Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
Chapter 6
6.1 The Wave Nature of Light [p.212]
1. Quantum Theory2. Quantum Mechanics3. Electronic Structure- arrangement of electrons in atom
4. Much of the present understanding of electronic structure comes from the analysis of light
a. Light is part of electromagnetic spectrum like radio waves, microwaves, infrared, Xrays etc.
i. Form of radiant energy - carries energy through space
ii. Travels at 3.00 x 108 m/s in a vacuum
iii. Periodic
1. Wavelength – distance between two adjacent peaks
2. Frequency – cycles (wavelengths)/ second [unit hertz =s-1]
iv. Properties of Waves are due to the periodic oscillations ( length of wavelengths
measured in meters)
1. Long wavelength = low frequency
2. Short wavelength = high frequency
3. Equation illustrating the relationship
c=λv
c= speed of light
λ = wavelength
v = frequency
6.2 Quantized Energy and Photons [p.215]
1. Hot Objects and the Quantization of Energy
a. Planck studies help the understanding of the relationship between temperature, intensity
, and wavelength of radiation
b. Stated energy was released and absorbed in discrete “chunks” of energy called quantum
c. Equation
E=hv
E= energy of a single quantum
h = Planck’s constant
v = frequency
d. Planck stated that energy is emit and absorbed in whole-number multiples
e. Energy is quantized- restricted to certain amounts
2. The Photoelectric Effect and Photons
a. Photoelectric Effect is the emission of electrons by shining a light on the surface of a
metal
i. Min frequency for each metal needed to cause the effect
ii. Below min. no emissions
iii. Photon = tiny packet of energy
iv. Emission only occurs if there is a enough energy to overcome the attractive
forces that hold the metal together (work function)
6.3 Line Spectra and the Bohr Model [p.218]
1. Line Spectra
1
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
a. Monochromatic – single wavelength
b. Spectrum – multi- wavelength light is separated into the different wavelengths
c. Continuous Spectrum- ROYGBIV
d. Line Spectrum – individual wavelengths that represent the spectrum of an element
2. Bohr’s Model
a. Orbits have certain different energy levels
b. Electrons in a permitted orbit has a specific any that is “allowed”
c. Electron’s can absorb and emit energy as it travels form one orbital to another
3. The Energy States of the Hydrogen Atom
a. Each principle energy has a specific amount of energy
b. The more negative the energy, the lower the energy
i. Lower energy = ground state
ii. Higher energy = excited state
iii. Zero energy = energy needed to remove the electron form the atom n = ∞
c. Energy transitions for electrons in the exited state can be calculated using the following
formula.
ΔE = Ef – Ei = Ephoton= hυ
4. Limitations of the Bohr Model
a. Bohr’s model and explanation only applied to the hydrogen atom
b. No explanation for why electron stops from falling into nucleus
6.4 The Wave Behavior of Matter [p.222]
1. De Broglie – (wave characteristic of matter)
h
a. Equation  
mv
λ= wavelength
h = Planck’s constant
m = mass
v = velocity
b. De Broglie’s equation shows the inverse relationship between wavelength and mass.
c. The smaller the mass the more obvious the matter wave characteristics. The larger the
mass the more difficult it is to observe the matter waves.
2. The Uncertainty Principle- it is inherently impossible to know simultaneously to know the
exact momentum and exact location in space of extremely small masses due to their wave
characteristics. (Heisenberg)
6.5 Quantum Mechanics and Atomic Orbitals [p.224]
1. Background
a. Schrodinger’s equation incorporated the wavelike and particle-like structure of an atom.
b. Wave functions (ψ) – series of mathematical functions that describe the electrons in an
atom.
c. Probability density or electron density (ψ2) = the probability of finding an electron at
that location
2. Orbitals and Quantum Numbers
a. Orbitals – define a set of wave functions and corresponding energies
2
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
i. Each orbital has unique energies and shape
ii. Three quantum numbers are used to describe the orbital
1. the principal quantum number (n) – greater the number, the higher the
energy level
2. the angular momentum number (l) –
a. represented by an integral number from 0 to n-1 (n = the principal
energy level
b. Values are generally represented by corresponding letter
Values of l
0
1
2
3
Letter used
s
p
d
f
3. The Magnetic quantum number (m) –
a. integral value between –l and l
b. represents orientation of the orbital in space
iii. Electron Shell = the orbitals with the same principal energy level ( 3s, 3p, 3d)
iv. Subshells or sublevels = set of orbitals with the same n and l values (principal
and quantum number)
v. Ground State = electrons occupying the lowest energy level
vi. Exited State = an electron occupies a higher energy level.
