400 Chemistry
Atomic Structure and Electron Notes
Key questions:
What is matter made up of?
How tiny are the particles?
How do you discover and describe tiny particles that you can’t see?
All matter is made up of ATOMS
Atom = smallest particle of an element that maintains the properties of that element in a
chemical reaction
Theories about how to describe atoms have evolved over time…
Scientist(s)
Democritus
(Greek
philosopher)
Dalton
(English
chemist and
teacher)
Time
frame
~400
B.C.
Key Discoveries or Ideas (items in bold are still known to be
true today)
Atoms are indivisible and indestructible
~1800
Dalton’s Atomic Theory (from experiments & scientific
method):
1. Elements are made up of tiny, indivisible atoms
2. All atoms of the same element are identical
3. All atoms of one element are different from those of
any other element
4. Atoms can mix together physically or chemically
combine in simple, whole number ratios to form
compounds
5. Chemical reactions occur when atoms are separated,
joined, or rearranged
6. Atoms of one element are never changed into atoms
of another element in a chemical reaction
 Did experiments with cathode ray tubes (when high
voltage was applied a stream of negatively charged
particles was produced)
 Discovered electrons (negatively charged particle in
the atom) (rays would bend toward a positive plate)
 Described atoms using the “plum pudding model”
(positive atom with negative electrons sprinkled
throughout)
 Did famous “gold foil experiment”
 Atom is mostly empty space
 Atom has a small, dense region in the center
(nucleus)
 Nucleus contains protons
J.J. Thomson
(English
physicist)
~late
1800s
Ernest
Rutherford
(Thomson’s
student)
~1911
Niels Bohr
(Danish
physicist and
Rutherford’s
student)
Various
scientists
ideas
combined
(Heisenberg,
de Broglie,
Schrodinger)



Electrons surround nucleus and are spread out
No explanation of chemical properties of the elements
~1913
Electrons are found in specific paths (circular orbits)
around the nucleus
 Development of the Bohr Model:
o Electron in a H atom moves around the
nucleus only in certain allowed circular orbits
o Calculated radii for these orbits based on
classical physics
o Tendency of electron to fly off nucleus due to
acceleration must balance its attraction for
the nucleus
o Classical physics said a charged particle that
accelerates should radiate energy, so an electron
should emit energy (light) and collapse into the
nucleus (doesn’t, so Bohr came up with a model
to account for this… more on that later)
o Model fits what is observed for H’s energy
levels but DOES NOT accurately describe
atomic structure (more on that later)
~1920s Quantum mechanical model:
 Heisenberg’s uncertainty principle: There is a
fundamental limitation to how precisely we can
know both the position and momentum of a particle
at a given time
 Electrons are NOT located in specific, fixed, circular
orbits as the Bohr Model suggests (we don’t know
exactly how the electrons are moving)
 Electrons have certain allowed energies and one can
determine how likely it is to find an electron with
that particular amount of energy
 Looks at probabilities of locating an electron at a
point in space (a region of high probability for
finding an electron = an orbital)
 Use quantum numbers to describe the various
properties of an orbital
Atomic Structure Basics
In the nucleus: protons and neutrons
Outside the nucleus: electrons
Protons = positively charged (+1 relative charge), mass of 1 atomic mass unit (1 amu)
Neutrons = neutral (no charge), mass of 1 atomic mass unit (1 amu)
Electrons = negatively charged (-1 relative charge), negligible mass (very tiny)
How do we determine how many protons, neutrons, and electrons an atom has?
Atomic number = Whole number on the Periodic Table, indicates the # of protons
In a NEUTRAL atom (atom with no charge), the # of protons must = the # of
electrons…. So for neutral atoms, atomic number = # of electrons
Atomic mass or average atomic mass = Decimal number on the Periodic Table
Mass number = Atomic mass from Periodic Table rounded to a whole number
Since only protons and neutrons have significant mass…
Mass number = # of protons + # of neutrons
So, to summarize, for a neutral atom:
# of protons = # of electrons = atomic number
# of neutrons = mass number - # of protons = mass number – atomic number
2 types of variations can occur: ions and isotopes
Note: with these variations, the number of protons NEVER changes- it gives the
element its identity
Ion = an atom with the usual number of protons and a different number of electrons,
resulting in a net charge
Isotope = an atom with the usual number of protons and a different number of neutrons,
resulting in a different mass
Ions
Ions are formed when an atom loses or gains electrons
If an atom loses 1 electron, it has a +1 charge; loses 2 electrons, a +2 charge, etc.
