outer shell electrons in a many electron atom. The effective nuclear charge is
less than the actual positive nuclear charge due to the shielding, or screening
effect of the core electrons.
Periodic Trends: Ionic Radii
The radius of a cation is always smaller than that of the atom from which it is
derived
The radius of an anion is always larger than that of the atom from which it is
derived
Positive and negative ions of elements in the same group increase in size
down the group
Periodic Trends: Ionization Energies


XII.

Ionization energy is the energy required to remove an electron from an atom
in its gaseous state.
For s- and p-block elements, first ionization energies (the energy required to
remove one electron from the neutral atom) generally decrease down a group
and increase across a period.
Periodic Trends: Electron Affinities
Electron affinity is a measure the attraction an atom has for an additional
electron. EA is the energy change when an electron is added to a gaseous
atom to form a -1 ion.