6.6_Representations of Orbitals [p.228]
1. The s Orbitals
a. Spherically symmetric
b. l=0
c. ml= 0, only one “s” orbital
d. Radial Probability Function- indicates the probability of finding an electron any point
that is a distance “r” form the nucleus
e. Node-point at which there is zero probability of locating an electron
f. As “n” increases there is an increase in the size of the orbital
2. The p Orbitals
a. Dumbbell shaped (two lobes of the orbital are separated by a node)
b. n=2 or greater, l=1, ml = -1,0,1 (3 orbitals)
c. each orbital shaped like a dumbbell lies on one of the axes
d. increase in size as you move away from the axis
3. The d and f Orbitals
a. “d” orbitals
i. n= 3 or greater, l= 2, ml= -2,-1,0,1,2
ii. comprised of 5 different shapes, four are four leaf clovers
b. “f” orbitals
i. n= 4 or greater, l= 3, ml= -3,-2,-1,0,1,2,3
ii. comprised of 7 different shapes
6.7 Many-Electron Atoms [232]
1. Orbitals and their Energies
a. In many-electron atoms, for a given value of “n” , the energy of an orbital increases
with increasing value of “l”
b. Orbitals with in a subshell have the same energy
3
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
c. Degenerate – term used when orbitals have the same energy
2. Electron Spin and the Pauli Exclusion Principle
a. Electron Spin- causes each electron to behave as if it were a tiny sphere spinning on its
own axis
b. Spin magnetic quantum number (ms) indicates the spin on an electron as +1/2 or -1/2
c. The two opposite magnetic fields lead to the splitting of the spectral lines
d. Pauli Exclusion Principle –
i. no two electrons in an atom can have the same set of four quantum numbers n, l,
ml, and ms
ii. an orbital can hold a maximum of two electrons and they must have opposite
spins
6.8 Electron Configurations [234]
1. Background
a. Electron Configurations- the way electrons are distributed around an atom
b. Most stable electron configuration is the ground state
c. Orbitals are filled in order of the increasing energy, with no more than two electrons per
orbital
d. Orbital diagram – use of boxes and arrows to denote placement of the electrons
e. Paired – orbitals that contain two electron
f. Unpaired electrons- orbitals that contain one electron
2. Hund’s Rule
a. Customary to placement one electron as an up-spin
b. Filled shells of an energy level are very stable
c. Hund’s Rule- for degenerate orbitals, the lowest energy is attained when the number of
electrons with the same spin is maximized
i. Parallel spins- half-filled orbitals with the same electron spin
ii. Electron-electron repulsions is the reason for Hund’s rule
3. Condensed Electron Configurations
a. Ne has a complete 2p subshell making it very stable
b. Stable Eight – eight electron’s in the outermost shells
c. Condensed Electron configurations use the stable configurations of group 18 to shorten
the electron configuration
Na: [Ne]3s1
Li: [He]2s1
i. Core electrons – inner shell electrons
ii. Valence electrons- outer shell electrons
iii. Helps focus on the outermost shell which aids in seeing chemical similarities
4. Transition Metals
a. Argon is the core electrons for potassium which ends in 3p6
b. Potassium’s electron configuration exhibits the first sign of orbital overlapping because
the 4s fills before the 3d
c. Transition elements (transition metals) are the metals that fill the 3d subshell
i. follow Hund’s rule
ii. 4p is filled after 3d
5. The Lanthanides and Actinides
a. Lanthanide elements (rare earth metals)
4
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
i. Mark the filling of the 4f subshell
ii. Properties for these elements are very similar
iii. These elements occur together in nature
b. Actinide elements- mark the filling of 5f subshell
6.9 Electron Configurations and the Periodic Table [240]
1. Background
