Chapter 11 Modern Atomic Theory
You should be able to:
 Understand the Bohr model of the atom
 Understand the concept of electrons in shells and the use of quantum numbers
 Understand the use of the terms s, p, d and f and their use in orbital notation
 Recall and understand the rules for filling orbitals (Aufbau, Pauli and Hund) and
determining electronic configuration including the Pauli exclusion principle,
Hund's rule of maximum multiplicity and notable exceptions
 Be able to construct the electronic configuration of the elements using the s, p and
d and f notation
 Be able to construct the electronic configuration of the elements using the noble
gas core
 Be able to construct the electronic configuration of simple ions (including d block
ions)
 Recall the shapes of the s, p and d orbitals
 Recall that orbitals are electron probability maps
 Be able to describe electronic configurations using the electrons in boxes notation
 Recall the meanings of the terms paramagnetic, diamagnetic and isoelectronic
 Know what are meant by the terms, "group" and "period", when applied to the
periodic table
 Be able to recall the group names of groups 1, 2, 17 and 18
 Understand that regular, repeatable patterns occur in the periodic table
 Appreciate that these patterns sometimes have notable exceptions
 Recall and understand that the noble gases have full outer shells that represent
stable electronic configurations
 Recall the definition of ionization energy
 Recall the definition of electron affinity
 Recall and understand the variation in ionization energy when moving about the
periodic table
 Be able to predict the group an element is in from ionization energy data
 Recall how and why atomic and ionic size vary when moving about the periodic
table
 Understand how many physical properties change gradually when moving about
the periodic table
Vocabulary- electromagnetic radiation wavelength frequency
photon
quantized energy levels
wave mechanical model
orbital
sublevels
principal energy levels
Pauli exclusion principle
electron configuration
orbital (box) diagram
valence electrons
core electrons
metals
lanthanide series
actinide series
nonmetals
metalloids (semimetals)
main group (representative) elements
atomic size ionization energy
11.1 Rutherford’s Model of the Atom (fig. 11.1, p. 323)
Describe:
11.2 Energy & Light
Electromagnetic Radiation (fig. 11.4, p. 325)-
Draw fig. 11.3, p. 324
Draw fig. 11.4, p. 325
WavelengthFrequencySpeed of lightPhotons11.3 Emission of Energy by Atoms
Excited stateGround state
A quantum package of light is emitted when an electron loses energy and goes from an
excited state to a less excited state (maybe all the way down to its ground state).
Draw fig. 11.8
11.4 The Energy Levels of Hydrogen
Different wavelengths of light carry different amounts of energy per photon
Red light- less energy
Blue light- more energy
Draw fig. 11.10, p. 328.
Energy levels are quantized- only certain energy levels are allowed. Why?
11.5 The Bohr Model of the Atom
Niels Bohr- Bohr model of atom (only really works for Hydrogen. Why?)
Electrons are not really in circular orbits around the nucleus (even though it is convenient
to draw them this way)
Draw fig. 11.17, p. 329
11.6 The Wave-Mechanical Model of the Atom
E. Schrodinger & L.V. de BroglieWave-mechanical model-Electrons can behave with wave properties
OrbitalHeisenberg Uncertainty Principle-
11.7 The Hydrogen Orbitals
Principal Energy Levels- 1, 2, 3, …
(fig. 11.21, p. 334)
Sublevels1st principal energy level has 1 sublevel
2nd principal energy level has 2 sublevels
3rd principal energy level has 3 sublevels, etc.
The first sublevel is always designated _____
The second is designated_____; The third is designated _____;The Fourth is designated
_____. (no need for any more…)(why?)
s has 1 orbital; p has ___ orbitals; d has ___ orbitals; f has ___ orbitals.
(see figures 11.24, 11.25, 11.26, 11.27, and 11.28, p. 334-336)
Shape of s
shapes of p
shapes of d
11.8 The Wave Mechanical Model: Further Development
Pauli Exclusion Principle-
Aufbau Filling Order1s2 2s2 2p6 3s2 3p6 ** 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
** overlap occurs
Example: True or False?
_____ an s orbital is always spherical in shape
_____ the 2s orbital is the same size as the 3s orbital
_____ the number of lobes on a p orbital increases as n increases; that is, a 3p orbital has
more lobes than a 2p orbital
_____ level 1 has one s orbital; level 2 has two s orbital; level 3 has three s orbitals, and
so on.
_____ the electron path is indicated by the surface of the orbital
11.9 Electron Arrangements in the First Eighteen Atoms on the Periodic Table
Hund’s Rule(see Carbon, Chromium, and Copper)
Element At #
#p
#e
elec conf
orbital diagram
H
at. #1
1p
1e1s1
He
at. #2
2p
2e1s2
Li
at. #3
3p
3e1s22s1
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Valence electronsCore (noble gas inner core) electronsKernel & shorthand elec. configuration
Elements in same vertical group have same # of valence electrons (see the periodic table
we marked up)
Orbital DiagramsAn orbital is a box; an electron is represented by an arrow, one up and one down
(indicates opposite spins). No more than 2 arrows can be in a box. S orbital has 1 box; p
has 3; d has 5; f has 7.
Example 11.2,, p. 342- write the complete electron configurations and orbital diagrams
for elements aluminum through argon.(do above in the table under 11.9)
11.10 Electron Configurations and the Periodic Table
K
Ca
Cr (expected)
Cr (actual)
Cu (expected)
Cu (actual)
Lanthanide series
Actinide series
11.11 Atomic Properties & the Periodic Table
MetalsNonmetalsMetalloids (semimetals)Atomic size- size of atoms generally decrease as you go across the period from left to
right (increasing atomic number)(increased # of protons draws electrons in tighter)
Atomic size generally increases as you go down a vertical group. (higher principal energy
levels are further from the nucleus)(fig. 11.36, p. 350)
Ionization Energy- amount of energy needed to remove an electron from an atom of the
element in the gaseous state. (p. 352)(see charts/graphs in class)
(successive ionization energies)