Transcript Chapter 7
Atomic Structure and Periodicity
Section 7.1 Electromagnetic Radiation
• The electromagnetic spectrum organizes
waves by wavelength, frequency, and energy
l – wavelength- the length from a point on a wave
to a corresponding point later in the wave.
n – frequency- the number of times a full wave cycle
passes by a reference point in one second.
The electromagnetic spectrum
Relationships between l, n, and E
• Wavelength and frequency are inversely
proportional (a long wave will take longer to
pass by a reference point, thus making its
• Energy is directly related to frequency. Higher
frequency waves have more energy than lower
• The speed of a wave is constant in a vacuum.
(3.0 x 108 m/s = c, speed of light)
Show Me Problem
• A red light emits light of about 650 nm
wavelength. What is the frequency of the
650 nm 6.5 x 10-7m
(6.5x10-7)(n) = (3.0 x 108)
n = 4.61 x 1014 Hz
Section 7.2 The Nature Of Matter
• Max Plank
– Quantizes energy
• Matter can absorb energy, but only in whole
number ratios of the term hn.
DE = (h*n)
6.626 x 10-34 J*s
• Albert Einstein
– Quantizes Radiation
• E = mc2 … Energy has mass and velocity.
Electromagnetic Radiation must be made up of
particles called photons.
Duality of Light:
Electromagnetic radiation has the
capacity to behave both as a wave
and a particle.
• Louis DeBroglie
– Determines the Duality of Matter
• If waves act as particles, do particles act as waves?
• Set the Einstein and the Plank equations equal to each other.
mc2 = E = hc/l
m = h/cl
l = h/mn
Duality of Matter:
Electrons have the capacity to
behave both as a particle and a
How can we be certain?
• X Ray Crystallography
Different color and
shading patters appear as
the electron “waves”
with the x-rays that are
exposed to the crystal.
X Ray Diffraction Pattern of
2-terphenyl-4-yl-5-phenyl thiophene (PPPTP)
X-Ray Diffraction Pattern of Beryl.
Atomic Spectrum of Hydrogen
• Extensively studied by atomic theorists such as Bohr.
• High energy sparks cause hydrogen gas molecules (H-H) to
break apart suddenly, with some electrons in higher energy
levels than would be expected normally.
• As the electrons fall back to their ground states, energy is
• Each color in the spectrum relates to an electron in a different
• Plank’s equation can be used to determine the color of the light
produced or the energy of the electron that is being observed.
More on the atomic spectrum of hydrogen
Section 7.4 The Bohr Model
• Based upon the study of the Hydrogen Spectrum,
Bohr designs paths for electrons to travel while
orbiting the nucleus. [ORBITS]
• Each orbit corresponded to a different energy level.
E = 2.178 x10
energy of e-
18 Z 2
• An electron in a hydrogen atom moves from energy
level one to energy level 2. What is the change in
energy the electron experiences?
E1 = -2.178 x 10-18 (12/12)= -2.178 x 10-18
E2 = -2.178 x 10-18 (12/22)= -5.445 x 10-19
DE = E2-E1 = -5.445 x 10-19-(-2.178 x 10-18)
DE = 1.634 x 10-18 Joules
Endo or Exo?
Does this make
Section 7.5 The Quantum Mechanical
Model of the Atom
• Determine that if the electron acts as a standing
wave (a wave that is fixed in place), then there are
only certain orientations for it to exist without
causing destructive interference with itself.
Heisenberg Uncertainty Principle
• There is a fundamental limitation to just how
precisely we can know both the position and
momentum of a particle.
Dx * Dmn = h/4p
• The more certain you are of the location of an
electron, the less certain you can be of its
• The more certain you are of the momentum of an
electron, the less certain you can be of its position.
The Wave Equation
• Schrodinger’s wave equation is used to
define the location of electrons as waves.
• A wave function is called an orbital.
– ORBITALS = ORBITS
• Wave functions are impossible to visualize.
We picture the electron density map (aka:
electron probability diagram)
HY = EY
Section 7.6: Quantum Numbers
• Quantum numbers describe the properties of
Principal quantum 1, 2, etc
0 to (n-1)
l to -l
# of orbitals
-1, 0, 1
-1, 0, 1
-2, -1, 0, 1, 2
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
Common Orbital Shapes
Section 7.7 Orbital Shapes
• Areas of high probability are separate by
areas of low probability. (NODES)
• Degenerate orbitals have different orientation
or shape but the same ENERGY.
• Lowest available energy level for an electron
= ground state
• Higher energy levels than expected = excited
Section 7.8 Electron Spin and Pauli Principle
• Electrons exhibit a fourth quantum number.
• Electron spin quantum number (ms)
• Values of ½ and -½. Indicates magnetic moment
of electron. Electrons can only spin in one of 2
Pauli exclusion Principle:
In a given atom, no two electrons
Can have the same set of
3 Ways for Electron Configurations
Electron Configuration Rule Summary
• Electrons enter lowest energy orbitals first.
• Only two electrons per degenerate orbital.
• Electrons spread out among degenerate
orbitals before pairing.
Exceptions to the Configuration Rules
• A fully filled orbital is more stable than a
partially filled orbital.
• half-filled orbital is more stable than a
more/less partially filled orbital.
Copper and Chromium
Cr actual configuration:
Cu actual configuration:
Molybdenum and Silver
Mo actual configuration:
Ag actual configuration:
Europium and Americium
Eu actual configuration:
Am actual configuration:
Electron Configuration Diagrams
Noble Gas Notation
Expanded Titanium : 1s22s22p63s23p64s23d2
Noble Gas Titanium: [Ar] 4s23d2
Use the noble gas that comes before the
element as a benchmark, then tack on the
extra occupied orbitals.
Periodic Trends: Electron Configurations
• Counting down tells what energy orbital.
• Counting over tells how many electrons.
Periodic Trends Activity
Go to this excel sheet and click on the
bottom tab labeled “atom properties”
Periodic Trends: Atomic Size
Increasing Atomic Size
Periodic Trends: Ionization Energy
Increasing 1st Ionization Energy
Periodic Trends: Electron Affinity
Increasing electron affinity
• Negative ions indicate a gain of
• They are larger than the atom
from whence they are formed.
• Positive ions indicate a loss of
• They are smaller than the atom
from whence they are formed.