a. Periodic Table arranged so that elements in a column have the same valance electrons
b. Number of elements per row corresponds to the number of electrons in that energy level
i. 2 elements- two electrons in row 1
ii. 8 elements – 8 electrons in row 2 and row 3
c. Representative elements- s-block and p-block elements
d. Valance electrons for representative elements only are concerned with s and p subshells
2. Anomalous Electron Configurations- elements that orbitals out-of-order because the closeness
of the orbital energies. Usually occurs in heavy metals
Chapter 7
7.1 Development of the Periodic Table [256]
1. Mendeleev helped establish the initial use of a table to organize chemical and physical
properties
2. Moseley successfully arranged the elements by atomic number leaving holes for elements not
discovered
7.2Effective Nuclear Charge [257]
1. Electron’s are attracted to the positive nucleus (Attraction obeys Coulomb’s Law)
a. Greater the nuclear charge the greater the electron’s attraction to the nucleus
b. Further the electrons from the nucleus the weaker the attraction
2. Effective Nuclear Charge (Zeff)- net positive charge experienced by an electron in a manyelectron atom
a. Zeff < Z (actual nuclear charge) because it accounts for the shielding of the nucleus by
the inner electrons
i. Equation
Zeff = Z – S
Zeff = the effective nuclear charge
Z = the number of protons
S = pos. # called the screening constant
ii. S usually is close to the number of core electrons
b. “s” orbitals of an energy level are more attracted to the nucleus than “p” causing “s” to
have a lower energy level. Rule holds true for all orbital energies s<p<d<f.
c. The effective nuclear charge increases as we move across any row due to the increase in
the nuclear charge
d. Going down a column, the effective nuclear charge experienced by the valance
electrons charge is far less than across a row
7.3Sizes of Atoms and Ions [p.259]
1. Background
a. Nonbonding atomic radius- ½ the distance between the nuclei of 2 atoms that are not
bonded together
b. Bonding atomic radius
5
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
i. ½ the distance between the nuclei of two bonded atoms
ii. It is shorter than the nonbonding radius
iii. Used as the atomic radius for comparisons
iv. Atomic radius can be used to determine bond length
v. Covalent radii- come from the bonding atomic radii
2. Periodic Trends in Atomic Radii
a. Increases down a group due to the addition of a principal energy level
b. Decreases from left to right across the periodic table, due to an increase in the nuclear
charge pulling the valence electrons in closer
3. Periodic Trends in Ionic Radii
a. Cations are smaller than their parent atoms because the loss of the valence electrons
which results in less electron-electron repulsions.
b. Anions are larger than the their parent atoms due to an increase in the electron- electron
repulsion
c. Ions in the same group increase in ionic radius as you go down a group
d. Isoelectronic series- a group of ions containing the same # of electrons. Size of ionic
radii will be largest for the highest neg charge and lowest for the highest positive
charge.
7.4 Ionization Energy [p. 264]
1. Background
a. Ionization energy (I)- the minimum amount of energy needed to remove an electron
from the ground stat of the isolated gaseous atom or ion
i. Positive values
ii. Energy put into the atom to remove electrons
b. First ionization energy (I1)– energy to remove the outermost electron
c. Second ionization energy (I2) – energy to remove the 2nd electron from the outer shell
2. Variations in Successive Ionization Energies
a. Ionization energies increase with the removal of each electron because the ion is
becoming more positive
b. Once all of the valence electrons are remove an only the inner- shell or core electrons
remain the ionization energies increase greatly.
