300 Chemistry
Electrons and the Periodic Table Notes
Introduction to Electrons and Light
Recall Bohr’s idea: Electrons behave like particles and are found in discrete, circular
orbits (energy levels) which surround the nucleus
But- in reality, electrons do not have to stay in fixed, circular orbits, they can move
around
Heisenberg’s uncertainty principle: we can’t really identify exactly where electrons are at
any point in time (but we can get an idea of where they are likely to be)
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Electrons can change level by absorbing or releasing energy
An electron’s usual energy level is known as its “ground” state… if it absorbs
energy, it can jump up to a higher level and become “excited”
When an electron is in an “excited” state (higher energy level) and drops down to
a lower level, it releases energy in the form of light (a photon)
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
Because electrons in atoms release light when they are returning to the ground state after
being excited (energized), we can analyze the patterns of light they emit to learn about
them
Two types of emission spectra: continuous spectra and discrete spectra
 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)
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
The Quantum Mechanical Model
Bohr’s model is important- explains emissions spectral lines
But….2 major downfalls:
1. Does not account for wavelike behavior of electrons
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 which is based on a
math equation
Math equations have shapes…
(examples…)
Schrodinger’s equation was REALLY complicated…
Solution to Schrodinger equation is 4 sets of numbers- these are called quantum
numbers and they give us information about the probability of locating an electron
Probability distributions can have shapes…
(student example with 95% probability volume element)
Atomic orbital or electron cloud = region of high probability (likelihood) of finding an
electron with a particular energy (these have shapes)
To summarize, key points of quantum mechanical model:
 Determines allowed energies an electron can have (more flexible)
 Each orbital describes the probability of electron density in space and has
characteristic energy and shape
 Orbitals are NOT orbits!!!
Orbital Diagrams and Electron Configurations
Electron configurations (how electrons are arranged) give us insight into an element’s
properties and chemical behavior
Principles that help us understand electron configurations:
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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
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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
Principle energy levels actually are made up of sublevels containing orbitals with
different shapes and energies
Each orbital can hold 2 electrons (as long as the electrons have opposite spins)
There are 4 types of sublevels (s, p, d, and f), each made up of orbitals with their own
shapes
 s sublevel has 1 orbital
 p sublevel has 3 orbitals
 d sublevel has 5 orbitals
 f sublevel has 7 orbitals
Principal Energy
Level
1
2
3
4
Number of
Sublevels
1
2
3
4
Type of Sublevels
1s
2s, 2p
3s, 3p, 3d
4s, 4p, 4d, 4f
Maximum Number
of Electrons
2
8
18
32
Orbital diagram = a pictorial representation of how electrons are arranged in orbitals
around the nucleus
Electron configuration = system of showing the arrangement of electrons in an atom
using letters and numbers to represent electrons in the orbitals
Writing Electron Configurations and Drawing Orbital Diagrams
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The key is to apply Aufbau, Pauli, and Hund
Orbital diagram
o Use circles or boxes to show orbital
o Use up/down arrows to indicate electrons
o Always remember to half fill!
Electron configuration
o Large number on left = principle energy level
o Letter = sublevel (orbital)
o Superscript number = number of electrons in that sublevel
(examples, using handouts… starting with hydrogen)
Abbreviated Configuration (Noble Gas Abbreviation)
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Go up one row and all the way to the right to the noble gases
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Write that noble gas’s symbol in square brackets
Write the rest of the configuration
(examples)
Something to think about….
What would the electron configuration or orbital diagram for an ion look like?
Can certain things be “isoelectronic”, meaning that they have the same number of
electrons?
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 electrons go on top, then p’s are on the
other sides… add a dot to each of the other 3 sides until all sides are filled, then double
up)
(examples)
Periodic Table and Trends
Periodic Table Development
 Mendeleev organized the elements by their chemical and physical properties and
left blank spaces for undiscovered elements where he predicted their properties
 Later, upon noticing the pattern of the properties of the elements, the Periodic
Table was reorganized to list elements in order by atomic number
 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
Shielding Effect and Nuclear Charge (use pictures)
 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 and top to bottom on the table, nuclear charge increases
(more protons means more nuclear charge)
o For our purposes, we will ignore the increase in nuclear charge from top to
bottom and only focus on it from left to right because the increase in
shielding from top to bottom is so much more significant
All other periodic property trends result from these!
Periodic Trends (use pictures)
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)
o Cation size decreases from left to right (same reasons)
o Anion size decreases from left to right (same reasons)
o Ion size increases from top to bottom (same reasons)
o Cations are smaller than their parent atoms (why?)
o Anions are larger than their parent atoms (why?)
Ionization energy (energy required to remove an electron to form an ion)
 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
Electronegativity (ability of an atom to attract other electrons)
 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 and form cations when they react
 Nonmetals have a tendency to gain electrons and form anions 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!)