If an atom gains 1 electron, it has a -1 charge; gains 2 electrons, a -2 charge, etc.
So, you will have to use the charge to determine the number of electrons (protons
and neutrons will remain THE SAME)
Ex: Na1+ has a +1 charge and has lost 1 electron (so it has 11 protons, 12
neutrons, and 10 electrons)
Ex: N3- has a -3 charge and has gained 3 electrons (so it has 7 protons, 7 neutrons,
and 10 electrons)
Cation = positively charged ion (formed by losing electrons)
Anion = negatively charged ion (formed by gaining electrons)
Isotopes
Isotopes exist when an atom has a different number of neutrons and its mass changes
The number of protons and electrons will remain THE SAME as always
You might see isotopes written like this:
23
Na which means a sodium isotope with a mass number of 23, since Na always has 11
protons, it must have 23-11 = 12 neutrons
24
Na would have 11 protons, 11 electrons, and 24-11 = 13 neutrons
22
Na would have 11 protons, 11 electrons, and 22-11 = 11 neutrons
Note: Sometimes, they will include the atomic number underneath the mass number…
that just saves you the time of having to look it up on the Periodic Table
Isotopes and Average Atomic Mass
Masses on the Periodic Table are a WEIGHTED AVERAGE of all isotopes found in
nature, based on their percent abundance (how much they occur in nature)
Average atomic mass = weighted average of the mass of isotopes in a naturally occurring
sample of an element
Example:
Zinc has 5 isotopes that occur in nature as shown in the table…. Find its average atomic
mass.
Isotope
Mass number
Mass of isotope
Percent abundance
(amu) (sometimes
(%)
you might just have
the mass number
instead of these)
Zinc-64
64
63.929
48.89
Zinc-66
66
65.927
27.81
Zinc-67
67
66.927
4.11
Zinc-68
68
67.925
18.57
Zinc-70
70
69.925
0.62
Average atomic mass of Zn = 0.4889(63.929) + 0.2781(65.927) + 0.0411(66.927) +
0.1857(67.925) + 0.0062(69.925) = 65.39 = mass on the Periodic Table! Voila!
Energy and Electrons
Some scientists believed that electrons behave like waves…
Electromagnetic radiation = a way of energy traveling through space
All waves (gamma, xrays, UV, visible light, IR, microwaves, radio waves, etc.) have 2
characteristics: wavelength and frequency… also all waves in the electromagnetic
spectrum travel at the same speed (the speed of light)
Wavelength = distance between 2 consecutive peaks or troughs in a wave (represented by
Greek letter lambda λ, which has units of length)
Frequency = number of waves per second that pass a given point in space (represented by
Greek letter nu ν, which has units of 1/time or Hertz (Hz) which is 1/seconds)
(examples)
Notice that as wavelength increases, frequency decreases … and that as wavelength
decreases, frequency increases (inversely proportional)
c = λ ν where c is a constant equal to the speed of light (3 x 108 m/s)
As frequency increases, the energy of the wave also increases (directly proportional)
E = h ν = h c / λ where h is a constant called Planck’s constant (h = 6.63x10-34 J s)
Energy is measured in joules (J)
Electromagnetic spectrum = all different types of electromagnetic radiation, ordered by
their wavelengths and frequencies
Electromagnetic spectrum (wavelength shown in meters)
Gamma
10-12
X rays
10-10
Ultraviolet Visible
10-8
10-7
Infrared
10-4
Microwave Radar
10--2
1
Wavelength increases →
Frequency and energy decrease →
For visible light (wavelength shown in meters):
Violet
4x10-7
Green
5x10-7
Wavelength increases →
Frequency and energy decrease →
Orange
6x10-7
Red
7x10-7
Radio
100
(For visible light, color depends on wavelength and frequency)
(examples)
Atomic Emission Spectra and the Bohr Model
Light can be separated by frequency or wavelength
Ex: Rainbows- pass sunlight through a prism, pass sunlight through water droplets
(rain) - can see the spectrum
Emission spectroscopy = analyzing light emitted to identify substances based on the
frequencies of the light emitted




Provide a high energy spark to atoms- they absorb energy and become excited
(electrons jump up to a higher energy level or “excited state”)
Excess energy can then be released (electrons fall back down to “ground state”)producing emission spectrum
Continuous spectrum (ex: white light, made up of all wavelengths of visible light)
looks like stripes of colors blending into each other
Discrete (line) spectrum (ex: specific elements, made of particular wavelengths
only, corresponding to the energies allowed for the electrons since energy is
quantized) looks like detached stripes of color
An emission spectrum for an element is like a fingerprint (each line represents one
frequency, the unique pattern can be used to identify the element)