i. This supports the concept that only the valance electrons are involved with
bonding because the noble-gas core is to tightly held together
3. Periodic Trends in First Ionization Energies
a. Left to right across a period first ionization energy increases due to the larger positive
nuclear charge
b. Down a group ionization energy decreases due to the shielding of the nucleus by the
inner- shell electrons
c. “s” and “p” elements show a larger range of values than the transition metals
d. A few abnormalities occur due to the electron pairing in the orbitals or the energy level
of lone electrons
4. Electron Configurations of Ions
a. Electrons are moved from the highest principal energy with the highest quantum
number
b. Electrons gained are placed in the empty or partially filled orbital having the lowest
principal energy level
6
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
7.5 Electron Affinities [p.270]
1. electron affinity
a. measures the attraction of an atom for an added electron
b. usually involves a release of energy making it a negative value
c. the greater the attraction the more negative the value
2. positive values indicate that the anion is higher in energy than the separated atom and electron
(very unstable)
3. Halogens have the highest electron affinities (most neg)
4. Noble gases have pos affinities along with Be, Mg and N
7.6 Metals, Nonmetals, and Metalloids
1. Metals
a. Located on the left side of the periodic table (exclude H)
b. Account for aprox. ¾ of the elements
c. Metallic character generally increases as one goes down a group, and is the highest the
closer to the left of the table an element exists
d. Properties
i. Shiny luster
ii. Malleable and ductile
iii. Good conductors
iv. Most metal oxides are ionic solids that are basic
v. Tend to form cations
vi. All but Hg are solids at room temperature
vii. Many metals melt at high temperatures
viii. Have low ionization energies
e. Group 1 ion form 1+ ions, Group 2 form 2+ ions
f. Group 13-17 metals form + ions related to the values of losing the “s” electrons or the
“s” and “p” electrons
g. Transition metals have no obvious pattern
h. Compounds of metals and nonmetals tend to be ionic
2. Nonmetals (found on the right includes H)
a. Vary greatly in appearance
b. Not lustrous
c. Poor conductors
d. Lower melting points than metals
e. 7 form diatomic molecules (5 gases, 1 liquid, one solid)
f. Excluding the noble gases and the diatomic molecules the rest are solids that can be
hard like a diamond or soft like sulfur
g. Tend to gain electrons when reacting with metals
h. Compounds composed of all nonmetals are typically molecular substances
i. Most nonmetal oxides are acidic (reason that acid rain forms)
3. Metalloids
a. Have both metallic and nonmetallic properties
b. Behavior is dependent on what it is combined with
7.7Group Trends for the Active Metals [p.276]
1. Group 1A: the Alkali Metals
a. Soft metallic solids
7
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
b.
c.
d.
e.
f.
g.
h.
i.
j.
Low densities and melting points
Radius increases down a group
Ionization energy decreases down a group (lowest in table)
Form 1+ ions
Only exist in nature as compounds
Can form metal hydrides
React vigorously with water ( reaction more violent for heavy members)
Forms water soluble oxides
Can form metal peroxides where O has a 1- charge and superoxides where O2- is
possible
k. Pure samples are usually stored in a liquid hydrocarbon due to there reactivity
2. Group 2A: the Alkaline Earth Metals
a. Harder solids that are more dense than alkali metals and have higher melting points
b. Low ionization energies
c. Mg and Be are the least reactive
d. Tend to lose two elctrons
e. Emit characteristic colors when excited
7.8 Group Trends for Selected Nonmetals [p.281]
1. Hydrogen
a. Nonmetal
b. High ionization energy
c. Reacts exothermically with nonmetals
d. Can gain electron to form metal hydrides
e. Can donate H+ ions in a reaction
2. Group 6A: the Oxygen Group
a. Oxygen
i. Only gas in group
ii. Diatomic
iii. Has two allotropes O2 and O3 ( different forms of elements in the same state)
1. different properties
2. ozone more dangerous
iv. Forms 2- ion
b. Sulfur
i. Exists in many allotropic forms
ii. Most common is yellow solid with a formula of S8
iii. Gains 2 electrons
iv. Not as reactive
3. Group 7A: the Halogens
a. Melting and boiling points increase with an increasing atomic number
b. All are diatomic
c. F and Cl are gases
d. Br is a liquid
e. I is a solid
f. High electron affinities
g. Tend to gain electrons
h. F and Cl are most reactive
8
Chapter 6 and 7 Notes- Electronic Structure of Atoms and Periodic
Properties of the Elements
i. Fluorine reacts dangerously exothermic
j. Chlorine most used halogen in industry, makes stable aqueous solutions
k. React with metals to form ionic halides (very soluble)
4. Group 8A: the Noble Gases
a. Gases
b. Monoatomic
c. Have filled “s” and “p” orbitals in the outer shell
d. Large first ionization energies
e. Unreactive
9