Observe atomic emission spectrum for an element (i.e., hydrogen) and see discrete lines
indicating that only certain energies are possible (energy is quantized) (i.e., if any energy
level were allowed, the emission spectrum would be continuous)
Bohr’s model produced an expression for the energy levels available to the electron in a
H atom (allows us to determine the energy, wavelength, and frequency of the lines):
First… a bit on the “radius” or size of the energy levels…
Radius is proportional to n2 , where n = the number of the principal energy levels (i.e., the
“principal quantum number”) … so the higher the principal quantum number, the larger
the atomic radius
Now… energy…
En= -RH/n2 Joules
Where En = energy, n is an integer representing the principal quantum number (i.e.,
principal energy level- more on that later), and RH is the Rydberg constant (2.180 x 10-18
Joules) …. Bohr used this to calculate energy levels that match what H shows
experimentally
Another expression of this equation: En= -313.6/n2 kcal/mol (using -313.6 kcal/mol as the
Rydberg constant)
Negative energy???? The negative sign means that the energy of an electron bound to the
nucleus is less than that of an electron infinitely far from the nucleus (where n = infinity
and E = 0) (tells us that the energy of an electron in any orbit is negative relative to that
reference state… in other words, the farther it is from the nucleus, the higher the
energy)
Using this equation, we can compare the energies of electrons in different energy levels
(using n = 1 to represent the ground state and higher n’s to represent various excited
states)
(examples… calculating the energies for n = 1, 2, 3, ….to infinity)
As electrons drop down from the excited state to the ground state, they give off energy:
Change in energy = Energy of final state – Energy of initial state (where energy of the
final and initial states are calculated using the equation above)
If an electron drops down to n= 2, it produces a line in the Balmer series (visible light)
If an electron drops down to n=1, it produces a line in the Lyman series (ultraviolet)
If an electron drops down to n=3, it produces a line in the Paschen series (infrared)
Understanding the signs: Change in energy will be negative when an electron drops to the
ground state from an excited state, indicating that the electron has lost energy by emitting
a photon and is now in a more stable state; change in energy will be positive when an
electron jumps up to an excited state from the ground state, indicating that the electron
has absorbed energy
To calculate the wavelength or frequency of the light emitted when this happens, use
∆E = h ν = h c / λ where h = Planck’s constant = 9.5x10-14 kcal s / mol and c = 3 x1010
cm/s
(examples)
Now- since it is the movement of electrons from one energy level to another that cause a
characteristic emission spectrum- will look at the arrangement of electrons within energy
levels
2 most important points about the Bohr Model:
1. Model correctly fits the quantized energy levels of H atom and postulates only
certain allowed circular orbits for the electron
2. As electron is more tightly bound (closer to the nucleus), its energy becomes more
negative (relative to reference state); as electron is brought closer to nucleus,
energy is released
But Bohr’s model did not work at all when applied to other elements! So, the quantum
mechanical model was developed…
The Quantum Mechanical Model
Bohr’s model is important- explains emissions spectral lines
But….2 major downfalls:
1. Does not account for de Broglie’s hypothesis that electrons behave like waves
2. Does not agree with Heisenberg’s uncertainty principle
(example… measuring the radius of an atom using a photon, height analogy)
So… a new model was needed… the quantum mechanical model aka Schrodinger model
aka wave mechanical model
Based on a math equation
Math equations have shapes…
(examples…)
Schrodinger’s equation:
Solution to Schrodinger equation yields 4 sets of numbers- these are called quantum
numbers
Set = group of numbers that share something in common
Quantum Numbers
Name
Principal quantum
number
Symbol Values
n
1, 2, 3, ….to
infinity
Secondary quantum
number
l
0, 1, 2, …to n-1
Magnetic quantum
number
Magnetic spin
quantum number
ml
+l….0…. to –l
ms
+/- ½
What it Represents
The distance the electron is from
the nucleus, the amount of energy
of the electron
Shape of probability plot for finding
the electron in space (probability
plot = orbital… more on that later)
Number of probability plots,
number of unique ml values
Spin of electron (clockwise or
counterclockwise)
What is an orbital really?
(student example with 95% probability volume element)
So… what are the shapes of the orbitals from Schrodinger’s equation?
l=0
“s” orbital
ml = 0
(drawing)
l=1
“p” orbitals
ml = +1, 0, -1 (3 unique values & 3 orbitals!)
(drawing)
l=2
“d” orbitals
ml = +2, +1, 0, -1, -2 (5 unique values & 5 orbitals!)
(drawings & models)
l=3
“f” orbitals
ml =+3, +2, +1, 0, -1, -2, -3 (7 unique values & 7 orbitals!)
(models… good luck trying to draw them!)
Now…let’s begin to approach how the electrons are arranged in these orbitals
Create a chart in your notes with 7 columns:
-n
-l
-ml
-ms
-specific name of orbital (sublevel)
-number of orbitals
-maximum number of electrons that can be held
To summarize, key points of quantum mechanical model:
 Schrodinger’s wave equation incorporates both the wave and particle properties of
electrons
 Solution to the equation gives wave functions (orbitals) and energies



Determines allowed energies an electron can have (more flexible)
Each orbital describes the distribution of electron density (probability) in space
and has characteristic energy and shape
Orbitals are NOT orbits!!!
Electron Configurations
Atomic orbital or electron cloud = region with a high probability (likelihood) of finding
an electron with a particular energy
Electron configuration = ways in which electrons are arranged in orbitals around the
nucleus
Electron configurations give us insight into an element’s properties and chemical
behavior
Principles that help us understand electron configurations:



Aufbau Principle: Electrons occupy orbitals of lowest energy first
Pauli Exclusion Principle: No more than 2 electrons can occupy any orbital; if 2
electrons are to be in one orbital, they must have opposite spins
Hund’s Rule: When filling degenerate orbitals (orbitals of the same energy), each
orbital in the group will half fill first (receiving 1 electron), before any will double
up
(handouts and diagrams)
Writing Electron Configurations and Drawing Orbital Diagrams



The key is to apply Aufbau, Pauli, and Hund
Electron configuration
o Large number on left = principle energy level, principal quantum number
o Letter = sublevel (orbital), matches letter assigned by secondary quantum
number
o Superscript number = number of electrons in that sublevel
Orbital diagram
o Use circles or boxes to show orbital
o Use up/down arrows to indicate electrons (spin quantum number of +/- ½)
(examples)
Abbreviated Configuration (Noble Gas Abbreviation)
 Go up one row and all the way to the right to the noble gases
 Write that noble gas’s symbol in square brackets
 Write the rest of the configuration
(examples)
Electron Configurations… Additional Notes
 Exceptions:
o Transition metals may have configurations other than what is predicted by
the Periodic Table if they can achieve a full or half-filled set of d orbitals
(examples: copper, silver, chromium)
o Involves a shift of 1 or 2 electrons into an orbital of similar energy
o Slight increase in stability with half-filled or completely filled d orbitals
 Pairing of electrons:
o If in an orbital diagram, an orbital has only 1 electron, this electron is said
to be unpaired
o Can count up total number of unpaired electrons in a substance
o If there are unpaired electrons, solid will be attracted to a magnetic field
and is called paramagnetic
o If no unpaired electrons, solid will be slightly repelled from a magnetic
field and is called diamagnetic
 Isoelectronic: If two species have the same number of electrons, they are said to
be isoelectronic (ex: representative element ions are isoelectronic with whatever
noble gas they attain the configuration of)
Valence Electrons and Lewis Structures
Properties, reactions, and compounds of elements are strongly related to their electron
configurations, specifically their outer electron configurations.
Valence shell = outermost principal energy level (has highest principle energy level
number)
Valence electrons = outermost electrons (those in the valence shell); for most A groups,
this matches their group number in Roman numerals
(examples)
Lewis structure (electron dot structure) = diagram that shows the valence electrons of an
element as dots around the element’s symbol (s outer electrons go on top, then, proceed
by adding a dot to each side for the p electrons until all sides are filled, then double up)
(examples)
Periodic Table and Trends
Periodic Table Development
 1869- Mendeleev and Meyer published classification schemes for elements and
observed that similar chemical and physical properties occur periodically when
elements are in order by atomic mass
 Mendeleev organized them and left blank spaces for undiscovered elements where
he predicted their properties


Periodic Law = when elements are arranged in order of increasing atomic number,
there is a periodic repeating pattern to their properties
Trends in properties are related to the electron configurations of elements and two
phenomena: shielding effect and nuclear charge
http://video.google.com/videoplay?docid=-2134266654801392897
Shielding Effect and Nuclear Charge
 Shielding effect = inner electrons block valence (outer) electrons from the pull of
the positively charged nucleus
o The more principal energy levels, the more “layers of” inner electrons
available to shield the valence electrons
o From left to right on the table, shielding is constant (electrons are being
added to the same principal energy level- no new levels are added)
o From top to bottom on the table, shielding increases (more principal
energy levels are being added)
 Nuclear charge = pull or attraction for electrons by the nucleus
o The higher the atomic number, the more protons in the nucleus, and the
stronger the pull of the nucleus (greater nuclear charge)
o From left to right on the table, nuclear charge increases (more protons
means more nuclear charge)
All other periodic property trends result from these!
Periodic Trends (focusing on representative elements, “A” groups)
Atomic radius (size of the atom)
 From left to right on the table, atomic radius DECREASES
o Why? Shielding is constant but nuclear charge increases
 From top to bottom on the table, atomic radius INCREASES
o Why? Shielding noticeably increases
Ionic radius (size of the ion)
 Cations are smaller than their parent atoms (fewer electrons compared to protons,
electrons pulled in tighter to nucleus)
 Anions are larger than their parent atoms (more electrons compared to protons,
electrons farther out from nucleus)
 Cation size decreases from left to right, anion size decreases from left to right (same
reasons as atomic radius)
 Cation size increases from top to bottom, anion size increases from top to bottom
(same reasons as atomic radius)
Ionization energy (energy required to remove an electron… first IE is the energy
required to remove the outermost valence electron, second IE is the energy to remove
the next one, etc. )
 From left to right on the table, ionization energy INCREASES
o Why? Shielding is constant but nuclear charge increases


o If the electrons are being held tighter, it is harder to remove them (takes more
energy)
From top to bottom on the table, ionization energy DECREASES
o Why? Shielding noticeably increases
o If the electrons are not being held tightly, it takes less energy to remove them
THINK about how the ionization energy might change depending on the electron
configuration of the atoms/ions we are examining (is first IE higher or is second,
third, etc.)
Electronegativity (ability of an atom to attract other electrons to form a bond)
 From left to right on the table, electronegativity INCREASES
o Why? Shielding is constant but nuclear charge increases
o If the electrons of the atom are being pulled more strongly, outside electrons
will also be pulled more strongly
 From top to bottom on the table, electronegativity DECREASES
o Why? Shielding noticeably increases
o If the electrons of the atom are not being pulled strongly, outside electrons
will not be either
Reactivity and Metallic Character
 Metals have a tendency to lose electrons when they react
 Nonmetals have a tendency to gain electrons when they react
 Elements with low ionization energies or high electronegativities will be the most
reactive since they have the strongest tendencies to lose and gain electrons,
respectively (the more reactive elements are in the lower left corner or upper right
corner)
 From left to right on the table, metallic character DECREASES (electrons are not
easily removed, nonmetals are on the right side)
 From top to bottom on the left side of the table, metallic character INCREASES
(electrons are more easily lost)
Periodic Table Review
 Metals, nonmetals, metalloids
o Properties of metals
 Shiny luster
 Malleable and ductile
 Good conductors of heat and electricity
 Form cations (lose electrons)
 High melting points
o Properties of nonmetals
 Dull
 Brittle, hard, or soft
 Poor conductors of heat and electricity
 Form anions (gain electrons)
 Low melting points (some are liquids or gases at room temp)
o Properties of metalloids


 Have properties of both
 Can form cations or anions
 Are often electrical semiconductors
States of matter of elements (how can you tell?)
Alkali metals, alkaline earth metals, transition metals or elements, halogens, noble
or inert gases, inner transition metals, rare earth elements, lanthanide series,
actinide series (be able to find them!)